INORGANIC CHEMISTRY. THE NATIONAL DISPENSATORY: CONTAINING THE NAT- ural History, Chemistry, Pharmacy, Actions and Uses of Medicines, including those recognized in the Pharmacopoeias of the United States, Great Britain, and Germany, with numerous references to the French Codex, By Alfred Stilß, M.D., LL.D., Professor Emeritus of the Theory and Practice of Medicine and of Clinical Medicine in the University of Pennsylvania, and John M. Maisch, Phar. D., Professor of Mat. Med. and Botany in Phila. College of Pharmacy, Sec’y to the American Pharmaceutical Association. Third edition, thoroughly revised and greatly enlarged. In one magnificent imperial octavo volume of 1767 pages, with 311 fine engravings. Cloth, $7.25; leather, $8.00; half Russia, open back, $9.00. With Denison’s “Ready Reference Index” $l.OO in addition to price in any of above styles of binding. A MANUAL OF CHEMICAL ANALYSIS, AS APPLIED TO THE Examination of Medicinal Chemicals and their Preparations. Being a Guide for the Determination of their Identity and Quality, and for the Detection of Impurities and Adulterations. For the use of Pharmacists, Physicians, Drug- gists, and Manufacturing Chemists, and Pharmaceutical and Medical Students. By F. Hoffmann, A.M., Ph.D., Public Analyst to the State of New York, and F. B. Power, Ph.D., Prof, of Anal. Chem. in the Phil. College of Pharmacy. Third edition, entirely re-written and much enlarged. In one very handsome octavo volume of 621 pages, with 179 illustrations. Cloth, $4.25. MEDICAL PHYSICS. A TEXT-BOOK FOR STUDENTS AND Practitioners of Medicine. By John C. Draper, M.D., LL.D., Professor of Chemistry in the University of the City of New York. In one octavo volume of 734 pages, with 376 woodcuts, mostly original. Cloth, $4.00. Just ready. CHEMISTRY, GENERAL, MEDICAL AND PHARMACEUTICAL; Including the Chemistry of the U. S. Pharmacopoeia. A Manual of the General Principles of the Science, and their Application to Medicine and Pharmacy, By John Attfield, Ph.D., Professor of Practical Chemistry to the Pharmaceuti- cal Society of Great Britain, etc. A new American, from the tenth English edition, specially revised by the Author. In one handsome royal 12mo. volume of 728 pages, with 87 illustrations. Cloth, $2.50; leather, $3.00. TEXT-BOOK OF PHYSIOLOGY. BY MICHAEL FOSTER, M.D., F.R.S., Professor of Physiology in Cambridge University, England. Third American from the fourth English edition, with notes and additions by E. T. Reichert, M.D. In one handsome royal 12mo. volume of 908 pages, with 271 illustrations. Cloth, $3.25; leather, $3.75. Just ready. Detailed Catalogue sent to any address on application to LEA BROTHERS & CO., Philadelphia. INORGANIC CHEMISTRY, BY EDWARD FRANKLAND, Ph.D., D.C.L., LL.D., F.R.S., ui’’ ’ ’ ’ PROFESSOR OF CHEMISTRY IN THE NORMAL SCHOOL OF SCIENCE, LONDON, AND FEANCIS E. JAPP, M.A., Ph.D., F. 1.&, ASSISTANT PROFESSOR OF CHEMISTRY IN THE NORMAL SCHOOL OF SCIENCE, LONDON. WITH 51 ILLUSTRATIONS AND A PLATE. PHILADELPHIA: LEA BROTHERS & CO. 1885. PREFACE. The Lecture Notes for Chemical Students, already published hy one of us and now in their third edition, were always intended to he the precursors of text-books on Mineral and Organic Chemistry. The present volume fulfils this intention so far as Inorganic Chem- istry is concerned. It is constructed on those principles of Classi- fication, Nomenclature, and Notation which, after an experience of nearly twenty years, have been found to lead most readily to the acquisition of a sound and accurate knowledge of elementary chem- istry. In the Introduction we have endeavored to present to the student a connected account of the chief chemical theories at present prevail- ing, introducing only so much descriptive matter as is necessary for the elucidation of the subject. Afterwards, in the descriptive part of the work, the necessary references to the theoretical portion are given. In some of the theoretical sections, we have followed modes of treat- ment adopted hy H. Kopp, Lothar Meyer, and Naumann in their well-known works. We have also to express our obligations to Fittig’s excellent “ Grundiss der unorganischen Chemie.” Although it would be out of place, in an elementary work like the present, to impart detailed instruction in the technical applications of chemistry, we have not hesitated to give brief outlines of some of the more important of these applications. Normal School op Science and Eoyal School of Mines, South Kensington, London September, ISB4. TABLE OF CONTENTS. INTROLUCTION. CHAPTER I. MATTER AND FORCE. Matter and motion. Forces of nature, .33 Distinguishing characteristics of chemical force, . 34 CHAPTER 11. ELEMENTS AND COMPOUNDS, Simple and compound matter, 37 Table of elements, . 38 CHAPTER 111. CHAPTER 111. Nomenclature of elements, 40 CHEMICAL NOMENCLATURE, Nomenclature of compounds, . 40 . CHAPTER IV. LAWS OF COMBINATION, Law of constant proportions, ,45 Law of multiple proportions, .46 Law of equivalent proportions, 46 CHAPTER V. CHAPTER V. Atoms, .48 THE ATOMIC THEORY, Molecules, 48 CHAPTER VI. MOLECULAR WEIGHTS, Boyle’s Law, 52 Law of Charles, 53 Law of Avogadro, 53 Law of Gay-Lussac, 54 Hofmann’s volume-symbols, 56 Determination of molecular weights, 59 VIII TABLE OF CONTENTS. CHAPTER VII. ATOMIC WEIGHTS. Deductions of the atomic weight of an element from the vapor-density of its com- pounds, 61 Apparent exception to Avogadro’s law, . ... . .63 Determination of atomic weights by means of isomorphism, 64 Determination of the atomic weights from the specific heats of the elements in the solid state, 67 CHAPTER VIII. CHAPTER VIII. CHEMICAL NOTATION. ATOMICITY. Symbolic notation, 75 Atomicity of elements, 78 Graphic notation, 82 Calculation of formulse, 84 CHAPTER IX. CHAPTER IX. COMPOUND RADICALS. List of compound radicals, 86 Atomic and molecular combination, ■ . . 87 CHAPTER X. CHAPTER X. CLASSIFICATION OF ELEMENTS. Classification of the elements according to atomicity, 1 88 Classification of the elements according to their atomic weights. The Periodic Law, 90 Curve of the atomic volumes of the elements, 95 CHAPTER XI. CHAPTER XI. RELATIONS BETWEEN CHEMICAL COMPOSITION AND SPECIFIC GRAVITY. Atomic and molecular volumes, 96 ATOMIC VOLUME. Molecular volume of gases, 96 Molecular volume of solids, 97 Molecular volume of liquids, 98 CHAPTER XII. CHEMICAL AFFINITY. Extent and intensity of chemical affinity, 102 Modes of chemical action, 102 Combination. Decomposition, ■ 103 Dissociation, 103 Electrolysis, 104 Electro-chemical equivalents, 107 TABLE OF CONTESTS. IX CHAPTER XIII. CHEMICAL HOMOGENEITY. Homogeneity of gases, 109 Homogeneity of liquids and solids, 109 CHAPTER XIV. ISOMERISM, METAMERISM, POLYMERISM, ALLOTROPY. Differences of chemical character in compounds of the same composition, . . 110 Allotropy, HI CHAPTER XV. CHAPTER XV. HEAT OF CHEMICAL COMBINATION.—THERMOCHEMISTRY. Laws of thermochemistry, 11l CHAPTER XVI. CHAPTER XVI. FUSION AND FUSING-POINTS. Change of volume accompanying fusion, 117 Effect of pressure in altering the fusing-point, 117 Latent heat of fusion, 117 Suspended solidification, 119 CHAPTER XVII. EBULLITION AND BOILING-POINTS. Taper tension, 119 Law regulating boiling-points, 120 Latent heat of vapors, 122 Liquefaction of gases, 123 CHAPTER XVIII. CHAPTER XVIII. SOLUTION. . Solubility of liquids, 124 Solubility of gases, 124 Solubility of solids, 125 Supersaturation or suspended crystallization, 128 CHAPTER XIX. Phenomena of diffusion, 128 DIFFUSION. Diffusion of liquids. Dialysis, ; 129 Diffusion of gases, 130 CHAPTER XX. Systems of crystals, .132 CRYSTALLOGRAPHY. X TABLE OF CONTEXTS. CHAPTER XXI. WEIGHTS AND MEASURES. French and English systems, 136 Conversion of French into English weights and measures, 136 The crith, 137 NON-METALS. CHAPTER XXII. MONAD ELEMENTS. Section I. Hydrogen, 140 Section 11. Chlorine, 151 Hydrochloric acid, 156 CHAPTER XXIII. CHAPTER XXIII. DYAD ELEMENTS. Section I. Oxygen, 160 Allotropic oxygen or ozone, 166 Compounds of oxygen with hydrogen, 169 Compounds of chlorine with oxygen and hydroxyl, 177 CHAPTER XXIV. Section I. Boron, * 185 TRIAD ELEMENTS. Compound of boron with hydrogen, 187 Compounds of boron with the halogens, 188 Compounds of boron with oxygen and hydroxyl, 190 CHAPTER XXV. CHAPTER XXV. TETRAD ELEMENTS. Section I. Carbon, 193 Compounds of carbon with oxygen, 200 CHAPTER XXVI. CHAPTER XXVI. Section I. Nitrogen, 211 PENTAD ELEMENTS. Compounds of nitrogen with oxygen and hydroxyl, 213 Compounds containing nitrogen, chlorine, and oxygen, 228 Compounds of nitrogen with hydrogen and hydroxyl, 230 Compounds of nitrogen with chlorine, bromine, and iodine, 236 The atmosphere, 237 TABLE OF CONTENTS. XI CHAPTER XXVII. HEXAD ELEMENTS. Section I. Sulphur, 243 Compounds of sulphur with hydrogen, 249 Compounds of sulphur with the halogens, 254 Compound of sulphur with carbon, 256 Compound of sulphur with carbon and oxygen, 258 Compounds of sulphur with oxygen and hydroxyl, ......... 259 Compounds of sulphur with oxygen and chlorine (oxychlorides, acid chlo- rides), ' 281 Selenium, 283 Compounds of selenium with hydrogen and chlorine, 285 Compounds of selenium with oxygen and hydroxyl, 286 Tellurium, 287 Compounds of tellurium with hydrogen, chlorine, and oxygen, .... 288 CHAPTER XXVIII. MONAD ELEMENTS, Section II (continued). Bromine, 290 Hydrobromic acid, 292 Compounds of bromine with oxygen and hydroxyl, 293 lodine, . 295 Hydriodic acid, . 298 Compounds of iodine with chlorine, 300 Fluorine, 306 Compounds of iodine with oxygen and hydroxyl, 301 Hydrofluoric acid, 307 CHAPTER XXIX. TETRAD ELEMENTS. Section I (continued). Silicon, 309 Compound of silicon with hydrogen, 311 Compounds of silicon with the halogens, 313 Compounds of silicon with oxygen and hydroxyl, 316 Compounds of silicon containing sulphur, 320 Tin, 321 Compounds of tin, 323 Compounds of tin with the halogens, . 324 Compounds of tin with oxygen and hydroxyl, 326 Compounds of tin with sulphur, 328 General character and reactions of the salts of tin, 329 Titanium, 330 Compounds of titanium with chlorine, 331 Compounds of titanium with oxygen and hydroxyl, 332 Compounds of titanium with nitrogen and with nitrogen and carbon, . . 332 General character and reactions of the titanium compounds, 333 TABLE OF CONTENTS. Zirconium, 333 Compounds of zirconium, 334 Thorium, . 334 Compounds of thorium, 334 CHAPTER XXX. PENTAD ELEMENTS. Section I. (continued). Phosphorus, 335 Compounds of phosphorus with hydrogen, 340 Compounds of phosphorus with the halogens, . 3-14 Compounds of phosphorus with oxygen and hydroxyl, 348 Compounds of phosphorus with chlorine and oxygen, 309 Compounds of phosphorus with sulphur, ............ 361 Compound of phosphorus with sulphur and chlorine, 362 Phosphorus compounds containing nitrogen, 363 Vanadium, 364 Compounds of vanadium with chlorine, 365 Compounds of vanadium with oxygen and hydroxyl, 365 Arsenic, 366 Compound of arsenic with hydrogen, 367 Compounds of arsenic with the halogens, 369 Compounds of arsenic with oxygen and hydroxyl, ......... 370 Compounds of arsenic with sulphur and hydrosulphyl, 373 General properties and reactions of the compounds of arsenic, .... 376 Riorum and Tantalum, 378 Compounds of niobium and tantalum, 378 Antimony, 378 Compound of antimony with hydrogen, 380 Compounds of antimony with the halogens, 381 Oxides and acids of antimony, 383 Compounds of antimony with sulphur, 387 Sulphantimonites, 389 General properties and reactions of the compounds of antimony, .... 390 Bismuth, 391 Halogen and oxyhalogen compounds of bismuth, 391 Compounds of bismuth with oxygen and hydroxyl, 392 Compounds of bismuth with sulphur, 395 General properties and reactions of the compounds of bismuth, .... 396 METALS. CHAPTER XXXI. DISTINGUISHING CHARACTERISTICS OF THE METALLIC ELEMENTS. Chief points of difference between metals and non-metals, 397 Relation of the metals to heat, 398 TABLE OF CONTENTS. XIII Relations of the metals to light, . 399 Spectrum analysis, 400 Relations of the metals to gravity, 406 Cohesive power, 407 Alloys, 410 CHAPTER XXXII. CHAPTER XXXII. MONAD ELEMENTS. Section 111. Potassium, 411 Compound of potassium with hydrogen, 413 Compounds of potassium with the halogens, 414 Compounds of potassium with oxygen, 414 Compound of potassium with hydroxyl, 415 Oxy-salts of potassium, 416 Compounds of potassium with sulphur, 420 Compound of potassium with hydrosulphyl, 421 Sulpho-ifalts of potassium, 423 Compound of potassium with nitrogen and hydrogen, 423 General properties and reactions of the compounds of potassium, . ■. . . 424 Sodium, 424 Compound of sodium with hydrogen, 426 Compounds of sodium with the halogens, 426 Compounds of sodium with oxygen and hydroxyl, 427 Oxy-salts of sodium, . 427 Compounds of sodium with sulphur and hydrosulphyl, ....... 435 Sulpho-salts of sodium, 435 Compound of sodium with nitrogen and hydrogen, 435 Lithium, 435 General properties and reactions of the compounds of sodium, 435 Compounds of lithium with the halogens, . 436 Compounds of lithium with oxygen and hydroxyl, 436 Oxy-salts of lithium, 437 Rubidium, 438 General properties and reactions of the compounds of lithium, 437 Compounds of rubidium, 438 Cassium, 439 Compounds of csesium, 440 General properties and reactions of the compounds of rubidium and csesium, 440 The Ammonium Salts, 440 Compounds of ammonium with the halogens, 441 Compound with hydroxyl, 442 Oxy-salts of ammonium, 442 Compounds of ammonium with sulphur and hydrosulphyl, 446 General properties and reactions of the ammonium salts, 446 Section IV. Silver, 447 Compounds of silver with the halogens, 452 Compounds of silver with oxygen, 456 Oxy-salts of silver, 456 XIV TABLE OF CONTENTS. Compounds of silver with sulphur, 459 Sulpho-salts of silver, 459 Compounds of silver with nitrogen and phosphorus, ........ 459 General properties and reactions of the compounds of silver, 459 CHAPTER XXXIII. CHAPTER XXXIII. DYAD ELEMENTS. Section 11. Barium, 460 Compounds of barium with the halogens, , 461 Compounds of barium with oxygen, 462 Compound of barium with hydroxy], 463 Oxy-salts of barium, 464 Compounds of barium with sulphur, 467 Compound of barium with hydrosulphyl, 467 Strontium, 468 General properties and reactions of the compounds of barium, ..... 468 Compounds of strontium with the halogens, 468 Compounds of strontium with oxygen and hydroxyl, 469 Oxy-salts of strontium, 469, Calcium, 471 General properties and reactions of the compounds of strontium, .... 470 Compounds of calcium with the halogens, 472 Compounds of calcium with oxygen, 474 Compound of calcium with hydroxyl, 474 Oxy-salts of calcium, 475 Glass, 480 Compounds of calcium with sulphur, 483 Compound of calcium with phosphorus, 483 General properties and reactions of the compounds of calcium, .... 484 On potable water and on the impurities occurring in natural waters, . . 484 Magnesium, 507 Compounds of magnesium with the halogens, 508 Compounds of magnesium with oxygen and hydroxyl, 509 Oxy-salts of magnesium, 509 Compounds of magnesium with sulphur and hydrosulphyl, 513 Compounds of magnesium with nitrogen and with boron, 513 Compound of magnesium with silicon, 513 General properties and reactions of the compounds of magnesium, . . . 513 Zinc, ■ 514 Compounds of zinc with the halogens, 516 Compounds of zinc with oxygen and hydroxyl, 517 Oxy-salts of zinc, 518 Compounds of zinc with sulphur, . . 519 Compound of zinc with the pentad elements, 520 General properties and reactions of the compounds of zinc, 520 Beryllium, 521 Compounds of beryllium with the halogens, 521 Compounds of beryllium with oxygen and hydroxyl, 522 Oxy-salts of beryllium, 523 TABLE OF CONTENTS. XV Compound of beryllium with sulphur, 523 General properties and reactions of the compounds of beryllium, .... 523 CHAPTER XXXIV. CHAPTER XXXIV. Section 111. Cadmium, 524 DYAD ELEMENTS. \ Compounds of cadmium with the halogens, • 525 Compounds of cadmium with oxygen and hydroxy], 525 Oxy-salts of cadmium, 625 Compound of cadmium with sulphur, 526 General properties and reactions of the compounds of cadmium, .... 526 Mercury, 527 Amalgams, 529 Compounds of mercury with the halogens, 530 Compounds of mercury with oxygen, 632 Oxy-salts of mercury, 533 Compounds of mercury with sulphur, 535 Compound of mercury with nitrogen, 536 Ammoniacal mercury compounds, 536 Characteristic properties and reactions of the compounds of mercury, . . 537 Copper, 538 Compound of copper with hydrogen, 542 Compounds of copper with the halogens, 542 Compounds of copper with oxygen and hydroxyl, 544 Oxy-salts of copper, 546 Compounds of copper with sulphur, 549 Compounds of copper with nitrogen, phosphorus, and arsenic, 550 General properties and reactions of the compounds of copper, ..... 550 CHAPTER XXXV. TRIAD ELEMENTS, Section 11. Goud, 551 Compounds of gold with the halogens, 553 Compounds of gold with oxygen and hydroxyl, 554 Oxy-salts of gold, 555 Compound of gold with sulphur, . 556 Thallium, 556 General properties and reactions of the compounds of gold, 556 Compounds of thallium with the halogens, 557 . Compounds of thallium with oxygen and hydroxyl, 558 Oxy-salts of thallium, 559 Compounds of thallium with sulphur, t ... . 560 General properties and reactions of the compounds of thallium, .... 561 Indium, 561 Compounds of indium with the halogens, 562 Compounds of indium with oxygen and hydroxyl, 662 Oxy-salts of indium, 563 Compounds of indium with sulphur 563 General properties and reactions of the compounds of indium, .... 563 XVI TABLE OF CONTEXTS. CHAPTER XXXVI. TETRAD ELEMENTS. Section 11. Aluminium, 564 Compounds of aluminium with the halogens, 566 Compounds of aluminium with oxygen and hydroxyl, 567 Oxy-salts of aluminium, 568 Ultramarine, 573 Porcelain and pottery, 573 Compound of aluminium with sulphur, 576 General properties and reactions of the compounds of aluminium, . . . 576 Gallium, 576 Compounds of gallium, 577 General properties and reactions of the compounds of gallium, 577 CHAPTER XXXVII. CHAPTER XXXVII. METALS OP THE RARE EARTHS.—TETRAD ELEMENTS. Section 111. Cerium, 578 Compounds of cerium, 580 PENTAD ELEMENTS. Section 11. Didymium, 581 Compounds of didymium, 581 Section IY. Lanthanum, 582 TRIAD ELEMENTS. Compounds of lanthanum, 582 Yttrium, 582 Compounds of yttrium, 584 Erbium, 584 Terbium, Scandium, Samarium, Decipium, 585 Compounds of erbium, 584 General properties and reactions of the rare earth metals, 585 CHAPTER XXXVIII. TETRAD ELEMENTS. Section IV. Platinum, 586 Compounds of platinum with the halogens, 588 Compounds of platinum with oxygen and hydroxyl, 589 Oxy-salts of'platinum, 590 Compounds of platinum with sulphur, 590 Ammonium compounds of platinum (platinaraines), 591 General properties and reactions of the compounds of platinum, .... 591 Palladium, 592 Compound of palladium with hydrogen, 593 Compounds of palladium with the halogens, 593 Compounds of palladium with oxygen, 594 TABLE OF CONTENTS. XVII Fallacious oxy salts, 594 Compounds of palladium with sulphur, 594 General properties and reactions of the compounds of palladium, .... 595 Iridium, 595 Compounds of iridium with the halogens, 596 Compounds of iridium with oxygen, 597 Oxy-salts of iridium, . . . 598 Compounds of iridium with sulphur, 598 General properties and reactions of the compounds of iridium, .... 598 Rhodium, 598 .Compound of rhodium with chlorine, 599 Compounds of rhodium with oxygen, 599 Oxy-salts of rhodium, 599 Compound of rhodium with sulphur, 599 General properties and reactions of the compounds of rhodium, .... 599 OCTAD ELEMENTS. Osmium, 600 Compounds of osmium with chlorine, 601 Compounds of osmium with oxygen, 601 Oxy-salts of osmium, 602 The osmates, 602 Compounds of osmium with sulphur, 602 General properties and reactions of the compounds of osmium, .... 602 Ruthenium, 602 Compounds of ruthenium with the halogens, 603 Compounds of ruthenium with oxygen, 603 Oxy-salts of ruthenium, 604 Euthenates and perruthenates, 604 Compound of ruthenium with sulphur, 605 General properties and reactions of the compounds of ruthenium, . . . 605 CHAPTER XXXIX. CHAPTER XXXIX. TETRAD ELEMENTS. Section Y. Lead, 605 Compounds of lead with the halogens, 607 Compounds of lead with oxygen, 608 Oxy-salts of lead, 610 Compound of lead with sulphur, 613 General properties and reactions of the compounds of lead, .613 CHAPTER XL. HEXAD ELEMENTS. Section 11. Uranium, 614 Compounds of uranium with the halogens, 615 Compounds of uranium with oxygen, 615 Oxy-halogen compounds of uranium, 616 Oxy-salts of uranium, 616 The uranates, 617 2 TABLE OF CONTESTS. Compounds of uranium with sulphur, 618 Molybdenum, 619 General properties and reactions of the compounds of uranium, .... 618 Compounds of molybdenum with the halogens, 619 Compounds of molybdenum with oxygen, 620 The molybdates, 621 Phospho-molybdic acid, 622 Compounds of molybdenum with sulphur, 623 General properties and reactions of the compounds of molybdenum, . . . 623 Tungsten, 623 Compounds of tungsten with the halogens, 624 Compounds of tungsten with oxygen, 625 The tungstates 626 Silico-tungstic acids, 627 The tungsto-tungstates, 628 Compounds of tungsten with sulphur, 628 General properties and reactions of the compounds of tungsten, .... 628 CHAPTER XLI. CHAPTER XLI. Section 111. Chromium, 629 HEXAD ELEMENTS. Compounds of chromium with the halogens, 630 Compounds of chromium with oxygen, 631 Oxy-salts of chromium, . 633 The chromites, . . . 634 The chromates, 635 Compounds of chromium with oxygen and chlorine, 638 Compound of chromium with sulphur, 639 Compound of chromium with nitrogen, 639 General properties and reactions of the compounds of chromium, . . . 639 Manganese, 640 Compounds of manganese with the halogens, 641 Compounds of manganese with oxygen, 642 Oxy-salts of manganese, 646 The manganates, 647 Permanganic acid and permanganates, 648 Compound of manganese with oxygen and chlorine, 649 Compounds of manganese with sulphur, 649 General properties and reactions of the compounds of manganese, . . . 650 Iron, 650 Compounds of iron with the halogens, 655 Compounds of iron with oxygen, 657 Oxy-salts of iron, 659 The ferrates, 661 Compounds of iron with sulphur, 661 Cobalt, 663 General properties and reactions of the compounds of iron, 662 Compounds of cobalt with the halogens, 664 Compounds of cobalt with oxygen, 665 TABLE OP CONTENTS. Oxy-salts of cobalt, g66 Compounds of cobalt with sulphur, 667 Ammonium compounds of cobalt (cobaltamines), 668 Nickel, General properties and reactions of the compounds of cobalt, 669 Compounds of nickel with the halogens, 672 Compounds of nickel with oxygen, 672 Oxy-salts of nickel, 673 Compounds of nickel with sulphur, 673 General properties and reactions of the compounds of nickel, 674 Norwegium, 674 INORGANIC CHEMISTRY CHAPTER I. MATTER AND FORCE. In the most cursory observation of the objects surrounding us, our attention is arrested by two things—matter and motion. We see clouds drifting over our heads and rain falling from these clouds. The descending water flows in river beds or plunges in cataracts down precipices, making its way in both cases to the sea. The surface of that sea is in constant motion, whilst ships driven by wind or steam make their way through its waters. On land, animal life everywhere exhibits matter in motion. The air is rarely still, and many of the heavenly bodies are constantly changing their places in the sky. All this we cannot help observing; a some- what more minute examination, however, shows us that matter not only thus suffers a change of place, but that it also frequently undergoes other changes. Thus water becomes ice or steam, iron rusts, coal burns, and certain substances such as glass and sealing- wax acquire, when rubbed, the property of attracting light bodies. Now this motion of matter and these changes which matter undergoes are all brought about by what is termed force. This force assumes several different forms, which are sometimes regarded and generally described as distinct forces: thus the transfor- mation of water into ice and steam is due to the operation of two of these forces which act antagonistically to each other, and are termed cohesion and heat; the rusting of iron and the burning of coal are brought about by chemical force; the impression pro- duced upon the eye by the combustion of coal is due to light; the attractive power of the glass and sealing-wax is the effect of the electric force; whilst the motion of the heavenly bodies and that of water from the clouds to the sea are the result of the action of a force called gravity. The department of knowledge which deals with these phenomena is termed Natural Science. Natural science studies and investigates the whole range of sensible objects. It teaches us the properties of these objects and the various changes which they undergo, either in the ordinary course of nature or by the application of extraordinary and arti- 34 INORGANIC CHEMISTRY. ficial means. This vast field of observation and research has been divided into two great sections, viz.: 1. Statical sciences. 2. Dynamical sciences. The statical sciences study objects in a state of rest with reference to their form, magnitude, situation, structure, and other properties; such branches of science are Descriptive Astronomy and Geology, Mineralogy, Botany, Zoology, Animal and Vegetable Anatomy. The dynamical sciences take into consideration the changes to which sensible objects are subject. They are subdivided into two groups. The first group studies these changes without reference to their causes: such are Physical Astronomy and Geology, and Animal and Vegetable Physiology. The second group investigates the changes which bodies undergo with special reference to the causes of such changes. These are Physics, and Chemistry. This classification of the natural sciences, however, must not be taken in too strict a sense, especially in the case of the second section, for the astronomer and geologist are nowadays rarely content to observe changes without inquiring into their causes: the same is still more frequently the case with the physiologist, and thus physics and chemistry are continually appealed to in the development of astronomy, geology, and physiology. The force to which the phenomena of chemistry are primarily ascribed, and which is commonly termed chemical affinity, is there- fore closely associated with the other great forces of nature, but it is sharply distinguished from them, in the first place, by producing permanent changes in the properties of the bodies subject to its action. The other forces do not permanently alter the properties of matter, but when substances are brought under the influence of chemical force, they are scarcely if at all afterwards recognizable by the unaided senses. The presence of the bright, hard, colorless and heavy metal iron, could not even be suspected in the dull, soft, brown, and comparatively light rust, into which it is converted by exposure to the air; still less, perhaps, could the rust be credited with the presence of the colorless and invisible gas, oxygen, which is held in combination with the iron by chemical energy. The change is such as is not produced by mixture only. Mechanical mixture, however intimate, does not conceal the properties of iron and sulphur, for instance. The magnetic quality of the iron is as marked as ever, and the two constituents may be distinguished and even separated from each other under the microscope. But after these substances have been subjected to chemical action, the most powerful microscope is incompetent to detect either sulphur or iron, and the magnetic property of the metal almost entirely disappears. This change of properties is manifested in various ways: sometimes liquids or gases are converted into solids, or vice versd, sometimes a change in color, taste, odor, or medicinal pro- MATTER AND FORCE. 35 perties is produced, and there is always a change of temperature, sometimes in the direction of heat, and sometimes in that of cold. With all these changes, however, there is never the slightest altera- tion in the weight of the materials operated upon. In the second place, chemical affinity cannot act through an appreciable intervening space. Heat, light, and electricity affect bodies at considerable distances, whilst gravity acts through spaces inconceivably great; but if two substances, between which the chemical force is energetically exerted when they are in contact, be placed at the smallest appreciable distance from each other, no chemical action whatever occurs, even after they have been in close proximity for Of all other forces, cohesion alone requires this intimate contact. If two pieces of plate glass be gently placed one upon the other, the slightest effort suffices to separate them, but if they be pressed together, they markedly cohere, and if strongly pressed for a long time, they can no longer be separated. The two pieces have become one by cohesion, but the properties of the glass are unaltered, and cohesive action is thus sharply distin- guished from chemical action. The most distinguishing characteristic of the chemical force, however, is the limitation of its action to fixed and definite quanti- ties of matter. Each chemical compound not only always contains the same kinds of matter, but its constituents are always present in exactly the same proportions, although the specimens of the compound may have been derived from the most widely different sources. Thus water obtained from melting snow, from rain, from steam or from the artificial combination of its constituents, always consists of oxygen and hydrogen in the proportion of one part by weight of the latter to eight parts of the former. Again, common salt, whether obtained naturally from the mines of Cheshire or Poland, from the brine springs of Germany or America, from the salt lakes of Russia or Australia, from sea water, or prepared artificially from its constituents, always consists of chlorine and sodium in the proportion of 35.5 parts of the former to 23 parts of the latter. When two bodies combine chemically or become united together by the chemical force, they do so in fixed and definite proportions. The materials composing our universe are bound together by a force which, whether regarded as attraction or as pressure, pro- duces three sets of phenomena differing so much from each other as to lead to their being commonly referred to three of the distinct forces already mentioned. One of these is gravitation, which acts through distances inconceivably great. The second is cohesion, which acts only through spaces too small to be measured. The third is chemical attraction or chemical affinity which, like cohe- sion, also acts through distances too small to be measured, but which, as already mentioned, is distinguished, both from gravita- tion and cohesion, by producing a change of properties in the oiatter upon which it acts. Thus a lump of ice presses towards the centre of the earth, 36 INORGANIC CHEMISTRY. being pulled in that direction by the attraction of gravitation, which can be overcome by mechanical means. The lump of ice is made up of smaller pieces, for it can be broken up into an immense number of particles by mere mechan- ical effort, and thus cohesive attraction, like gravity, is overcome by mechanical means; but only partially, for each particle is made up of smaller particles still bound together by the same force. If, however, heat be applied to the ice, another well-marked step in the conquest of cohesion is gained, and liquidity is induced—a condition in which the particles of the water move freely about and amongst each other. But even here cohesion is not completely vanquished, and the particles still cling to each other with a con- siderable amount of tenacity. By the application of a greater amount of heat, the complete conquest of cohesion is at last achieved. In the condition of steam, the particles of water no longer stick together: they are entirely freed from all cohesive force, and are only restrained from flying asunder to infinite dis- tances by gravitation and external impediments. In all these operations, the properties of the water have not been essentially or permanently altered. Even steam is, like water, uninflammable and incapable of supporting combustion. More- over, on cooling, it is reconverted into water with all its properties unimpaired. By heat, cohesion has thus been gradually but completely over- come, and the question now arises, can any further effect be pro- duced upon water by the same agent? Experiment answers this question in the affirmitive, for if steam be subjected to the intense heat of a stream of electric sparks, it is resolved into a mixture of oxygen and hydrogen gases which refuses to condense to water on cooling, and which explodes by contact with flame. The proper- ties of the steam have thus been entirely altered, and by this in- tense heat another remarkable step has been taken in the conquest of attractive force; each particle of steam has been broken up, and by the change of properties which has followed the rupture, the attraction overcome is recognized as that of chemical affinity. The attractive forces thus operating within a mass of ice are enormous. They may be expressed in terms either of heat or of mechanical effort. In terms of heat ice requires as much heat to melt it, that is to convert it into liquid and ice-cold water, as would raise the temperature of an equal weight of water from 0° C. to 79.2° C. Water at 0° requires to convert it into steam as much heat as would raise its temperature to 637° C. if no steam were formed. But the separation of the oxygen from the hydrogen absorbs as much more heat as would raise the temperature of the steam to 10,315° C. if no separation occurred. In terms of mechan- ical effort the force required to convert 9 lbs. of ice into water is equal to that required to raise a weight of one ton to a height of 433 feet, to overcome the remaining cohesion and convert the water into steam requires a force sufficient to raise one ton to a height of 2,900 feet, whilst the power required for the separation ELEMENTS AND COMPOUNDS. 37 of the two constituents of steam would raise one ton a height of no less than 22,320 feet. CHAPTER 11. ELEMENTS AND COMPOUNDS. All kinds of matter which we meet with on the earth may be divided into two classes, those which are capable of resolution into other simpler kinds of matter, and those which defy our attempts so to resolve them. The former are termed compounds; the latter, simple bodies or dements. For example, if red oxide of mercury be heated, the heat will exert, as in the case of steam already described, a disintegrating or decomposing action: the red oxide will break up into two substances—a colorless gas, oxygen; and a white heavy liquid, mercury. If the mercury and the ox}Tgen be carefully weighed, it will be found that their weights are together exactly equal to that of the oxide of mercury employed; from which it may be concluded that none of the products of decomposition have escaped observation—that the liquid metal and the colorless gas, and nothing beyond these, went to make up the red powder. This opinion is confirmed by the fact that it is possible, under suitable conditions, to reproduce the red powder from oxygen and mercury. The process of resolving a com- pound into its constituents is known as analysis; that of building it up from its constituents is termed synthesis. Red oxide of mercury is therefore a compound, and its com- ponents are mercury and oxygen. Can these components be re- solved into still simpler bodies? The answer is, the resources of chemical science have not as yet been able to effect any such resolution. Both mercury and oxygen may be brought into union with various other bodies, and may be led by complicated processes from one combination to another; but at the end of their course they always emerge unchanged, and, if they do possess constituents, none of these have been dropped by the way. As no other kinds of matter can be extracted from them, it is agreed to regard them as elements. It is quite possible that the elements merely denote the present limits to our powers of effecting chemical decomposition. The only criterion which we have of the elementary nature of a body is, as above stated, the purely negative one of our inability to decompose it; and the history of the science shows us that this criterion is not necessarily trustworthy. The following is a list of the seventy elements at present recognized. The twenty-two most important of these are distin- guished by the largest type, those next in importance by medium type, whilst the names of elements which are either of rare occur- INORGANIC CHEMISTRY. rence, or of which our knowledge is very imperfect, are printed in small type: Name. Symbol.* Atomic weight.* Name. 1 Symbol. Atomic weight. ALUMINIUM A1 27 Nickel . . Ni 58.6 Antimony . . . Sb 120 Niobium . . Nb 94 Arsenic .... As 75 NITROGEN N 14 Barium .... Ba 137 Norwegium Ng 214 Beryllium . , . Be 9 Osmium . . Os 198.6 Bismuth . . . Bi 208.2 OXYGEN. O 16 Boron .... B 11 Palladium . Pd 105.7 BROMINE. . Br 80 PHOSPHORUS P 31 Cadmium . . . Cd 112 Platinum Pt 194.4 ' Caesium .... Cs 133 POTASSIUM . K 39 CALCIUM. . Ca 40 Rhodium . . Rh 104 CARBON . . C 12 Rubidium . . Rb 85.3 Cerium .... Ce 140.5 Ruthenium Ru 104 CHLORINE * Cl 35.5 Samarium . . Sm 150 Chromium . . . Cr 52 Scandium . . Sc 44 Cobalt .... Co 58.6 Selenium . . Se 79 COPPER . . Cu 63.2 SILICON . Si 28.2 Decipium, . . . Dp 159 SILVER . Ag 107.7 Didymiura . . . Di 146 SODIUM . Na 23 Erbium .... Er 165.9 Strontium . Sr 87.5 FLUORINE . E 19 SULPHUR S 32 Gallium .... Ga 68.8 Tantalum . . Ta 182 Gold Au 196 Tellurium . . Te 125 HYDROGEN H 1 Terbium . . Tb 148.8 . Indium .... In 113.4 Thallium . . T1 204 IODINE . . . 1 127 Thorium . . th 233.4 Iridium .... Ir 192.5 Tin .... Sn 118 IRON .... Fe 56 Titanium . . Ti 48 Lanthanum. . . La 138.5 Tungsten . . W 184 LEAD .... Pb 203.5 Uranium . . u 238.5 Lithium .... Li 7 Vanadium . . V 51.3 Magnesium . . Mg 24.4 Ytterbium . . Yb 172.8 MANGANESE Mn 55 Yttrium . . Y 89.8 MERCURY . Hg 200 ZINC . . . Zn 65.3 Molybdenum . . Mo 95.5 Zirconium . . Zr 90 It is usual to divide these elements into two great classes— metals and non-metals, the latter being sometimes also termed metalloids. The division is a somewhat arbitrary one, and the boundary-line between the two classes has been variously drawn by different chemists. Arsenic, selenium, and tellurium have been assigned to either category, according as the physical or the chemical characteristics formed the basis of the classification. Hydrogen, on the strength of its physical properties, is almost invariably classed as a non-metal; but its entire chemical beha- vior would lead to its being placed among the metals. An arrangement of the elements in their electro-chemical order, or a division into well-marked chemical groups, would perhaps be more logical. * For an explanation see Chapter VIII. CHEMICAL NOMENCLATURE. 39 CHAPTER 111. CHEMICAL NOMENCLATURE. The study of every science necessitates an acquaintance with the system of names and peculiar modes of expression which have been found most convenient to denote the materials and to describe the phenomena which form its objects. Such names and modes of expression constitute the groundwork of the language of every science, and upon the right employment of these depend the precision and accuracy of scientific definition. The nomenclature of a science ought to be distinguished by clearness and simplicity; but it is by no means easy to secure these conditions in a science like chemistry, where the rapid progress of discovery necessitates the continual addition of new and the fre- quent alteration of old names. The .chemical name of a substance should not only identify and individualize that substance, but it should also express the composition and constitution of the body, if a compound, to which it is applied. The first of these conditions is readily attained; but the second is much more difficult to secure, inasmuch as our ideas of the constitution of chemical compounds— of the mode in which they are built up as it were—require fre- quent modification- On this account all attempts to frame a perfectly consistent system of chemical nomenclature have hitherto been only partially successful. The names of the elements can scarcely be said to have been given according to any rule; many of them are derived from some prominent property of the bodies themselves, whilst others have a mythological origin. An attempt has been made to distinguish the metals by the termination urn, as potassium, sodium, etc.; but the common metals, such as gold, copper, and iron, still retain their original names; and one substance, selenium, which at the time of its discovery was regarded as a metal, has been suffered to retain its name unchanged, although further research has divested it of all metallic attributes. An important group of electro-negative* non-metals—flouriue, chlorine, bromine, and iodine—have received the termination ine; three are distinguished by the terminal syllable on, viz., carbon, silicon, and boron; and three others have gen for their final syllable, viz., oxygen, hydrogen, and nitrogen, these last names being derived from Greek words denoting the property possessed by these elements of generating respectively acid, water, and nitre. When two elementary bodies unite together, they form a chemical compound of the first order, to which the name binary compound has been applied. The names of these compounds are formed from those of their constituents, the name of the positive* * See Electrolysis, Chapter XII. 40 INORGANIC CHEMISTRY. constituent or some abbreviation thereof, with the termination ic, preceding that of the negative* constituent, which is made to ter- minate in ide, thus: Potassium and Sulphur form Potassic sulphide Sodium “ Oxygen “ Sodic oxide. Silver “ Chlorine “ Argentic chloride Zinc “ lodine “ Zincic iodide. Calcium “ Chlorine “ Calcic chloride. But the same elements frequently form with each other two compounds, in which case the one which contains the smaller pro- portion of the negative element is distinguished by changing the terminal syllable of the name of its positive constituent into ous, the terminal ic being retained for the compound containing the larger proportion of the negative element. Thus: One atom of tin and two atoms of chlorine form stannous chloride. One atom of tin and four atoms of chlorine form stannic chloride. Sometimes, however, the same elements form with each other more than two compounds. In these cases the prefixes hypo and per are employed as further marks of distinction; but their use is very rarely required. If a binary compound contains oxygen, and forms an acid when made to unite with water, or a salt when added to a base, it is termed an anhydride. Thus: One atom of carbon and two atoms of oxygen form carbonic anhydride. Two atoms of nitrogen and five atoms of oxygen form nitric anhydride. Two atoms of nitrogen and three atoms of oxygen form nitrous anhydride. One atom of sulphur and three atoms of oxygen form sulphuric anhydride. One atom of sulphur and two atoms of oxygen form sulphurous anhydride. In the following cases, the systematic names have not displaced the trivial and irregular names used for the same substances: Systematic name. Trivial or irregular name. Hydric oxide, . . . . Water. Hydric sulphide,. . . Sulphuretted hydrogen. Hydric selenide, . . . Seleniuretted hydrogen. Hydric telluride,. . Telluretted hydrogen. Hydric chloride, . . Hydrochloric acid. Hydric bromide, . . . Hydrobromic acid. Hydric iodide, . . . Hydriodic acid. Hydric fluoride, . . Hydrofluoric acid. Hydric carbide, . . f Marsh-gas or light carburetted * \ hydrogen. Hydric nitride, . . . Ammonia. Hydric phosphide, . . Phosphoretted hydrogen. Hydric arsenide, . . . Arseniuretted hydrogen. Hydric antimonide, . Antimoniuretted hydrogen. The term acid was originally applied only to substances possess- * See Electrolysis, Chapter XII. CHEMICAL NOMENCLATURE. 41 ing a sour taste like vinegar; but analogy has necessitated the application of the same name to a large number of compounds which have not this property. In the modern acceptation of the name, an acid may be defined as a compound containing one or more atoms of hydrogen, which become displaced by a metal when the latter is presentecl to the compound in the form of a hydrate. The hydrogen capable of being so displaced may be conveniently termed displaceable hydrogen. An acid containing one such atom of hydrogen is said to be monobasic, one containing two such atoms dibasic, etc. Acids of a basicity greater than unity are frequently termed polybasic acids. Thus nitric acid gives, with sodic hydrate, sodic nitrate: N OsH + ONaH = N03Na + OH2* Nitric acid. Sodic hydrate. Sodic nitrate. Water. Sulphuric acid gives, with potassic hydrate, potassic sulphate: S04H2 + 20KH = S04K2 + 20H2. Sulphuric acid. Potassic hydrate. Potassic sulphate. Water. And hydrochloric acid gives, with potassic hydrqte, potassic chloride: HCI + OKH = KCI + OH2. Hydrochloric Potassic Potassic Water, acid. hydrate. chloride. When an acid contains oxygen, its name is generally formed by adding the terminal ic either to the name of the element with which the oxygen is united, or to an abbreviation of that name; thus sulphur forms, with oxygen, sulphuric acid; nitrogen, nitric acid; and phosphorus, phosphoric acid. But it frequently happens that the same element forms two acids with oxygen; and when this occui's, the acid containing the larger amount of oxygen receives the terminal syllable ic, whilst that containing less oxygen is made to end in ous. Thus we have sulphurous acid, nitrous acid, and phosphorous acid, each containing a smaller proportion of oxygen than that necessary to form respectively sulphuric, nitric, and phosphoric acids. In some instances, however, the same element forms more than two acids with oxygen, in which case the two Greek words hypo, under, and hyper, over, are prefixed to the name of the acid. Thus an acid of sulphur containing less oxygen than sulphurous acid is termed hyposulphurous acid; and another acid of the same element containing, in proportion to sulphur, more oxygen than sulphurous acid and less than sulphuric, might be named either hypersul- phurous acid, or hyposulphuric acid; but the latter term has been adopted. The prefix per is frequently substituted for hyper; thus in the case of chlorine, which forms the following four acids with oxygen, viz., hypochlorous acid, chlorous acid, chloric acid, * For an explanation of these formulae see Chapter YIII. 42 INORGANIC CHEMISTRY. and by perchloric acid, the latter is generally named perchloric acid; but per can only be used as a prefix to the acid containing the largest proportion of oxygen. Some acids do not contain oxygen amongst their constituents, but consist of sulphur or hydrogen united with other elements. This peculiarity of composition is expressed in their nomenclature by the prefixes sulpho or sulph (or the equivalent Greek prefixes thio or thi), and hydro or hydr: thus sulpharsenic acid and sulpho- stannic acid denote acids composed respectively of sulphur, hydro- gen, and arsenic; and sulphur, hydrogen, and tin; whilst the names hydrochloric acid and hydriodic acid are given to acids composed, the first of hydrogen and chlorine, and the second of hydrogen and iodine. The terminals ous and ic are also applied to these acids in exactly the same manner as to the oxygen acids; thus we have sulpharsenious and sulpharsenic acid, the latter con- taining a larger proportion of sulphur than the former; but the application of the first of these terminals has not hitherto been found necessary in the case of hydrogen acids, since no element has yet been observed to form more than one acid with hydrogen. The term anhydride (cf. p. 40) is applied to the residue obtained by the abstraction (in combination with oxygen as water) of all the displaceable hydrogen from one or two molecules of an oxygen acid. Thus, SG4H2 0H2 = S03; Sulphuric acid. Water. Sulphuric anhydride. 2N03H OH2 = N205. Nitric acid. Water. Nitric anhydride. The term anhydro-add or pyro-add is applied to such acids as are formed from two molecules of a polybasic acid (see p. 41) by elimination of water: 2POJT3 - OH2 = P207H4; Phosphoric acid. Water. Pyrophosphoric acid, 2S04H2 OH2 = S207H2. Sulphuric acid. Water. Pyrosulphuric acid. (Nordhausen sulphuric acid.) These acids are thus partial anhydrides. The prefix pyro originally referred to their mode of formation, heat being employed to drive off the water; but its use has been extended to acids which have been prepared by other means, and it is to be under- stood generally as denoting partial anhydricity between two molecules of the parent acid.* * This sense of the prefix pyro must not be confounded with that in which it is employed in organic chemistry, as in pyrotartaric acid, pyromucic acid, etc. Here the mode of formation by the action of heat is alone indicated, the compounds having for the most part nothing further in common, and not being formed from the parent acid—tartaric acid, mucic acid—according to any fixed rule. CHEMICAL NOMENCLATURE. 43 The term base is applied to three classes of compounds, all of which are converted into salts by the action of acids. These are: Ist. Certain compounds of metals with oxygen, such as sodic oxide (Na2o), zincic oxide (ZnO), etc. 2d. Certain compounds of metals with the compound radical hydroxyl (HO), such as sodic hydrate (Na(HO)), zincic hydrate (Zn(HO)3), etc. 3d. Certain compounds of nitrogen, phosphorus, arsenic, and anti- mony, such as ammonia (NH3). There are also a few organic compounds to which the name base is sometimes given, but which are not included in the above classes; it is, however, unnecessary further to allude to them here. The bases of the first class are named in accordance with the rules already given for compounds of two elements. The following bases, however, still retain their irregular names: Systematic names. Irregular names. Baric oxide, Strontic oxide, Calcic oxide, Magnesic oxide, Alurainic oxide, Beryl!ic (Glucinic) oxide, . . . Beryl lia (Glucina). Zirconic oxide, The names of the bases belonging to the second class are formed by changing the terminal syllable of the name of the metal into ic or ous, and the word hydroxyl into hydrate. Thus caesium and hydroxyl from csesic hydrate (Cs(HO)); barium and hydroxyl, baric hydrate (Ba(HO)2); and iron and hydroxyl, ferric hydrate (Fe2(HO)6). A few of these bases have trivial or irregular names, which are almost in variably used instead of the systematic names : Systematic names. Irregular names. Potassic hydrate, . . . Potash. Sodic hydrate, . . . . . Soda. Litliic hydrate, . . . . Lithia. The bases of the third class are distinguished by the terminal syllable me, except nitrine (]SrIT3), which retains its trivial name ammonia. These bases belong almost exclusively to the depart- ment of organic chemistry, and their nomenclature could not be advantageously discussed here. It has been already mentioned that by the mutual action of an acid and a base upon each other, a salt is produced. If the salt be free from oxygen and sulphur, like common salt (NaCl), it is termed a haloid salt; if it contain oxygen it is termed an oxysalt; 44 INORGANIC CHEMISTRY. and if this oxygen be replaced by sulphur, it is distinguished as a sulphosalt. The haloid salts are named according to the rules of binary compounds above given, thus: Name. Formula. Sodic chloride, . NaCI. Calcic iodide,. Ferrous bromide, . Cal2. . FeBr2. Ferric bromide, . . Fe2Br6. Oxysalts are divided into normal, acid, and basic. A normal salt is one in which the displaceable hydrogen of the acid (see p. 41) is all exchanged for an equivalent amount of a metal or of a positive compound radical. The following examples will serve to illustrate this definition of a normal, or as it is sometimes incorrectly called, a neutral salt, the displaceable atoms of hydrogen in the acid, and the metal by which they have been displaced in the salt, being printed in italics: Acid. Normal salt. »>*■ phosphoric acid, Kw.. : fpofe.",, Hypophosphorons acid, . . P02H27T, . Sodic hypophosphite, . P02H2iVa. Phosphorous acid, .... P03Hj!J2, . Potassic phosphite, . . P03HiT2. Metaphosphoric acid, . . . P03H, . . Lithic metaphospliate, . P03ih’. Pyrophosphoric acid, . . . P207H4, . Calcic pyrophosphate, . P207(7a//2. Nordhausen sulphuric acid, . S./J.JE,, . Sodic pyrosulphate,. » S207jVa2. Unknown acid, Cr2()7JJ2, . Potassic dichromate, . Cr207/C2. An acid salt is one in which the displaceable hydrogen of the acid is only partially exchanged for a metal or positive compound radical. The following examples illustrate the constitution and nomen- clature of these salts: Acid. Acid salt. Sulphnric acid, . . S04H2, Hydric sodic sulphate, . . . SO4HNa. Carbonic acid, . . CO,//, ? Hydric potassic carbonate, . CO JTK. (Hydric disodic phosphate, . V04IINa2. Dihydric sodic phosphate, . POiHiNa. Microcoamic salt, .... P(J4 7/( (Hydric ammonic sodic phosphate.) Acid salts are produced almost exclusively from polybasic acids. When the number of bonds* of the metal or compound positive radical contained in a salt exceeds the number of atoms of displace- able hydrogen in the acid, the compound is usually termed a basic salt—as, for instance: * For an explanation of this term see Chap. VIII. LAWS OF COMBINATION. 45 Acid. Basic salt. r\ i • . i tt f Malachite, •••••* CO^llnOw^o. ar ’ ‘ 3 2’ (Bine cupric carbonate, . . C 20811.2Cfu//3. Sulnlmric acid SO H J Tribasic cuPric sulphate, . S08H40m//3. oulpliunc acid, . . SU4ii2, j Turpeth minerai( .... SO These and most, if not all, other basic salts do not differ essentially in their constitution from the normal and acid salts. This will be seen from the arrangement of their atoms given under the heading of the different metals entering into their composition. The molecular compounds (q.v.) which various substances form with water of crystallization may be conveniently termed aquafes. The nomenclature of organic bodies is founded upon the same principles as that of inorganic compounds; but its discussion could not be conveniently introduced here. CHAPTER IV. LAWS OF COMBINATION. As soon as chemists began to realize that the various changes which matter undergoes when two or more substances are extracted from some other substance, or unite to form this substance, are not changes in the ultimate nature of matter itself, but only in its mode of combination, it was natural that they should have recourse to the balance in order to determine the quantities of the different kinds of matter entering into each such combination. The results of these determinations are embodied in the following numerical laws, which form the groundwork of the science. Law of Constant Proportions.—lt has already been mentioned that each chemical substance contains its elements always in the same fixed proportions. Red oxide of mercury consists of 12.5 parts by weight of mercury and 1 of oxygen, this proportion being absolutely unvarying. In like manner hydrochloric acid gas always contains 35.5 parts of chlorine to 1 of hydrogen. And in the same propor- tions in which the elements of a compound may be separated from each other by analysis, they may by synthesis be made to combine. An excess of any one of the elements over and above the quantity required to unite with the rest, will remain unacted upon. If 40 parts of chlorine be brought into contact with 1 part of hydrogen under the conditions which are necessary for the formation of hydro- chloric acid, 4.5 parts of chlorine will remain unchanged, and cannot be made to enter into combination. The above law is known as the Law of Constant Proportions. It was in the course of the experimental development of this law that the great fact first became clear, that matter is indestructible, and, as far as experience goes, uncreatable. When carbon is burnt in a vessel containing oxygen it seems to disappear; but if nothing be allowed to escape, and if the vessel be accurately weighed both before 46 INORGANIC CHEMISTRY. and after the combustion, the weight will be found not to have changed. The carbon has merely combined with the oxygen to form the invisible gas carbonic anhydride. If a burning piece of the metal sodium be now plunged into the carbonic anhydride thus formed, the sodium will combine with the oxygen of the carbonic anhydride, and the carbon will reappear as a fine black dust. In every series of chemical processes, however complicated, the sum of the weights of the final products will be neither more nor less than that of the initial substances. Law of Multiple Proportions.—ln the course of their quan- titative researches, chemists found that in some cases the same two elements combined with each other in two or more different propor- tions, to form totally distinct compounds; but as these proportions were always constant for each such compound, this new fact did not in any way contradict the law just stated. A very simple numerical relation regulates this variation. Mercury, for example, forms two compounds with oxygen—the red oxide, in which the proportion of mercury to oxygen is as 12.5 : 1; and a black oxide, in which the proportion is as 25 :1. The mercury in the first compound is, there- fore, to that of the second as 1: 2. With nitrogen, oxygen forms no fewer than five different compounds : Nitrons oxide, Parts by weight of nit rogen. 1 Parts by weight of oxygen. 0.571 N itric oxide,. 1 1.142 Nitrous anhydride, 1 1.714 Nitric peroxide, Nitric anhydride, . 1 2.285 1 2.857 The relative proportions of the oxygen uniting with a constant weight of nitrogen in these five compounds are as 1:2:3 : 4 : 5. In all cases in which one element unites with another in two or more different pro- portions these proportions are found to be simple multiples of some common factor. This law is known as the Law of Multiple Proportions. Law of Equivalent Proportions.—The foregoing numerical law was discovered by comparing the different weights of the same element which combine with a given weight of some other element. But when the weights of different elements which combine with a given weight of various other elements were compared, new and sur- prising numerical relations became manifest. Thus— 1 part of chlorine 1 part of bromine 1 part of iodine 1 part of oxygen 1 part of sulphur Combines with Hydrogen,. . . 0.02817 0.0125 0.00787 0.125 0.0625 Sodium,.... 0.6479 0.2875 0.1811 2.875 1.4375 Potassium,. . . 1.099 0.4875 0.3071 4.875 2.4375 Copper, .... 0.891 0-395 0.249 ' 3.95 1.975 Lead, .... 2.908 1.2906 0.813 12.906 6.453 47 LAWS OP COMBINATION. The numbers in each vertical column bear to each other the same proportion ; thus, in all the columns— Hydrogen : Sodium : Potassium : Copper : Lead. as . . 1 : 23 : 39 : 31.6 : 103.25 It will be noticed that the numbers for hydrogen, sodium, and potassium are the same as those attached to these elements in the column headed “Atomic weight” in the table of elements, p. 38, whilst those for copper and lead are less by one-half than the num- bers in the table. The reason of this will be explained later. (See Chapter XII., Electro-chemical Equivalents.) On the other hand— 1 part of hydrogen 1 part of sodium 1 part of potassium 1 part of copper 1 part of lead Combines with Chlorine, . . . 35.5 1.544 0.91 1.123 0.343 Bromine, . . . 80 3.478 2.05 2.531 0,774 Iodine, .... 127 5.522 3.256 4.019 1.229 Oxygen, . . . 8 0 348 0.205 0.253 0.0774 Sulphur, . . . 16 0.696 0.41 0.506 0.1548 Here again, in all the vertical columns— O / Chlorine : Bromine : lodine : Oxygen : Sulphur. as . . 35.5 : 80 : 127 : 8 : 16 The numbers which express the proportions of chlorine-, bromine, and iodine are those given in the table on p. 38 ; whilst those of oxygen and sulphur are less by one-half. This law may be expressed thus : The relative proportions by weight in which the members of any series of elements combine with the same quantity of another element are the same for their combina- tions with any other element. 35.5 parts by weight of chlorine, 80 parts by weight of bromine, 127 parts by weight of iodine, 8 parts by weight of oxygen, and 16 parts by weight of sulphur are said to be equivalent, as each of these weights serves to satisfy the chemical affinity of 1 part by weight of hydrogen. In like manner 1 part by weight of hydrogen, 23 parts by weight of sodium, 39 parts by weight of potassium, 31,75 parts by weight of copper, and 103.5 parts by weight of lead are equiva- lent. But the members of the first series are also equivalent to those of the second : thus 23 parts by weight of sodium combine with 35.5 parts by weight of chlorine, 39 parts by weight of potassium with 80 parts by weight of bromine, etc., as may easily be calculated from the last table. Thus every element may have an equivalent weight as- signed to it, according to which it combines with other elements, the equivalent weight of hydrogen being taken as unity. 48 INORGANIC CHEMISTRY. CHAPTER Y. THE ATOMIC THEORY. In order to account for the remarkable relations just described, chemists have adopted a theory concerning the ultimate constitution of matter which is to be found in the systems of some of the ancient Greek philosophers, but which first received a scientific form at the hands of Dalton. Dalton supposed matter to consist of exceedingly minute particles, incapable of further division—atoms from a privative, and ri,avw} I cut). These atoms possess different weights in the different kinds of elementary matter, but have always the same weight for the same kind. The juxtaposition of different elementary atoms constitutes chemical combination. Thus if the relative weights of the atoms of potassium and chlorine are as 39 to 35.5, and if the formation of potassic chloride consists in the juxtaposition of one atom of the one element to one of the other, then it is evident that potassic chloride can contain its elements only in the proportion of 39 parts by weight of potassium to 35.5 parts by weight of chlorine. If the relative weights of the atoms of mercury and oxygen are as 200 to 16, and if red oxide of mercury is a combination of one atom of each of its elements, it must contain mercury and oxygen in the pro- portion of 200 to 16. Again, if the black oxide of mercury is a combination of two atoms of mercury with one of oxygen, the pro- portion of the former to the latter must be as 400 to 16, or the pro- portion of mercury in the black oxide is to that in the red as 2 to 1 for equal weights of oxygen. Thus by the hypothesis of atoms, which possess the same weights for the same elementary kind, but different weights for the different elementary kinds of matter, the three great experimental facts of Constant Proportion, Multiple Proportion, and Equivalent Propor- tion are referred to one general law. The atomic theory has, since its adoption by Dalton, undergone many developments, particularly in the sharp distinction of atoms from molecules {molecula, diminutive of moles, a mass). The atoms which enter into chemical combination are supposed to be grouped into molecules—“ little masses.” These latter are again grouped together to form the masses of matter recognizable by the senses. Thus a solid piece of ice, which contains the atomic weights of hy- drogen and oxygen in the proportion of 2 to 1, is not to be regarded as having its atoms thrown together indiscriminately ; it is supposed to be made up of a vast number of small independent systems, each containing two atoms of hydrogen and one of oxygen. The atoms within the molecule are held together by chemical attraction: the molecules are kept in their places by cohesion. Neither the atoms within the molecule nor the molecules within the mass, are supposed to be in actual contact. When a body expands by heat the distance THE ATOMIC THEORY. 49 between its molecules is increased, and when it contracts by cooling this distance is diminished. Neither the atoms nor the molecules in a solid body are to be conceived as occupying their positions in a state of rest; various considerations, chiefly of a physical nature, lead to the conclusion that they execute some sort of vibratory motion about their positions of equilibrium. The amplitude of vibration increases with the temperature. If the amplitude of vibration of the molecules becomes too great for stability, the molecules detach themselves from their positions of equilibrium, desert the immediate sphere of attrac- tion of the neighboring molecules, and wander about till they fall under the dominion of other molecules, to be again released by their intensity of vibration. This state of things corresponds to liquidity: cohesion is alternately overcome and restored, and hence is weakened. If, however, the energy of the molecules becomes so great as to carry them beyond the reach of their mutual attraction, they shoot forward in straight lines until they strike against other molecules or against the sides of the containing vessel, in which case they rebound and change their direction, sometimes imparting, sometimes receiving energy. This represents the gaseous condition of matter. Up to this point the atoms which compose the molecule have been considered as keeping together during the wanderings of the molecule itself; but if the temperature be raised still higher, it may happen that the vibra- tion of the atoms within the molecule will carry these also beyond the reach of their mutual attraction, in which case some of them may separate from the parent molecule, forming among themselves simpler molecules more capable of existing at a high temperature. This is the phenomenon of decomposition by heat. It is probable that, at sufficiently high temperatures, only elementary matter can exist, and it is possible that even the molecules of the elements (for, as will be shown later, the atoms of the same element combine with each other to form molecules) break up into their component atoms. (See lodine.) The motions of the molecules are manifested in the phenomena of the diffusion of liquids and gases. In order to give some conception of the aims and scope of the atomic theory in its most recent developments, it may be mentioned that modern chemistry seeks to determine not only the nature and number of the atoms in the molecule, but also their arrangement. That there must be a special arrangement is shown by the fact that two or even more totally distinct compounds may exist having the same number of the same atoms in the molecule. Such compounds are termed isomeric. The molecule is to be looked upon as a system composed of various members held together by chemical attraction, just as the members of one of the cosmical systems are held together by gravitation. The molecule of acetic acid, for example, contains two atoms of carbon, four of hydrogen, and two of oxygen. To continue the astronomical illustration, the two atoms of carbon are supposed to be united by mutual attraction like the two suns of a double star. One of these suns possesses three planets in the shape of three atoms of hydrogen ; the other has two atoms of oxygen as planets; whilst one 50 INORGANIC CHEMISTRY. of the oxygen planets has an atom of hydrogen annexed to it as a satellite. Of course all the members of such a system must attract each other; but the attraction will be greatest between those which, cceteris paribus, are by virtue of their position most subject to each other’s influence. When the molecule is divided at any point, the two parts, provided the reaction by which the separation has been effected is not too violent, retain their previous arrangement : thus, by heating potassic acetate with caustic alkali, it is possible to divide the molecule of acetic acid at the junction of the two carbon atoms, in which case the one carbon atom retains its three hydrogen atoms, and the other its two oxygen atoms—one of these with an atom of potassium in the place of the hydrogen of acetic acid. In like man- ner, by the action of phosphorous chloride, the molecule of acetic acid may be divided so as to split off the atom of oxygen with its hydrogen atom attached. Both parts again remain unchanged as regards their internal arrangement.* The facts on which these assertions are based could not with ad- vantage be introduced into this chapter. They will be fully treated of in their proper place. To the unscientific mind there is something peculiarly repellent in the atomic theory and in the physical conceptions which it involves. Our notions of a multitude of minute unconnected particles are de- rived from the sand-heap—the symbol of instability—and to realize that a solid mass, such as an ingot of steel, consists of minute particles suspended in space without actual contact, is certainly at first sight difficult. But the student of science must dismiss from his mind all crude analogies, and learn above all things to distrust bis unaided senses, which in scientific matters are by no meaqs so infallible as they are considered to be in everyday life. In transmitting to the mind the phenomena of the external world, the senses first translate these phenomena into a language of their own, which, however ad- mirably adapted for its purpose, is only a symbolical representation of the phenomena themselves. Sound as heard by the ear has no resemblance to the vibrations of the air; red and violet light as they affect the eye are in no way like longer and shorter waves of ether: yet this is what science tells us concerning these phenomena as they exist outside the sentient subject. And the same holds of the other forces of nature. But the object of science is to perceive the pheno- mena as they are in themselves—stripped of the interpretation put upon them by the senses. Hence it is that many of the greatest discoveries have apparently contradicted the evidence of the senses. The magnificent generalization of the conservation of energy, a pendant to that of the indestructibility of matter, has given to the dynamical sciences a unity which they formerly lacked, and has laid down the lines of their future progress. Just as, when we have led an element through a series of combinations with other elements and * Lucretius (De Rerum Natura) has a remarkable passage, which might almost be regarded as an anticipation of the views of modern chemists regarding the con- stitution of compounds. “It matters much,” he says, “ with what others and in what position the same atoms are held together.” THE ATOMIC THEORY. 51 find that the increase of weight due to the accession of this element has in all cases been the same, and that we can extract the origi- nal quantity of the first element, unaltered in all its properties, from its last combination, we conclude that these various compounds, in spite of the difference of their characteristics, all actually contained this given quantity of the same kind of matter; so, when we trans- form the motion of a mass of matter into the various other forms of energy and find that the quantities are in every case equivalent, and that each of these equivalent quantities can (or could, were it possible to operate without loss) be transformed back into the original quan- tity of motion of matter, we conclude that all these manifestations of energy actually consisted of the same thing—motion of matter. When the motion of a mass is suddenly arrested, this motion is converted into heat—a motion of the molecules. And in all cases of convertible forms of energy, the amount of this energy, as expressed in terras of the masses and of the velocities, will be the same, whether the masses be sensible masses, or whether they be molecules. A further refinement of speculation as to the nature of atoms has been introduced by Sir William Thomson in the hypothesis that the ultimate atoms of the elements consist of various forms of vortex rings in a perfect fluid, the ether. This would reduce the different kinds of matter to varieties of motion in one kind of matter, and would account among other things for the indestructibility of matter; it being mathematically demonstrable that a vortex ring in a perfect fluid is indestructible. But it is not necessary in a work like the pres- ent to do more than refer to this hypothesis. Fascinating as all these speculations are, they must never be taken at more than their true value. Even the atomic theory, which ex- plains perhaps as many heterogeneous facts as any other theory, not excepting that of gravitation and the undulatory theory of light— these two theories surpassing it, however, in the important point of their far higher mathematical development—must not be looked upon as more than the best existing explanation of the facts as at present known. It may represent the absolute truth; it may be nothing more than a symbolical expression of certain aspects of the truth. The real object of a theory is to group the facts round some central idea from which we may start in our search for fresh facts. The deductions from the theory are the objects of experiment, and by ex- periment the theory stands or falls. The greater the number of new facts a theory predicts, the better is the theory; but that is all that can be said of it. No number of verified predictions can establish the absolute truth of a theory. Of course this does not refer to those particular cases in which the theory itself may be an ultimately veri- fiable matter of fact. It can scarcely be so with the atomic theory. No one has ever seen an atom or a molecule, and from theoretical considerations derived from the undulatory theory of light, it is almost certain that no one ever will. The opposed conception is, that matter fills space continuously and homogeneously. It is impossible to review here the vast array of physical evidence which speaks against this conception and in favor 52 INORGANIC CHEMISTRY. of the atomic theory : the chemical evidence forms the subject of this work. One chemical fact may, however, be specially mentioned at this point. It has already been stated that the same quantities of the same kinds of matter frequently combine so as to produce two or more totally different compounds. With matter homogeneously fill- ing space this would be inconceivable. Such a difference bespeaks, as was said before, an arrangement of parts. Furthermore, as in the state of the finest mechanical subdivision the particles of a chemical compound all display the same qualities, the parts, by the juxtaposi- tion and arrangement of which the compound is produced, must be exceedingly small. We are thus led back to the atomic theory. How small the ultimate parts of matter are supposed to be may be judged from Sir William Thomson’s calculation that in solids and liquids the mean distance between the centres of contiguous molecules is less than t4eowotto and greater than 46oA“o"ooAo °f a centimetre. The molecular vibrations, to which reference lias already been made, must of course take place through a correspondingly small range. CHAPTER YI. MOLECULAR WEIGHTS. All bodies in the gaseous state are affected equally by pressure. If a given volume of hydrogen and a given volume of chlorine be measured at the pressure of one atmosphere, and if the pressure in each case be then doubled, it will be found that the volume of each has been reduced by one-half. If, on the other hand, the pressure be reduced to half an atmosphere, the original volume of each will be doubled. This relation is expressed by saying that the volume of a gas is, cceteris 'paribus, inversely proportional to the pressure under which it is measured. This law is named from its discoverer Boyle’s Law. Exceptions to it occur in the case of gases and vapors in the neighborhood of their point of condensation to liquids, when the gaseous condition is imperfect. In these the volume decreases more rapidly than the pressure increases. In like manner, all bodies in the gaseous state are affected equally by change of temperature. Every gas, when measured at 0° C., expands of its original volume when heated to 1° C., supposing the pressure to remain constant during the operation. This fraction is called the co-efficient of expansion of gases. The dilatation takes place in the same ratio for every further increase of temperature: thus if the volume of a gas at 0° be equal to 1, the volume at t° will be 1 T This might also be expressed by saying that, the pressure being constant, the volume of a gas is proportional to its temperature measured from 273°. Thus the volume of a gas at 20° is to its 53 MOLECULAR WEIGHTS. volume at 70° as 273 +2O : 273 -f- 70. This law holds for all gases, subject to the deviations mentioned in the ease of Boyle’s Law. The relation of the volume of gases to temperature was discovered by Charles. The kinetic theory of gases, a theory at present almost universally accepted by physicists, explains the elasticity and pressure of a gas as the result of the shock of its molecules against the sides of the vessel in which it is contained. If the volume of the gas be reduced by one-half, the number of molecules which strike against the unit of surface in unit of time will be doubled ; and hence the pressure will be doubled. If the temperature be raised, the velocity of the molecules, and hence their energy, will be increased : the shock against the sides of the vessel is more intense and also more frequent, hence the pres- sure will be greater. All gases behave in exactly the same manner in regard to temperature and pressure, and the only satisfactory ex- planation of this uniformity is the assumption that equal volumes of all gases at the same temperature and pressure contain an equal num- ber of molecules. In fact this assumption has been deduced as a law by strict mathematical processes from the kinetic theory of gases.* This law was first stated as a hypothesis by Avogadro in 1811. It excited little attention at the time, but is now one of the chief foun- dations of modern chemical theory. As equal volumes of all gases contain equal numbers of molecules, it is evident that the molecular weights of gaseous bodies will be pro- portional to the weights of equal volumes at the same temperature and pressure, i.e., to their specific gravities or vapor-densities. If the molecular weight of hydrogen, as the lightest known gas, were to be taken as unity, the molecular weights of other gases would be ex- pressed by the number of times that their specific gravity is greater than that of hydrogen. As will be shown later, however, the mole- cule of hydrogen consists of two atoms. Since, therefore, its atomic weight is taken to be equal to 1, its molecular weight will be 2. Let the unknown molecular weight of a gas be M, and let its specific gravity (referred to that of air as unity), as found by experiment, be d, then since the specific gravity of hydrogen is 0.0693: 0.0693 : 2 = d:M and M = 28.86 d, or, expressed in words, the molecular weight of a gas may be found by multiplying its specific gravity (referred to that of air as unity) by 28.86. From what has been said above, it is evident that the terra gas will here include the vapors of all substances, solid or liquid, capable of volatilizing without decomposition. it; on the other hand, the specific gravity of the gaseous body is referred to that of hydrogen as unity, then, calling this specific gravity L>, we should have * See Clerk Maxwell, Theory of Heat, 3d edition, p. 296. 54 INORGANIC CHEMISTRY. 1:2 = D: M or M= 2D. That is to say, the molecular weight of a substance is found by doubling its specific gravity in the gaseous state, the specific gravity of hydrogen being taken as unity. It is evident that the molecular weight will be equal to the sum of the atomic weights of all the atoms contained in the molecule. (See Atomic Weights.) Since in nearly every case of chemical action between two or more substances, it is the molecules of these substances which act on each other—either by exchange of atoms or by direct union—and since equal volumes of gas contain, cceteris paribus, equal numbers of mole- cules, it might be expected that in chemical action between gaseous bodies the volumes entering into reaction would present some simple relation to each other. Not only is this the case, but the gaseous volume of the product of the reaction also follows a very simple law. Thus: 1 vol. of hydrogen +1 vol. 1 “ “ +1 “ of chlorine yield 2 vols. of hydrochloric acid. “ “ hydrobroraie acid. bromine vapor 2 vols. “ +1 “ sulphur vapor “ “ sulphuretted hydrogen. 2 “ “ +1 “ oxygen “ steam. 3 “ “ +1 “ nitrogen “ “ ammonia. The law of combination by volume was discovered by Gay-Lussac. If the number of molecules in one volume be called n, the first of the above combinations might be written thus: n molecules of hy- drogen combine with n molecules of chlorine to form 2n molecules of hydrochloric acid. As each of the 2n molecules of hydrochloric acid contains both hydrogen and chlorine, each of the n molecules of hydrogen and each of the n molecules of chlorine must have been divided into two parts in order to furnish hydrogen and chlorine for these 2n molecules. The molecule of hydrogen therefore consists of at least two atoms of hydrogen. The molecule of chlorine is likewise at least diatomic. Reasons will be given latter for the belief that the number of atoms in the molecules of these elements is not greater than two.* The combination by volume may therefore be written : 2n atoms of hydrogen combine with 2n atoms of chlorine to form 2n molecules of hydrochloric acid ; or, dividing by 2a,: 1 atom of hydrogen com- bines with 1 atom of chlorine to form 1 molecule of hydrochloric * The supposition that the molecules of the great majority of the elements con- sist of mutually combined elementary atoms, throws light upon a number of other- wise inexplicable phenomena. Thus elements in the so-called nascent state—that is, at the moment at which they are released from their combinations—display much more powerful affinities, and are much more capable of effecting chemical changes than when in the free state. The explanation is that in the nascent state, it is the single atoms which are released from combination, and that being endowed with free affinities they are especially ready to enter into any fresh combination ; whereas in the case of the free element, the atoms have combined with each other to form molecules: not only therefore have the atoms no longer any free affinities, but their mutual combination has to be broken up before they can enter into union with other elements. MOLECULAR WEIGHTS. 55 acid. That is to say, if we represent in this case the atomic propor- tion of each of the combining elements by one volume, the molecular proportion of the resulting compound will be represented by two volumes. The same holds of all the combinations given in the above list; thus we may write: 2 atomic proportions (or volumes) of hy- drogen combine with 1 atomic proportion (or volume) of oxygen to form 1 molecular proportion ( = 2 volumes) of steam. This is what is meant by the elliptical and somewhat misleading expression frequently employed, that the molecule of a compound occupies in the gaseous state two volumes. In every case, if we take such proportions by volume of the gaseous elements as will represent the atomic proportions* of these elements uniting to form a compound, the molecular proportion of this compound, if measured in the gase- ous state, will occupy two volumes. Further, as equal volumes of all gaseous substances contain an equal number of molecules, it is evident that the molecular proportion of these various combining elements will also be represented in the gaseous state by two volumes. But though the molecular proportion may in every case be represented by two volumes, it by no means follows that the atomic proportion of the gaseous elements may always be repre- sented by one volume, though this happens to be the case in the series of combinations given in the foregoing list. In order to ascertain what volume of a gaseous element corresponds to its atomic proportion when the molecular proportion is represented by two vol- umes, it is necessary first to ascertain how many atoms the molecule of that element contains. This may be found by dividing the mo- lecular weight, as deduced from the vapor-density, by the atomic weight, as determined by one or more of the methods given in the next chapter. Name of element. Molecular weight. Atomic weight. Number of atoms in molecule. Mercury, I 200 200 1 Cadmium, 112 112 1 Zinc, 65 65 1 Hydrogen, 2 1 2 Oxygen, 32 16 2 “ (as ozone), . . . 48 16 3 Chlorine, 71 35.5 2 Bromine, 160 80 2 Iodine, 254 127 2 Nitrogen, 28 14 2 Sulphur (at 524°), . . . 192 32 6 “ (at 860°), . . . 64 32 2 Selenium, 158 79 2 Tellurium, 256 128 2 Phosphorus, 124 31 4 Arsenic, 300 75 4f * See following paragraph. t This list contains all the elements of which the vapor-density has been deter- mined, and,consequently, all the elements of which the molecular weight is known ; for though other methods of ascertaining the molecular weight will be described, 56 INORGANIC CHEMISTRY. The number of atoms contained in the molecules of the various elements is therefore not always the same.* Thusin the case of mercury, cadmium, and zinc, the molecular weight is identical with the atomic weight: the molecules of these elements are monatomic. With hydrogen, oxygen, chlorine, nitrogen, and various other ele- ments, the molecular weight is twice as great as the atomic weight: the molecules are diatomic. In oxygen in the form of ozone, on the other hand, the molecule is triatomic. Phosphorus and arsenic are examples of tetratomic molecules, while the molecule of sulphur is hexotomic at 524°, and diatomic at 860°, the heavy hexatomic mole- cule breaking up into three lighter diatomic molecules as the tem- perature rises. Kundt and Warburg, by a determination of the velocity of sound in mercury vapor, have shown that in the case of this vapor there is no increase of “specific heat at constant volume” due to motion of atoms within the molecule, as is the case with gases having molecules containing more than one atom. The molecule of mercury in the gaseous state must therefore be assumed to be truly monatomic. From this it follows that diatomic molecules really contain only two atoms, triatomic molecules only three atoms, etc. It is evident that, whatever volume of a gas is adopted to repre- sent its molecular proportion, the volume required to represent its atomic proportion will be inversely as the number of atoms in the molecule of that gas. If, therefore, the molecular proportion is represented by two volumes, the volume corresponding to the atomic proportion will be found by dividing this molecular volume by the number of atoms in the molecule. Thus we find that for a monat- omic gas, the volume representing one atomic proportion—or, as it may be termed, the atomic volume—is two volumes; for a diatomic gas one volume; fora tetratomic gas, half a volume, and so on. A very convenient expression of these relations is afforded by a notation devised by A. W. Hofmann. In this notation one volume of an element in the gaseous state is represented by a square D within which is written the symbol of the element in question, the atomic volume of this element being unity; two volumes by a double square, open in the middle ; and half a volume by a tri- angle . Thus in the case of the elements, these symbols would be employed as follows : only that of vapor-densities is applicable in the case of elements. All other elements are either non-volatile or volatilize at temperatures and under conditions such as to render the determination of their density in the gaseous state a problem beyond the present resources of chemistry. Silver, for example, is volatile only at the temper- ature of the oxyhydrogen flame. Again, potassium and sodium, though volatile at relatively low temperatures, yield vapors which attack and combine with the mate- rial of the vessels employed, and in this way furnish discrepant and untrustworthy results. Hence the molecular weight of all elements other than those contained in the above table is at the present moment purely a matter of surmise. * From this it follows that the vapor-density alone of an element furnishes no clue to its atomic weight. MOLECULAR WEIGHTS. 57 Atomic volume in the Molecular volume in Name of element. gaseous state. the gaseous state. Mercury, i Hg 1 Hg i Cadmium, . ... ! Cd i Cd i Zinc, i Zn i 1 Zn, i Hydrogen, H 1 Ha i * Oxygen, . . . . ... . O Oa .... 1 Chlorine, ...... Cl a2 i Bromine, Br i Bra l Iodine, I i h , i Nitrogen, N n2 i Sulphur (at 860°), 8 i $z I Selenium, . .'».•■* . Se i f Sea I Tellurium, . . . . . Te Tea i Phosphorus, .... P\ i i Arsenic, As\ i 1. In the ease of compounds, the symbol of the compound (see Chemical Notation) is written within the double square representing the molecular volume in the gaseous state, thus: Molecular volume in the Name of compound. gaseous state. Hydrochloric acid, HCl i Water, ok2 i Ammonia, etc., nh3 1 These volume-symbols may be combined into equations (see Chemical Notation), which will thus express the relative volumes * The small subscript Arabic numeral indicates how many atoms of the element represented by the atomic symbol are present (see Chemical Notation). 58 IXORGANIC CHEMISTRY. of the gaseous elements or compounds taking part in any chemical action, and the volume of the resulting product or products. Thus; 1 H Cl HQ l_ I or one volume of hydrogen combines with one volume of chlorine to form two volumes of hydrochloric acid. "=1 T —{ + o = oh2 H J or two volumes of hydrogen combine with one volume of oxygen to form two volumes of steam. H h| + m = f nh3 H or three volumes of hydrogen combine with one volume of nitrogen to form two volumes of ammonia. T R T Hg + r=\ 1 a 1 or two volumes of mercury vapor combine with two volumes of chlorine to form two volumes of the vapor of mercuric chloride. a + a = pcis a or half a volume of phosphorus vapor combines with three volumes of chlorine to form two volumes of the vapor of phosphorus chloride. Of course in reality these chemical reactions take place not between atoms, but between molecules, and the reaction of hydrogen with chlorine, for example, would therefore have to be written: 5 + aa I = h'q + iicl j- J I . J i- 1 1 MOLECULAR WEIGHTS. 59 but the above simplified mode of expression has been adopted in order that the molecule of the resulting compound may in every case be represented by two volumes. To the definitions of the terms molecule and atom already given, the following may be added : The molecule of an element or of a compound is the smallest por- tion capable of existing in a free state—at all events during any appreciable interval of time. An atom of an element is the smallest part of that element capable of entering into or being expelled from a chemical compound—the smallest part that exists in the molecule of any of its compounds. The atomic weight of an element expresses the number of times its atom is heavier than the atom of hydrogen. The molecular weight of an element or compound is, as already stated, the sum of the atomic weights of the atoms in its molecule. The various methods of determining vapor-densities will be fully described in the part of this work relating to organic chemistry : they are of great importance in fixing the molecular weights of organic compounds. The principles involved in these methods may be stated in a few words. The method of Dumas, applicable both to gases and to vapors, consists in ascertaining the weight of that quantity of the substance which in the gaseous state occupies a known volume. In the method of Gay-Lussac, which can be employed only in the case of vapors, the reverse principle—that of ascertaining the volume occupied in the gaseous state by a known weight of substance—is employed. In both cases the temperature of the gas or vapor, and the pressure at which it is measured, must be carefully noted. The relation of the weight of a given volume of substance in the gaseous state to the weight of an equal volume of air or hydrogen at the same temperature and pressure, constitutes the vapor-density of the substance. In order that results obtained in the measurement of gases and vapors may be comparable, it is usual to calculate what the volumes would have been had the measurement been made under the pressure of 760 millimetres of mercury (this being the average pres- sure of the atmosphere), and at the temperature of 0° C. This pro- cess is known as “ reduction to standard temperature and pressure.” It is employed even in cases where the substance does not exist in the gaseous state under these conditions of temperature and pressure. Any other temperature and pressure might have been chosen, and the relations of the volumes of different gases so reduced would have remained exactly the same. If v be the volume of a gas or vapor measured at the temperature of t° C., and under the pressure of p millimetres of mercury, its volume Vat 0° C and 760 millimetres will be; V = L- -760(1 + 273)* This formula may easily be deduced from the laws of Boyle and Charles. • 60 INORGANIC CHEMISTRY. All other direct methods of determining vapor-densities are modi- fications of the two just mentioned. The method of ascertaining the molecular weight from the vapor- density is unfortunately limited in its application. Allusion has already been made (p. 56) to the practical impossibility of determin- ing the vapor-density in the case of the great majority of the elements. As regards compounds, many of these decompose in assuming the gaseous state, so that their vapors consist of molecular mixtures more or less heterogeneous, from the density of which no conclusion can be drawn as to the molecular weight of the original compound. In the case of such compounds, an indirect method has to be re- sorted to. It will be best to illustrate the application of this method by a case in which the molecular weight has already been deduced from the vapor-density. The analysis of a compound gives a certain percentage composition, from which an empirical formula may be calculated. In this way the empirical formula CH20 is obtained for acetic acid. But it is evident that any multiple of this formula, C 2H+02, C 3H603, etc., would correspond equally well with the same percentage composition, and the question therefore arises, which is the true molecular weight? Experiment shows that 107.7 parts by weight, or 1 atom of silver, may be substituted for 1 part of hydrogen in acetic acid ; and further, that in this manner one-fourth part of the entire hydrogen present in the acid may be displaced. As fractions of atoms do not exist, the only legitimate conclusion is that the number of atoms of hydrogen in the molecule of acetic acid is four, or some integer multiple of four. At this point the decision is rendered easy by the knowledge derived from other .sources that acetic acid belongs to the class of the mono- basic acids in the molecule of which only one atom of hydrogen can be displaced by silver. Hence the molecular formula of acetic acid must be G2H402. Adding together the atomic weights (see table, p. 38) of all the atoms in the molecule, the molecular weight 60 is ob- tained. Now the vapor-density of acetic acid determined at 300° has been found to be 2.08 (air = 1). Substituting this value for din the formula AT = 28.9 X d, we find If =60.1 as the molecular weight of acetic acid, a number which agrees very well, within the limits of experimental error, with that deduced above. As the operations of weighing, on which the determinations of the atomic weights depend, can be performed with greater accuracy than those involved in ascertaining vapor-densities, it is usual to select as the most trustworthy the molecular weight obtained by adding to- gether the atomic weights of all the atoms in the molecule, using the vapor-density only to decide between two or more possible molecular weights. Thus in the case of acetic acid, the formulte CH,O, C 2H+02, and C 3H6Os would represent the molecular weights 30, 60, and 90 respectively. The number 60.1 obtained from the vapor-den- sity leaves no doubt as to which of these is the true molecular weight.. Melissic acid is a compound of high molecular weight, not volatile ATOMIC WEIGHTS. 61 without decomposition. Its whole chemical behavior shows that it belongs to the same class of acids as acetic acid; this knowledge is of use in determining the molecular weight. The empirical formula is Cl 5H300, which would correspond to the molecular weight 226. We have already seen that 107.7 parts of silver can displace I part of hydrogen in 60 parts of acetic acid. In like manner experiment shows that 1 part of hydrogen in 452 parts of melissic acid may be displaced by 107.7 parts of silver. The molecular formula of this acid is therefore C3OH60O2 = 452, or twice as great as the empirical formula, as was also the case with acetic acid. When a substance is not volatile without decomposition, and is moreover incapable of forming compounds from which conclusions can be drawn as to its molecular weight, the determination of this latter is beset with still greater difficulties. In this case it is neces- sary to take the compound, as it were, to pieces, either by breaking it up into two or more known compounds, or by destroying one part and leaving the rest intact, the object being in every case to arrive at compounds of known molecular weight. In this way more or less trustworthy conclusions as to the molecular weight of the original compound may sometimes be arrived at; but this method is far in- ferior in the certainty of its results to the two already described. CHAPTER VII. ATOMIC WEIGHTS. 1. Deduction of the Atomic Weight of an Element from the Vapor-density of its Compounds. The atomic weight of an element is that weight which is the greatest common divisor of the various weights of that element oc- curring in the molecules of its compounds, the atomic weight of hy- drogen being taken as unity. The atomic weights are thus relative, not absolute weights. As the molecular weights of volatile elements and of those com- pounds which can be vaporized without decomposition have alone been determined with certainty (all other methods, whatever proba- bility of accuracy their results may possess, being based more or less on analogy), it is necessary, in order to determine the atomic weight of an element according to the above definition, that it should form a number of compounds volatile without decomposition. The fol- lowing tables show the application of this method: INORGANIC CHEMISTRY. 1 mol. Mol. weight. Contains parts by weight. Mol. formula. Hydrogen, .... Chlorine, Oxygen, Sulphur, .... | Nitrogen, Hydrochloric acid,. . Hydrocyanic acid, . . Nitric oxide, .... Nitrous oxide, . . . Water, Carbonic oxide, . . . Carbonic anhydride, . Methylic hydride, , . Methylic chloride, . . Methylenic dichloride, Chloroform, .... Carbonic tetrachloride, Dicarbonic hexachloride, Acetone, Methylic oxalate, . . Sulphuretted hydrogen, Disulphur dichloride,. Sulphurous anhydride, Boric fluoride, . . . Silicic fluoride, . . . 2 71 32 64 192 28 36.5 27 30 44 18 28 44 16 50.5 85 119.5 154 237 58 118 34 135 64 68 104.2 2 Hydrogen, 71 Chlorine, 32 Oxygen, 64 Sulphur, 192 Sulphur, 28 Nitrogen, 1 Hydrogen, 35.5 chlorine,. . . 1 Hydrogen, 12 carbon, 14 nitrogen, 14 Nitrogen, 16 oxygen, .... 28 Nitrogen, 16 oxygen, .... 16 Oxygen, 2 hydrogen, .... 12 Carbon, 16 oxygen, 12 Carbon, 32 oxygen, 12 Carbon, 4 hydrogen, .... 12 Carbon, 3 hydrogen, 35.5 chlorine, 12 Carbon, 2 hydrogen, 71 chlorine, 12 Carbon, 1 hydrogen, 106.5 chlorine, 12 Carbon, 142 chlorine, .... 24 Carbon, 213 chlorine, .... 36 Carbon, 6 hydrogen, 16 oxygen, 48 Carbon, 6 hydrogen, 64 oxygen, 32 Sulphur, 2 hydrogen, .... 64 Sulphur, 71 chlorine, .... 32 Sulphur, 32 oxygen, .... 11 Boron, 57 fluorine, 28.2 Silicon, 76 fluorine, .... h2. Cl,. 0,. 8,. S6. N,. HC1. HCN. NO. N20. oh,. CO. co2. CH4. CHSCI. CH2C12. CHClo. CC14. C2C16. C31I60. c4h60, sh2. S2CI2. so2. BF3. SiF4. In the next table the above results are arranged so that the atomic weights of the various elements under discussion may be deduced. The first column contains the name of the element; the second, the relative weights of it occurring in the molecules of its compounds above enumerated—the smallest of these weights, which generally coincides with the atomic weight, being placed first; and the third, the greatest common divisor of these numbers, this last being iden- tical with the atomic weight: Element. Relative weights. G. C. D. Hydrogen, 1, 2, 3, 4, 6, 1 Chlorine, 35.5,71,106.5,142,213, 35.5 Oxvgen, 16, 32, 64, 16 Sulphur, Nitrogen, 32, 64, 192, 32 14, 28, 14 Carbon, 12, 24, 36, 48, 12 Fluorine, 57, 76, 19 In this way the atomic weights of these elements have been deter- mined. It will be noticed that the smallest relative weight of fluorine occurring in the molecule of either of its compounds above mentioned is thrice its atomic weight. A compound, hydrofluoric acid, con- ATOMIC WEIGHTS. 63 taining one atom of fluorine to one of hydrogen, has long been known, but, though capable of existing as a gas even at ordinary temperatures, its vapor-density could not be ascertained, owing to its property of attacking the vessels of glass or porcelain in which it has to be meas- ured. Latterly, however, the problem has been solved, and hydro- fluoric acid is found to possess the molecular formula HF = 20,* and to consist of 19 parts of fluorine to lof hydrogen. Organic compounds of fluorine, containing only one atom of this element in the molecule, have also been discovered. They are volatile and do not attack glass, so that their vapor-density may be determined in the ordinary way. The existence of these compounds places the number now accepted as the atomic weight of fluorine on a much surer basis. It is evident that the above method alone can never afford absolute certainty as to the atomic weights of the elements, since we can never be sure that a compound will not be discovered containing in its molecule either a smaller relative weight of an element than that which has been deduced from the known compounds of that element, or some relative weight which is not a rational multiple of the re- ceived atomic weight. If, for example, a compound containing 8 parts, or 24 (or any odd multiple of 8) parts of oxygen in the mole- cule were to be discovered, it would be necessary to change the atomic weight of oxygen from 16 to 8. Fortunately, however, two other methods of fixing the atomic weight are known (see pp. 65 and 67), and the agreement prevailing between the numbers determined by these three totally independent methods, increases enormously the probability of their correctness. Apparent Exceptions to Avogadro s Laic.—There are cases in which the molecular weights as deduced from the vapor-densities give values which are less than the sum of the weights of the smallest possible number of whole atoms which can go to form the compound. The following three substances, at ordinary temperatures solids, will serve as illustrations : The vapor-density of ammonic chloride has been found to be 0.89 (air = 1). The molecular weight would therefore be M= 28.9 X 0.89 = 25.7. The smallest stoechiometricf molecule is NH4CI = 53.5 =2 X 26.75. The molecular weight deduced from the vapor-density would therefore correspond to the formula N|H2Clj :in other words, the accepted atomic weights of nitrogen and chlorine would have to be halved. Phosphoric chloride has a vapor-density of 3.65, or only half of that required by its smallest stoechiometric formula PC15. The formula * The above is the molecular weight of hydrofluoric acid at 100°. At 25° it has the molecular weight 40, corresponding to the molecular formula H2F2. This in no invalidates the foregoing conclusions. t Stoechiometric, pertaining to the atomic weights. INORGANIC CHEMISTRY. would therefore have to be written PiClf, and the atomic weights of phosphorus and chlorine would have to be halved. A still worse complication is introduced by the vapor-density of ammonic carbamate, which is 0.89, or only one-third of that which its smallest possible formula lST2H6C02 demands. The molecular formula would therefore be N|H2Cso|. In order to introduce whole numbers of atoms into this last formula, and at the same time into that of ammonic chloride, NjH2CJj, it would be necessary to give to the atomic weight of nitrogen a value only one-sixth of that now assigned to it, or 2.33 instead of 14. This would further involve the assumption that nearly all the other com- pounds of nitrogen contain at least six atoms of nitrogen. Fortunately, however, these alterations, which would introduce indescribable confusion into chemistry, would also be erroneous. It has been proved that all these compounds decompose in volatilizing. The molecule of ammonic chloride (NH4CI) breaks up into one mole- cule of ammonia (NH3) and one of hydrochloric acid (HCI). The vapor thus contains twice as many molecules as it would have done had no decomposition taken place; it therefore occupies twice the volume, and consequently possesses only half the density. The same holds good concerning phosphoric chloride (PCIS), which breaks up into equal molecules of phosphorous chloride (PCI3) and free chlorine (Cl 2). Ammonic carbamate (JST2H6CO2) decomposes into two mole- cules of ammonia (NIT 3, NH3) and one of carbonic anhydride (CO2), so that the volume is three times, and the density only one-third as great as would be the case if no decomposition had taken place. Since in all these cases the products of decomposition recombine on cooling to form the original compound, the difficulty lay in proving that a decomposition had really taken place. However, this has been satisfactorily accomplished by various methods, both direct and indirect; so that it is not necessary either to doubt the validity of Avogadro’s law, as some chemists were inclined to do, or to intro- duce intricate and contradictory changes in the accepted atomic weights. 2. Determination of the Atomic Weights by means of Isomorphism.* Many different compounds crystallize in the same or nearly the same forms. For example, the salts Plumbic nitrate, .... PbNA. Baric nitrate, .... BaN206. Strontic nitrate, .... SrN206. crystallize in the same forms of the regular system (see Crystallog- raphy). As any given form of the regular system has invariably the same angles, the identity of form in the above three cases is absolute. Again : * The selection of examples of isomorphism is borrowed from Kopp’s Theoretische Ghemie. ATOMIC WEIGHTS 65 Nickelous sulphate, . . . . . . NiS04,60H2, Niekelous seleniate, . . . NiSe04,60 H2, Zincic seleniate, crystallize in the same quadratic forms, with angles almost identical in the three cases, and with the same cleavage.* The following com- , ' o o pounds: Zincic sulphate, .... Nickelous sulphate, . . . . . . NiS04,70H2, Magnesic sulphate, . . . . . . MgS04,70H9, Magnesic seleniate, . . . . . . MgSe04,70H9, Magnesic chromate, . . . . MgCr04,70H2, crystallize in very similar forms of the rhombic system, with almost the same angles. Compounds which, like the above, crystallize in the same or nearly the same forms, and possess similar constitution, are termed isomorphous. In an isomorphous group those elements which occur in all the mem- bers are called the common elements; those which may be varied with- cut producing a change of crystalline form, the corresponding elements. The corresponding elements are frequently termed the isomorphous dements, although they do not, when isolated, necessarily crystallize in the same forms. The sense in which the term isomorphous is used when applied to compounds must not be confounded with that which it bears m reference to elements. In the former case it means : u possessing the same form;” in the latter, “producing the same form.” In each of the above groups It will be noticed that all the compounds contain the same number of atoms. It has further been found by ex- periment that in an isomorphous group, the corresponding elements occur in the relative proportions of their atomic weights as determined by Avogadro’s law. Hence it is only necessary to know the atomic Weight of one of the corresponding elements in a group of isomorphous compounds in order to determine the atomic weights of all the rest. Hut before illustrating this, it will be necessary to describe the various groups of isomorphous elements. In such a group the analogous com- pounds which the various members form with the same element or ele- ments are frequently, but not necessarily, isomorphous. 1. Sulphur, Selenium, Manganese, Chromium.—Sulphides and selen- ides are frequently isomorphous, for instance : PbS and PbSe, Ag.JS and Ag2Se. The salts of sulphuric, selenic, manganic, and chromic acids, with the same base, and containing the same number of molecules °f water of crystallization, are generally isomorphous. 2. Magnesium, Calcium, Manganese, Iron, Cobalt, Nickel, Zinc, Cad- mium, Copper.—The carbonates of these metals crystallize in rhombo- nedra with rhorabohedral cleavage. The cleavage rhombohedra have almost the same angles. The sulphates are also for the most part iso- . Cleavage is the tendency which some crystallized substances display when broken, to rJ?|t m directions parallel to the faces of certain crystalline forms of these substances., he artificial forms thus produced are known as “cleavage forms.” 66 INORGANIC CHEMISTRY. morphous, and the same is the case with the double sulphates of these metals with potassium'and ammonium. 3. Manganese and Iron, both members of the preceding group, also form another group with Chromium and Aluminium. The three ses- quioxides Fe2Oa, Cr203, and A1203, are isomorphous. The sesqui- oxides of these four metals combine with monoxides of the general formula R"0 to form the spinelles, which all crystallize in the regular system and possess the general formula Rr/0, W,rf)a. The sesquioxides also enter into the composition of the alums, which all crystallize in the regular system. 4. Calcium has also isomorphous relations with Strontium, Barium, and Lead. All four are connected by their carbonates (calcium as arragonite); calcium and lead by their tungstates; strontium, barium, and lead by their anhydrous sulphates. A simple enumeration of some of the remaining isomorphous groups must suffice: 5. Tungsten and Molybdenum. 6. Tin and Titanium. 7. Palladium, Platinum, Iridium, and Osmium. 8. Potassium and Ammonium. 9. Sodium and Silver. 10. Phosphorus, Arsenic, and Antimony. 11. Chlorine, Bromine, lodine. Elements which are isomorphous with the same element are not neces- sarily isomorphous with each other. It would be incorrect, for exam- ple, to say that iron and sulphur must be isomorphous because they are both (in different ways) isomorphous with manganese. Only those elements can be said to be isomorphous which occur in the same true group of isomorphous compounds; and in a true group of isomorphous compounds all the members possess the same crystalline form and an analogous atomic composition. It only remains to give an illustration of the method of applying the law of isomorphism to the determination of the atomic weights. From the vapor-density of their compounds, chlorine and sulphur have been found to possess the atomic weights Cl = 85.5 and S = 32. In the isomorphous sulphates and manganates (isomorphous group 1), the corresponding elements occur in the proportion of 32 parts by weight of sulphur to 55 of manganese. In the isomorphous perchlorates and permanganates, the proportion in which the corresponding elements occur is 35.5 parts of chlorine to 55 of manganese. The atomic weight of manganese is therefore 55. But the metals of the 2d isomorphous group are contained in their isomorphous carbonates and sulphates in the following relative proportions: manganese 55, magnesium 24.4, calcium 40, iron 56, cobalt 58.6, nickel 58.6, zinc 65.3, cadmium 112, copper 63.2; and these are therefore the atomic weights of those ele- ments. In like manner it is only necessary to refer the proportions in which the metals of the 4th isomorphous group occur in their isomor- phous compounds to the atomic weight of calcium just deduced, Ca=4o, in order to determine the atomic weights of barium, strontium, and lead, which are thus found to be Ba = 137, Sr = 87.5, Pb = 206.5. ATOMIC WEIGHTS. 67 The foregoing enumeration of isomorphous groups includes only some of the most prominent. There are many others which serve as connecting links, so that it is possible by means of the law of isomorph- ism to determine the atomic weights of nearly all the elements. T O J isomorphous compounds possess the property of crystallizing together m various proportions to form homogeneous crystals belonging to the same system as the compounds themselves. These crystals are generally distinguished by possessing simpler forms—less variety of faces—than the crystals of the pure compounds. If the angles of the latter differ slightly from each other, the angles of the mixed crystals will possess values which lie between those of the pure compounds. Thus the ter- minal angle of the cleavage rhombohedron of pure calcium carbonate is 105° s'; that of pure magnesic carbonate, 107° 25'; whilst in the case of their isomorphous mixtures, this angle varies between these two limits, inclining in the direction of the compound which predominates m the mixture. A substance which crystallizes in two different forms not reducible to the same system is termed dimorphous. It sometimes happens that two dimorphous compounds are isomorphous, in which case the two distinct forms frequently correspond in the two compounds. This double isomorphism is known as isodimorphism. Antimonious oxide, Sb2o3, occurs naturally in regular octahedra as senarmontite, and in rhombic prisms as valentinite. Arsenious anhydride, As203, is found m nature in regular octahedra as arsenic bloom and in rhombic prisms ns claudetite, these two forms respectively corresponding with those of antimonious oxide, with which arsenious anhydride is thus isodi- morphous. The law of isomorphism was first enunciated by Mitscherlich, in 1819. The determinations of atomic weights by means of this law are not always absolutely certain. This uncertainty has its root in the fact that various undoubtedly isomorphous compounds are known in which the number of atoms in the molecule is different. Thus the salts of potas- sium (K) and ammonium (NH4) are isomorphous. Baric permanganate, DaMn208, is isomorphous with anhydrous sodic sulphate, Na2S04. In none of these compounds can the corresponding dements be said to be substituted for each other in the proportion of their atomic weights. 3- Determination of the Atomic Weights from the Specific Heats of the Elements in the Solid State. If a kilogram of water at 100° be mixed with a kilogram of water at o°, the temperature of the mixture will be 50°, the mean of the other two temperatures. If a kilogram of iron filings at 100° be mixed a kilogram of water at o°, the temperature of the whole will not be higher than 10°. As, therefore, a given weight of water in cooling through 50° can raise the temperature of an equal weight of water through 50°, and as a given weight of iron filings in cooling through can raise an equal weight of water through only 10°, it is evident that equal weights of iron and water at the same temperature contain 68 INORGANIC CHEMISTRY. very different amounts of heat. Calculated from the above figures, the quantities of heat contained in equal weights of water and iron at the same temperature will be as to or as Ito And as the heat which a body gives off in cooling is equal to that which it has taken up in heating, it will require 9 times as much heat to raise the temperature of a given weight of water through a given number of degrees, as it will to raise the same weight of iron through an equal number of de- grees. The relative capacities of bodies for heat are known as their specific heats, that of water being taken as unity. For many reasons it is useful to have a unit of heat, by means of which the heat evolved or absorbed in chemical or other processes may be measured. For this purpose that quantity of heat required to raise the temperature of 1 gram of water from 0° to 1° C. is employed as the standard of measurement, and is known as the unit of heat, thermal unit, or calorie. As the specific heat of water is the unit of the specific heats, it is evident that in order to find how many units of heat are required to raise the temperature of a body through any number of degrees of the centigrade scale, it will only be necessary to multiply together the weight of the body expressed in grams, its specific heat, and the number of degrees through which its temperature has been raised.* Thus the quantity of heat required to raise the tempera- ture of 2 grams of iron through 90°, or of 180 grams through I°, or of 1 gram of water through 20°, or of 2 grams through 10°, is in every case the same, namely 20 thermal units. Dulong and Petit were the first to determine the specific heats of a number of the chemical elements, and they arrived at the remarkable result, that the specific heats of the elements in the solid condition are in- versely as their atomic weights. If instead of determining the specific heat of equal weights of the elements, the latter be taken in the propor- tion of their atomic weights, the specific heats of these atomic weights will be equal, or as this may be expressed : the capacities for heat of the atoms of different elements in the solid state are equal: all the elements in the solid state have the same atomic heat. The atomic heat may be found by multiplying the specific heat of an element by its atomic weight. The average value of the atomic heat for the different elements is 6.4. The slight variations which the atomic heats of the various elements display, arise first from the difficulty of determining accurately the specific heat, and secondly from difference of physical condition in the elements—the chief disturbing influence depending upon the fact that the specific heat of an element rises with the temperature, being greatest near the fusing point, whilst the specific heats are generally determined between 0° and 100°, and consequently at varying distances from the fusing points of the different elements. It is evident that the law of Dulong and Petit must offer a very valuable means of checking doubtful atomic weights, and of determin- ing such as are not within the reach of the other two methods. Thus, * This mode of calculation is based on the assumption that the specific heat of a body is the same at all temperatures, which is oply approximately correct. As will he shown later, the specific heat increases with the temperature. ATOMIC WEIGHTS. 69 gold forms no volatile compounds, and its isomorphism with other ele- ments is not sufficiently marked to be available as a means of fixing its atomic weight. But the specific heat of gold has been found to be 0.0324, and this number multiplied by 196, the accepted atomic weight of gold, gives 6.35, closely approximating to the average atomic heat of the elements, from which it may be concluded that 196, and no mul- tiple or sub-multiple of this number, is the true atomic weight of gold. A glance at the table of specific heats on p. 73, in which the elements are arranged in the order of their atomic weights, will show that the deviations from the law of Dulong and Petit follow a certain rule. In the case of the elements of high atomic weight, the agreement is almost always good, and with these elements it is to be noted that the varia- tion of the specific heat with the temperature at which it is determined 18 but small. The notable exceptions to the law are to be found among the elements which combine the two properties of low atomic weight and loio atomic volume {q. v.). In the following list of these exceptional ele- ments, the specific heats have been determined at temperatures below 100° C. (212° F.). The brackets denote indirect determinations (see Aehmann’s Law, p. 70): Name of element. Atomic heat. Aluminium, . . 5.7 Phosphorus, . . 5.3 Sulphur, Nitrogen, . . 5.1 . . (5) Fluorine, . . (5) Oxygen, . . (4) Silicon, . . 3.8 Beryllium (Glucinura), . . 3.7 Boron, . . 2.7 Hydrogen, . . (2.3) Carbon (as diamond and in its compounds), . . 1.8 A reference to Lothar Meyer’s curve of the elements (see diagram, Classification of the Elements according to their Atomic Weights) will show, that the whole of these exceptional elements are to be found in the lower portions of the first three periods of the curve—a position which, from the nature of this curve, falls to these elements in virtue of their i°w atomic weight and low atomic volume. That low atomic weight alone is not sufficient to produce deviation from the law of Dulong and I etit, is very clearly shown by the fact that three elements of low atomic*weight—lithium, sodium, and potassium—which, however, owing to their relatively high atomic volume, form maxima of the curve, perfectly conform to the law. A straight dotted line, cutting the curve, has therefore been drawn to indicate the “limit of validity ct the law of Dulong and Petit.” The exceptional elements are all to oc found below this line. It is probable, however, that even for these exceptional elements there is a temperature at which they conform to this law. H. F. Weber, ho has carefully determined the specific heats of carbon and silicon °r a great range of temperature, finds that the specific heat rapidly 70 INORGANIC CHEMISTRY. increases with the temperature until a point is reached at which these elements approximately obey the law; that is to say, the deviations are not much greater than in the case of aluminium, thus leaving no reasonable doubt about the atomic weight. Above this point the specific heat rises only very slowly with the temperature. This lower limit of conformity to the law lies in the case of silicon at about 200° C., in the case of carbon about 600° C. It is worthy of note that the various modifications of carbon, which at ordinary temperatures possess widely different specific heats, have the same specific heat as soon as the above limit is reached. Boron shows a similar rapid rise of specific heat; but the observations have not been carried to temperatures sufficiently high to determine the lower limit of conformity in the case of this element; it, however, probably lies between 500° and 600° C. Dulong and Petit tried without success to extend the law of specific heat to compounds. This was finally accomplished by Neumann (1831), who showed that chemically equivalent quantities of similar compounds have the same capacities for heat. If the product of the molecular weight into the specific heat be termed the molecular heat of a compound, this law may be expressed : Similar compounds have the same molecular heeds For example: Compound. Mol. formula. Mol. weight. Sp. heat. Mol. heat. Lithic chloride, . . Li Cl 42.5 0.2821 12 Sodic chloride, . . NaCl 58.5 0.2140 12.5 Potassic chloride, KC1 74.5 0.1730 32.9 Argentic chloride, . AgCl 143.2 0.0911 13 It is possible in this way to determine the atomic heat of elements which do not exist at ordinary temperatures in the solid state. Thus, by subtracting from the molecular heat of potassic chloride, 12.9, the atomic heat of potassium, 6.6, the atomic heat of chlorine is found to be 6.3. A study of the above-mentioned chlorides shows that the atomic heat of chlorine thus deduced varies according to the chloride employed; but the method of calculating its value by subtracting the atomic heat of the other element exaggerates these errors. It is further evident that the danger of error in this indirect method of determining the specific heat of an element will be greater the greater the relative num- ber of atoms of other elements contained in the molecule of the com- pound employed. But if the molecular heat of a compound be di- vided by the number of atoms in the molecule, the variations caused by difference of physical conditions in different compounds will be dis- tributed among the atomic heats of the several atoms in the molecule (which are probably all affected in the same direction by such varia- tions), and the average atomic heat of the elements contained in that compound will be obtained. Thus the molecular heats of the above chlorides divided by 2 give numbers varying from 6 to 6.5, suf- ficiently approximating to 6.4, the average atomic heat of the elements in the solid state. ATOMIC WEIGHTS. 71 In this way Neumann’s law has been successfully applied in verifying the atomic weights of elements, the specific heats of which had not been directly determined. Thus in the case of barium, strontium, and calcium, chemists were in doubt whether these elements possessed the atomic weights Ba = 137, Sr = 87.5, and Ca = 40; or, only the half of these weights, ba = 68.5, sr = 43.8, and ca = 20—these smaller values being formerly universally employed. In these two cases the formulae of the chlorides would be respectively: Formula. Mol. weight. Formula. Mol. weight. BaCl2, . .... 208 bad, . . . . . 104 SrCJ2, . .... 158.5 srCl, . . . . . 79.3 Cad2, . .... Ill caCl, . . . . . 55.5 The number of atoms in the molecule is in the first case 3, in the second 2, The specific heats of these compounds were found to be: Baric chloride, 0.0902 Strontic chloride, 0.1199 Calcic chloride, 0.1642 vr ~ . molecular weight X specific heat ~ . JNow the expression - v—- ought to number ot atoms in molecule be approximately equal to 6.4, the average atomic heat. Substitut- ing in this expression the above values, we find for bad, . . . 104 X 0.0902 = 4.7, 2 srCl, . . . 79.3 X 0.1199 = 4.75, 2 caCl, . . . 55.5 X 0.1642 = 4.55; 2 and for BaCl2, . . 208 X 0.0902 _ 3 6.23, SrCl2, . . . 358.5 X 0.1199 _ 6.33, * 3 CaC]2, . . 111 X 0.1642 _ 3 6.07. The values 6.23, 6.33, and 6.07 approximate with sufficient close- ness to 6.4; whereas, 4.7, 4.75, and 4.55 differ widely from this num- ber. The formulae of the chlorides must, therefore, be written BaCl2, ®rCl2, and CaCl2, and the three elements must possess the atomic weights Ba = 137, Sr = 87.5, and Ca = 40. Only a few years ago the specific heat of metallic calcium was determined for the first time by Bunsen, and was found to be 0.1704. This number, multiplied by 40, the 72 INORGANIC CHEMISTRY. atomic weight of calcium, gives 6.82 as the atomic heat of this element, thus directly proving the correctness of the above deduction. In applying Neumann’s law to compounds in which any of the ex- ceptional elements occur, it is necessary to introduce the special value for the atomic heat in calculating the molecular heat of the compound. In the case of the other elements, the average atomic heat, 6.4, may be employed without sensible error : Name of Molecular Molecular heat. compound. formula. Calculated. Found. Antimonious sulphide, Potassic pvrophos- Sb2S3. (2X6.4) + (3X5.1) =28.1 28.6 phate, k4p2o7- (4 X 6.4) + (2 X 5.3) + (7 X 4) = 64.2 63.1 Calcic fluoride, . . CaF2. 6.4 +(2X5) =16.4 16.3 Cupric oxide, . . . CuO. 6.4 + 4 = 10.4 10.2 Silicic anhydride, Si()2. 3.8 + (2X4) = 11.8 11.5 Boric anhvdride, . . B./)3. (2X2.7)+(3X4) =17.4 16.6 Sodic metaborate,. , Dicarbonic hexachlo- NaBOa. 6.4 + 2.7 + (2X4) = 17.1 16.9 ride, C2CI6. (2X1.8) + (6X6.4) — 42 42.2 Succinic acid, . . . c4h604. (4 X 1.8) + (6 X 2.3) + (4 X 4) = 37 36.9 Tims, the molecular heat of a compound is the sum of the atomic heats of its elements.* This law, like the law of Dulong and Petit, of which it is a corollary, is only an approximate law. It generally holds in the case of chlorides, but is an unsafe guide in the case of oxides, especially if tfoe number of atoms in the molecule be large (see page 70); indeed, in some cases, the attempt to deduce the atomic heat of an element from the molecular heat of its oxide has led to fallacious results. The following table contains the specific and atomic heats of all ele- ments for which the determination has been made. In the case of carbon, silicon, and boron, the values obtained at higher temperatures are em- ployed. The elements are arranged in the order of their atomic weights. The bracketed numbers represent indirect determinations: * The law of Neumann that the molecular heat of a compound is the sum of the atomic heats of its elements, taken in connection with the fact that the known elements possess an atomic heat approximating to 6.4, has a direct bearing upon the view some- times advanced that many or all of the known elements are in reality compounds. It is evident either that these supposed compounds do not contain as constituents any of the known elements, since these have already an approximate atomic heat of 6.4, and the resulting “ compound ” element would necessarily possess a higher atomic heat; or that the mode of combination is totally different from any yet known to chemists. Further, as all the known elements have approximately the same atomic heat, the con- clusion appears almost unavoidable, on the “ compound ” theory, that they are all com- pounds of exactly the same complexity—containing the same number of constituent atoms, a degree of uniformity which nature does not usually exhibit. Kundt and Warburg’s proof (p. 56) that the molecule of mercury has no internal motion of parts, and is, therefore, in all probability truly monatomic, also appears to militate against the “ compound ” theory of the elements. ATOMIC WEIGHTS. 73 Table of the Specific Heat of the Elements in the Solid State. Name of element. Atomic weight. Specific heat. Atomic heat. f Hydrogen, 1 (2.3) (2.3) Lithium, 7 0.94 6.6 Beryllium (Glueinum), 9 045 4.0 Boron, 11 0.5* ? 5.5 Carbon, 12 0.46 5.5 Nitrogen, 14 (0.36) (5) Oxvgen, 16 (0.25) (4) Fluorine, 19 (0.26) (4.9) Sodium, 23 0.29 6.7 Magnesium, 24.4 0.25 6.1 Aluminium, 27 0.21 5.7 Silicon, 28.2 0.20 5.6 Phosphorus, 31 0.17 5.3 Sulphur, 32 0.16 5.1 Chlorine, ... 35.5 (0.18) 6.4 Potassium, 39 0.17 6.6 Calcium, 40 0.17 6.8 Titanium, 48 (0.13) 6.2 Chromium, 52 (0.12) 6.2 Manganese, 55 0.12 6.6 Iron, 56 0.11 6.2 Nickel, 58.6 0.11 6.4 Cobalt, 58.6 0.11 6.4 Copper, 63.2 0.094 5.9 Zinc, 65.3 0.094 6.1 Gallium, . , , 68.8 0.079f 5.4 Arsenic, 75 0.081 6.1 Selenium, 79 . 0.075 5.9 Bromine, 80 0.084 6.7 Rubidium, 85.3 (0.077) (6.6) Strontium, 87.5 (0.074) (6.5) Zirconium, 90 0.066 5.9 Molybdenum, 95.5 0.072 6.9 Rhodium, 104 0.058 6.0 Ruthenium, 104 0.061 6.4 Palladium, 105.7 0 059 6.2 Silver, 107.7 0.056 6.0 Cadmium, 112 0.057 6.4 ■ * Indium, 113.4 0.057 6.5 Tin, 118 0.056 6.6 Antimony, 120 0.051 6.1 Tellurium, 125 0.047 59 Iodine, 127 0.054 6.9 Barium, 137 (0.047) (6.4) Lanthanum, 138.5 0.045 6.2 Cerium, 140.5 0.045 6.3 Didymium, 146 0.046 6.7 Tungsten, 184 0.033 6.1 Iridium, 192.5 0.033 6.4 Platinum, 194.4 0.033 6.4 Gold, 196 0 032 6.3 Osmium, 198.6 0.031 6.2 Mercury, . . 200 0.032 6.4 Thallium, 204 0.034 6.9 Lead, 206.5 0.031 6.4 Bismuth, 208.2 0.031 6.4 Thorium, 233.4 0.028 6.5 Uranium, 238.5 0.028 6.7 * This is a hypothetical value deduced from the experiments of Weber, t This value was obtained from a determination performed within a limit of eleven degrees—a very narrow range of temperature. 74 INORGANIC CHEMISTRY. Another mode of expressing the above facts consists in stating what weight of each element has the same capacity for heat as 7 parts by weight of lithium, 7 being the atomic weight of that metal. If the law of Dulong and Petit were a perfectly strict law, the weights which satisfy these conditions would be identical with the atomic weights. In the following table the atomic weights are given side by side with these “specific heat equivalents” in order to indicate clearly in every case the extent of the discrepancy between the two values: Name of element. Specific heat. Weights con- taining equal quantities of heat. Atomic weight. Lithium, 0.94 7 7 Beryllium (Glucinum), . 0.45 14.6 9 Boron, 0.5 13.2 11 Carbon, 0.46 14.3 12 Sodium, 0.29 22.7 23 Magnesium, 0.25 26.3 24.4 Aluminium, 0.21 31.3 27 Silicon, 0.20 32.9 28.2 Phosphorus, 0.17 38.7 31 Sulphur, 0.16 41.1 32 Potassium, 0.17 38.7 39 Calcium, 0.17 38.7 40 Manganese, . . t . . 0.12 54.8 55 Iron, 0.11 59.7 56 Nickel, . . 0.11 59.7 58.6 Cobalt, . 0.11 59.7 58.6 Copper, 0.094 70.0 63.2 Zinc, 0.094 70.0 65.3 Gallium, 0.079 83.3 68.8 Arsenic, 0.081 81.2 75 Selenium, 0.075 87.7 79 Bromine, 0.084 78.3 80 Zirconium, 0.066 99.7 90 Molybdenum, .... 0.072 91.4 95.5 Rhodium, 0.058 113 104 Ruthenium, 0.061 108 104 Palladium, 0.059 112 105.7 Silver, 0.056 118 107.7 Cadmium, 0.057 115 112 Indium, 0.057 115 113.4 Tin, 0.056 118 118 Antimony, 0.051 129 120 Tellurium, 0.047 140 125 Iodine, 0.054 122 127 Lanthanum, 0.045 146 138.5 Cerium, ...... 0.045 146 140.5 Didymium, . . * . . 0.046 143 146 Tungsten, 0.033 199 184 Iridium, 0.033 199 192.5 Platinum, 0.033 199 194.4 Gold, 0.032 206 196 Osmium, 0.031 212 198.6 Mercury, 0.032 206 200 Thallium, 0.034 194 204 Lead, 0.031 212 206.5 Bismuth, 0.031 212 208.2 Thorium, 0.028 235 233.4 Uranium, ...... 0.028 235 238.5 Specific Heat Equivalents of Solid Elements. CHEMICAL NOTATION. ATOMICITY. 75 CHAPTER VIII. CHEMICAL NOTATION. ATOMICITY. The use of symbols in place of words, for recording the composition of chemical compounds, and of equations for expressing chemical changes, has long been necessary to accurate description, and has con- tributed in an important degree to the development of chemistry into an exact science. Unfortunately there has been, and still is, much diversity of opinion amongst chemists as to the best kinds of symbols to be used, and the extent to which these should be employed for ex- pressing the constitution, as well as the composition, of chemical com- pounds. It would serve no useful purpose and would only confuse the student to review the various systems of notation in actual use amongst chemists, and the description will therefore be here confined to two of those systems, which have been extensively used for many years, and as these systems are based on the doctrine of atomicity, this subject has been introduced into the present chapter. Symbolic Notation.—Every element is represented by a symbol, which is frequently the initial letter of the name of the element; but as, in some cases, the names of two or more elements begin with the same letter, it is necessary to distinguish them by the use of a second letter in small type, which is either the second letter of the word, or some other letter prominently heard in its pronunciation : thus carbon, cad- mium, cobalt, and cerium all begin with the same letter ; but they are distinguished by the symbols C, Cd, Co, and Ce. In the use of the single letters, the non-metallic elements have the preference; thus oxygen, hydrogen, nitrogen, sulphur, phosphorus, boron, carbon, iodine, and fluorine are expressed by the single letters O, H, N, S, P, B, C, I, and F ; whilst the metals osmium, mercury, nickel, strontium, platinum, bismuth, cobalt, iridium, and iron are symbolized by two letters each; thus Os, Hg (hydrargyrum), Ni, Sr, Pt, Bi, Co, Ir, and Fe (ferrum). In the selection of the single letters for other cases, preference is given to the most important element; thus, sulphur, selenium, and silicon are all non-metallic elements, beginning with the same letter ; but sul- phur being the most important, the single letter S is assigned to it, whilst selenium and silicon are denoted respectively by Se and Si. The symbols of compounds are formed by the juxtaposition of the symbols of their constituent elements. Such a group of two or more symbols is termed a chemical formula. Thus : Argentic chloride, AgCl. Zincic oxide, ZuO. The symbols not only represent the elements for which they are used, but they also denote a certain definite proportion by weight of each element; the formula HCI, for instance, does not merely denotea com- pound of hydrogen and chlorine, but it signifies a molecule of that compound containing one atom (1 part by weight) of hydrogen, and 76 INORGANIC CHEMISTRY. one atom (35.5 parts by weight) of chlorine. When, therefore, the molecule of a compound contains more than one atom or combining proportion of any element, it is necessary to express the fact in its formula : this is done by the use of a small subscript coefficient placed after the symbol of the element; Zincic chloride, ZnCh. Ferric chloride, Fe2CI6. Stannous chloride, Stannic chloride, SnCl4. When it is necessary to denote two or more molecules of any com- pound, a large figure is placed before the formula of the compound ; such a figure then affects every symbol in that formula : thus 3S04H2 means three molecules of the compound S04H2. The changes which occur during chemical action are expressed by equations, in which the symbols of the elements or compounds, as they exist before the change, are placed on the left, and those which result from the reaction on the right. Thus, taking an example from each of the five kinds of chemical action (see Chemical Affinity) we have (1) Zn + Cl3 == ZnCl2. Zinc. Chlorine. • Zincic chloride. (2) 2HCI + Zn = ZnCl2 + Ev Hydrochloric Zinc. Zincic Hydrogen, acid. chloride. (3) SO.Cu + (M)3)2Ba = S04Ba + (NO3)2Cu Cupric Baric nitrate. Baric Cupric sulphate. sulphate. nitrate. (4) (CN)O(NH4) = N2H4(CO). Amnionic cyanate. Urea. (5) 20H2 02 -f- 2F2. Water. Oxygen. Hydrogen. The sign -J-, as seen from the foregoing examples, is placed between the formulae of the molecules of the different substances which are brought into contact before the reaction, and of those which result from the change. This sign must never be used to connect together the con- stituents of one and the same chemical compound.. The sign is only very rarely used in chemical notation, but when employed it has the ordinary signification of abstraction; thus, S04F2 oh2 = so3. Sulphuric Water. Sulphuric acid. anhydride. Use of the Bracket.—The bracket has been employed in various senses in chemical formulae; but in the present work it is used in notation for one purpose only, viz., for expressing chemical combination between CHEMICAL, NOTATION. ATOMICITY. 77 two or more elements which are placed perpendicularly with regard to each other, and next to the bracket in a formula. Thus in the follow- ing cases, I. 11. 111. fc H3 fCH3 N030) \ CH3 JO Ba V (CH3 NOaOj the formula No. I. signifies that two atoms of carbon are directly united with each other, No. 11. that two atoms of carbon are linked together, as it were, by an atom of oxygen, the latter being united to both carbon atoms; whilst in like manner, No. 111. indicates that one atom of oxygen in the formula of the upper line is linked to another atom of oxygen in the formula of the lower line, by an atom of barium. Use of Thick Letters.—As a rule, the formulae in this book are so written as to denote that the element represented by the first symbol of a formula is directly united with all the active bonds (see p. 81) of the other elements or compound radicals following upon the same line: thus the formula S02(OH)2 (sulphuric acid) signifies that the hexad atom of sulphur is combined with the four bonds of the two atoms of oxygen, and also with the two bonds of the two semimolecules of hydroxyl. Such a formula is termed a constitutional formula.* Occasionally, however, owing to the atomic arrangement of a com- pound not being known, its formula cannot be written according to this rule; and in order to prevent such formulae, whether molecular or empirical,f from being mistaken for constitutional formulae, the first symbol of a constitutional formula will always be printed in thick,type. As a rule, the element having the greatest number of bonds will occupy this prominent position. Thus : Sulphuric acid, . . . . so2(OH)2. Water, . . OHo. Nitric acid, . . . . . N02(0H). Microcosmic salt, . . . PO(OH)(ONH4)(ONa). * For further information on this subject see Atomicity of Elements and Com- pound Radicals. f A molecular formula, sometimes called rational, is one in which the atomic compo- sition of a molecule is expressed, but without reference to the manner in which the elements are combined amongst themselves. An empirical formula merely expresses, by the smallest integers, the proportional number of atoms of each element entering into the composition of a compound. Thus the three formulae of ferric hydrate are written: Empirical formula,FFes Molecular formula, Fe2H606. Constitutional formula, Fe2(OH)6. Constitutional or rational formulae are therefore essentially molecular formulae, whilst empirical formulae afford no indication of the number of atoms contained in a molecule ; they are, in fact, only used to express the composition of substances, the molecular Weights of which are either unknown or cannot be inferred from analogy. 78 INORGANIC CHEMISTRY. Atomicity of Elements. It has been already stated that the atomic weight of an element is the smallest proportion by weight in which that element enters into oris expelled from a chemical compound. The atoms of the various elements, the relative weights of which are thus expressed, possess very different values in chemical reactions. Thus, an atom of zinc is equivalent to two atoms of hydrogen, for when zinc is brought into contact with steam at a high temperature, one atom of zinc expels from the steam two atoms of hydrogen, and occupies their place, thus: 0H2 + Zn = OZn + H2. Water. Zincic oxide. Again, when zincic oxide is brought into contact with hydrochloric acid, the place of the zinc becomes once more occupied by hydrogen, but two atoms of hydrogen are found to be necessary to take the place of one atom of zinc: OZn + 2HCI = ZnC]2 + OH2. Zincic Hydrochloric Zincic Water, oxide. acid. chloride. In like manner, one atom of boron can be substituted for three atoms of hydrogen, one of carbon for four, one of nitrogen for five, and one atom of sulphur for no fewer than six atoms of hydrogen. This combining value of the elementary atoms, which was first dis- covered in the compounds of certain metals with organic radicals, is termed their atomicity, equivalence, valency, or atom-fixing power ; and an element, with an atom-fixing power equal to that of one atom of hydrogen is termed a monad, one with twice that power a dyad, with thrice a triad, with quadruple & tetrad, with quintuple a pentad, and with an atom-fixing power equal to six times that of hydrogen, a hexad. To avoid any speculation as to the nature of the tie which enables an element thus to attach to itself one or more atoms of other elements, each unit of atom-fixing power will be named a bond,—a terra which in- volves no hypothesis as to the nature of the connection. A monad element has, obviously, only one such bond ; a dyad, like zinc, two; a triad, like boron, three, and so on. The number of bonds possessed by an elementary atom may be usefully symbolized by lines in the fol- lowing manner; Hydrogen, .... . . . . H— Zinc, \ / Boron, . . . . B 1 Carbon, 1 . . . . — C— CHEMICAL NOTATION. ATOMICITY. 79 JN itrogen, JN \ / Sulphur, —S— / \ In symbolic notation, the same idea is conveyed by the use of dashes and Roman numerals placed above and to the right of the symbol of the element, thus: Hydrogen, . . . H', Carbon, . . . Civ, Zinc, Zn", Nitrogen, . . Nv, Boron, .... B//r, Sulphur, . . STI. Elements with an odd number of bonds are termed perissads, whilst those with an even number are named artiads. With very few exceptions, elements, either alone or in combination, are never found to exist with any of their bonds free or disconnected ; hence, the molecules of all elements with an odd number of bonds are generally diatomic, and always 'polyatomic ; that is, they contain two or more atoms of the element united together. Thus: Symbolic. Graphic. Hydrogen, . . . • • H2 . . . . . H—H Chlorine, . . . Cl2 . . . . . Cl—Cl Nitrogen, . . . . . Nv2 . . . III 11 Phosphorus, . . . . . . . . Ill 111 p=p An element, with an even number of bonds, however, can exist as a monatomic molecule, its own bonds apparently satisfying each other. Thus: Symbolic. Graphic. Mercury, . . . . . Hg" . . . r~HgD Cadmium, . . . Cd" . . . v J Zinc, . . . Zn" . . . f—Zn—^ It is, nevertheless, obvious that such an element may also exist as a polyatomic molecule. Oxygen furnishes us with an example of this ; for, in its ordinary condition, it is a diatomic molecule, and, in the al- lotropic form of ozone, a triatomic molecule : Symbolic. Graphic. Oxygen, . . . . . 0"2 .. . . . . 0=0 0—0 Ozone, . . . . . . 0"3 . • . • . \/ o 80 INORGANIC CHEMISTRY. In order to avoid the unnecessary use of atomicity-marks in sym- bolic notation, they will never be attached to a monad, or to oxygen, which, it must be remembered, is always a dyad. Neither will the atomicity coefficient be attached to the tetrad element carbon, in the formulae of organic bodies, unless this element plays the part of a dyad, an occurrence of extreme rarity. When not otherwise marked, there- fore, carbon must always be understood to be a tetrad. It will also, as a rule, be unnecessary to mark the atomicity of the elements which are expressed by symbols in thick type, because their atomicity is clearly indicated by the sum of the atomicities of the ele- ments or compound radicals placed to their right, or connected with them perpendicularly by a bracket. Thus, in the formula f 0C1„ ICCI3, each atom of carbon is united with three atoms of the monad chlorine, whilst the bracket indicates that the two atoms of carbon are also united by one bond of each, thus denoting C to be a tetrad element. From what has just been said with regard to carbon, it is evident that the atomicity of an element is, apparently at least, not a fixed and in- variable quantity; thus, nitrogen is sometimes equivalent to five atoms of hydrogen, as in ammonic chloride (NVH4CI), sometimes to three atoms, as in ammonia and sometimes to only one atom, as in nitrous oxide (ON2). But it is found that this variation in atomicity takes place, with very few exceptions, by the disappearance or develop- ment of an even number of bonds; thus, nitrogen, except in nitric oxide (NO), and dissociated nitric peroxide (NO2), is either a pentad, a triad, or a monad; phosphorus and arsenic, either pentads or triads; carbon and tin, either tetrads or dyads; and sulphur, selenium, and tellurium, either hexads, tetrads, or dyads. These remarkable facts can be explained by a very simple and ob- vious assumption, viz.: That one or more pairs of bonds belonging to the atom of an element can unite, and, having saturated each other, be~ come, as it were, latent. Thus, the pentad element, nitrogen, becomes a triad when one pair of its bonds becomes latent, and a monad, when two pairs, by combination with each other, are, in like manner, rendered latent,—conditions which may be graphically represented thus: Pentad. Triad. Monad. \l/ 1 O • n —N— N— /\ o o And in the case of sulphur Hexad. Tetrad. Dyad. V/ _o_ /\ o u CHEMICAL NOTATION. ATOMICITY. 81 Adopting this hypothesis, it will be convenient to distinguish the maximum number of bonds of an element as its absolute atomicity, the number of bonds united together as its latent atomicity, and the number of bonds actually engaged in linking it with the other elements of a compound as its active atomicity. The sum of the active and latent atomicities of any element must evidently always be equal to the absolute atomicity. Thus in sulphuric acid (Sv,02Ho2) the absolute and active atomicities are both = Yi, therefore the latent atomicity =O. In sulphurous acid ("SlvOHo2) the active atomicity = iy, and consequently the latent = VI IV = il; whilst in sulphuretted hydrogen (IVS//H2) the active and latent atomicities are respectively H and iy. The apparent exceptions to this hypothesis nearly all disappear on investigation. Thus iron, which is a dyad in ferrous compounds (as PeCI2), a tetrad in iron pyrites (FeS"2), and a hexad in ferric acid' (FeO2(OH)2), is apparently a triad in ferric chloride (FeCl3); but the vapor-density of ferric chloride shows that its formula must be doubled —that, in fact, the two atoms of the hypothetical molecule of iron (Fe2) have not been completely separated. The formulae of the ferrous and ferric chlorides and of ferric acid then become Symbolic. Graphic. O Ferrous chloride, . . lvFe"Cl2 Cl—Fe—Cl o o or „/"PeCI, Cl—Fe—Cl \ "Fed/ Cl—Fe—Cl o Cl Cl Ferric chloride, . . . '"Fe'"ClG | | Cl—Fe—Fe—Cl 1 1 Cl Cl Cl Cl f "FeCl3 Cl— fS-fI^-CI 1 ( Cl Cl or \ "Fed/ o 11 Ferric acid, . , Fevl02(OH)2. H- —O—Fe—0—II II 0 _ It will be remarked that the number of bonds supposed to be com- bined with each other in the atom of iron in ferrous chloride is expressed 82 INORGANIC CHEMISTRY. in one of the above formulae by the atomicity numeral iv placed to the left of the symbol, whilst the analogous union of three bonds of each atom of iron in ferric chloride is expressed by the three dashes to the left of the symbol Fe2. These coefficients of latent atomicity will not, however, be used in the case of the single atom of an element, the student being supposed to have made himself acquainted with the absolute atomicity of every element, as expressed in the Table given in Chap. X. For a similar reason it will also rarely be necessary to express the same idea in graphic notation. Thus, for instance, ammonia will be drawn H H and not H—N—H H—N—II It will be necessary, however, to employ these coefficients in sym- bolic formulae where two or more atoms of the same element are joined together under such circumstances that the number of bonds uniting them cannot be found by subtracting the coefficient of active atomicity from the absolute atomicity of the element, as in hydric persulphide (/’S'2H2), for instance, which might otherwise be viewed as "'S'gHg, or VS'2U2. In rare cases, in which oxygen links together two elements or radicals in the same line of a formula, a hyphen is placed before and after the symbol O, thus; / CH2-0-CMeO \ CH2-O-oMeO* Diacetic glycol. Graphic Xotation.—This mode of notation, although far too cumbrous for general use, is invaluable for clearly showing the arrange- ment of the individual atoms of a chemical compound. It is true that it expresses nothing more than the symbolic notation of the same com- pound, if the latter be written and understood as above described ; nevertheless the graphic form affords most important assistance, both in fixing upon the mind the true meaning of symbolic formulae, and also in making comparatively easy of comprehension the probable internal arrangement of the very complex molecules frequently met with both in mineral and organic compounds. It is also of especial value in rendering strikingly evident the causes of isomerism in organic bodies; and it is now almost universally employed by chemists in de- scribing the results of their new discoveries. Graphic notation, like the above method of symbolic notation, is founded essentially upon the doctrine of atomicity, and consists in representing graphically the mode in which every bond in a chemical compound is disposed of. Inasmuch, however, as the principles in- volved are precisely the same as those already described under the heads of symbolic notation and atomicity of elements, it is unneces- sary here to do more than give the following comparative examples of symbolic and graphic formulae : CHEMICAL NOTATION. ATOMICITY. 83 Water, . . . . Symbolic. oh2. Graphic. H—0—H Nitric acid, . . . N02(0H). O II N—O- O II -II Amnionic chloride, nh4ci. H—N—Cl /\ H H Sulphuric anhydride, so3. O II s=o II o Sulphuric acid, . . S02(0H)2. II—0- o II -S—O- II o -II Carbonic anhydride, co2. o= =c=o Potassic carbonate, . CO(OK)2. K—0- -c—o- II o II 1 -K Marsh-gas, , . . ch4. 1 H—C- 1 H -II Ammonic carbonate, CO(ONH4), H H \/ H—N—O- 1 H H H \/ _c—O—N—H 11 1 O H Zincic nitrate, . . no2o I Zn" y no2o j 0 11 N—O- 11 o -Zn—0 O II —N II O It must be carefully borne in mind that these graphic formalfe are intended to represent neither the shape of the molecules, nor the sup- 84 INORGANIC CHEMISTRY. posed relative position of the constituent hypothetical atoms. The lines connecting the different atoms of a compound, and which might with equal propriety be drawn in any other direction, provided they connected together the same elements, serve only to O show the definite disposal of the bonds, the latter again || being only a concrete symbolic expression of an abstract N—O—H train of reasoning; thus the formula for nitric acid indi- ]| cates that two of the three constituent atoms of oxygen are O combined with nitrogen alone, and are consequently united to that element by both their bonds, whilst the third oxygen atom is combined both with nitrogen and hydrogen. The lines connecting the different atoms of a compound are but crude symbols of the bond of union between them; and it is scarcely neces- sary to remark that no such material connections exist, the bonds which actually hold together the constituents of a compound being, as regards their nature, entirely unknown. It deserves also to be here mentioned that graphic, like symbolic formulae, are purely statical representations of chemical compounds: they take no cognizance of the amount of potential energy associated with the different elements. Thus in the formulae for marsh-gas and carbonic anhydride, H lI—C—II o=o=o I H Marsh-gas. Carbonic anhydride. there is no indication that the molecule of the first compound contains a vast store of force, whilst the last is, comparatively, a powerless mole- cule. Calcueation of Formulae.—By quantitative analysis the rela- tive weights of the various constituents of a compound body are dis- covered, and these relative weights are usually expressed in parts per 100. From these numbers the formula of the compound has to be calculated. The percentage composition expresses the relative proportions of the component elements in terms of a common unit; in the formula, the proportion of each element is expressed in terms of its atomic weight. In order, therefore, to ascertain in what proportion of their atomic weights the elements occur in the compound, it is only necessary to divide the proportion of each element in 100 parts of the compound by the atomic weight of that element. Thus the analysis of acetic acid yields the following percentage composition : Carbon, . . In 100 parts. . . 40.00 Hydrogen, . . 6.66 Oxygen, . . . . 53.33 99.99 COMPOUND RADICALS. 85 Dividing each of these numbers by the atomic weight of the element • .. - , 40.00 OOQ 6.66 i 53.33 qqq m question, we find: =3.33; -y— = 6.66; and -yy- = 3.33. Therefore the atomic proportion of carbon : hydrogen : oxygen in acetic acid is as 3.33 : 6.66 : 3.33, or as 1 : 2 : 1. The formula of acetic acid would thus be 0II20. This is, however, only the empirical formula, or smallest possible proportion of the atomic weights. We have already seen (p. 60) that the molecular formula of acetic acid is C 2H402, or twice as great as the above. CHAPTER IX. COMPOUND RADICALS. The term compound radical may be applied to any group of two or more atoms, which takes the place and performs the functions of an element in a chemical compound. In practice, however, it is only applied to any such group when met with in numerous chemical com- pounds. An element is a simple radical, and enters into combination in the following manner, a, b, c, and d being monad elements, a" a dyad, a 7" a triad, and aiv a tetrad element: a' -f- b = ab, a" + 26 = a"b2, a'" + 3 b = a"'b3, etc. etc. A group of elements replacing a, a", or a'" in the above equations is a compound radical, as in the following examples: {a"b) + h = (a"b)b, (.a"'b)" +2 b = {a’"b)"h.2, (■a'"bc) -f- b (a"'bc)b, {aivb)'" + 36 = {aivb)'"bs, \aiybo)" + 2b = {aiybG)"b2, -- 6 = (o)ybed)b. The group of elements (a"b) constitutes a compound monad radical equivalent to one atom of hydrogen or chlorine. The group (a"'b)" is a compound dyad radical, etc. It is therefore evident that a polyad element is essential to every compound radical; in fact a compound radical consists of one or more atoms of a polyad element in which one or more bonds are unsatisfied; and it is either a monad, dyad, triad, etc., radical, according to the number of monad atoms required to satisfy its active atomicity. Such a radical, when a monad, triad, or pentad, can- not exist as a separate group : like hydrogen or nitrogen, when isolated, it combines with itself, forming a duplex molecule. It is only by the 86 INORGANIC CHEMISTRY. union of two atoms or groups of atoms that the vacated bonds can in these cases be satisfied. From the above definition of a compound radical, it is evident that an almost infinite number of such bodies must exist; for in the com- pounds of every polyad element it is only necessary to vacate successive bonds to create each time a new compound radical. Thus marsh-gas CH4 minus one atom of hydrogen gives the compound radical methyl CH3; minus two atoms of hydrogen, it forms methylene (CH2),/; and by the abstraction of three hydrogen atoms it is transformed into the triad radical formyl (CH)'"; but, except in a few cases, it is not ad- vantageous thus to incorporate, as it were, compound radicals, which, instead of simplifying notation and nomenclature, would, if thus multi- plied, only embarrass them. No compound radical, therefore, ought to receive recognition as such, unless it can be shown to enter into the composition of a large number of compounds. The following are the names, symbols, and formulae of the inorganic compound radicals recognized in the notation of this volume: Hydroxyl, . . . Molecular formulae. • (OH)a Semimoleeular formulae. OH Semimoleeular symbols. Ho. Hydrosulphyl,. . . (8H)2 SH Hs. Ammonium, . (NH4)2 nh4 Am. Ammonoxyl, . . . (ONH4]2 onh4 Amo. Amidogen, . . . • (NH2)2 nh2 Ad. In addition to these, certain compounds which metals form with oxygen are also regarded as compound radicals—for instance, Molecular Semi molecular Semimolecular formulae. formulae. symbols. Potassoxy], . . (OK)2 OK Ko. f° Zincoxyl, . . . (02Zn) •< Zn" Zno". 0 The essential character of these last compound radicals is that the whole of the oxygen they contain is united with the metal by one bond only of each oxygen atom, as seen in the following graphic formulae : Hydroxyl, —O—H Potassoxyl, —O—K Zincoxyl, —O—Zn—O— The metal thus becomes linked to other elements by these dyad atoms of oxygen. The functions of such compound radicals will be sufficiently evident from the following examples of compounds into which they enter, and in which their position is marked by dotted lines. O : : II | I Nitric acidr . . . N-j-O—H j O I. i COMPOUND RADICALS 87 r:O : : I MM I Potassic sulphate, . K—O-i-S-.-O K I MM I O i i Or ■ O ii i i ii Baric nitrate, . . N-i-0—Ba—O-i-N ii i I ii O i. .i O 0 i' /0\ i 11 A \ i Zincic sulphate,. . Sq; bZn i Alv I It is not necessary to dignify all these metallic compound radicals with names; the chief point of importance about them is their abbrevi- ated notation, in which the small letter o is attached to the symbol of the metal, the atomicity of the radical being marked in the usual manner. Although the small letter o in these symbols of combining quantities has no more reference to the composition of the radical than the d in the corresponding symbol of amidogen, yet it may usefully remind the reader that oxygen is always a constituent of the compound radicals so symbolized. It must be borne in mind that the number of atomsof oxygen in any radical of .this class depends upon its atomicity: thus a monad contains only one atom of oxygen, a dyad two, and a triad always three atoms of oxygen. The use of any but monad and dyad metallic compound radicals is very rare. It is also in some cases convenient to recognize as a radical the atomic group which remains when all the hydroxyl is abstracted from an oxyacid, as for instance : Acid. Acid radical. Nitrous acid, . . . NOHo Nitrosyl, . . . . (NO) Nitric acid, . N02Ho Nitroxyl, , . . . (NO.) Sulphuric acid, . S02Ho2 Sulphury], . . , (S02)" Phosphoric acid, . . POHo3 Phosphoryl, (PO)'" It is evident that the atomicity of these elements must be the same as the basicity of the acids from which they are derived. Atomic and Molecular Combination. In all the cases of chemical combination already considered, a union of atoms has been invariably contemplated. This atomic union is gen- erally attended by the breaking up of previously existing molecules— two such molecules, by the mutual exchange of their atomic constitu- ents, producing two new and perfectly distinct molecules. Thus, when chlorine unites with hydrogen to form hydrochloric acid, a molecule of 88 INORGANIC CHEMISTRY. chlorine and one of hydrogen yield up their constituent atoms, forming two molecules of hydrochloric acid, Cl2 + F2 = 2HCI. In comparatively rare cases, two molecules combine to form only one new molecule; thus a molecule of carbonic oxide and one of chlorine combine to form one melecule of carbonic oxydichloride or phosgene gas: but the union is even here essentially atomic; for after combination both the oxygen and chlorine are directly united with the atom of carbon: 0"O + Cl2 = OivOCl2. Carbonic oxide. Chlorine. Phosgene gas. Chemists are, however, compelled to admit an entirely different kind of union, which not unfrequently occurs, and which in conformity with the atomic hypothesis, may be appropriately termed molecular union or mole- cular combination. In the formation of such compounds, no change takes place in the active atomicity of any of the molecules. It is this kind of combination which holds together salts and their water of crystalliza- tion, as, for instance, Sodic chloride crystallized at — 10° C., . . NaCl,20H2. Sodic bromide crystallized below -J- 30° C., . NaBr,20H2. Sodic iodide crystallized below + 50° C., . . NaI,20H2. Alum; • s4o8('Ai"' 2O6)viKo2,240H2. Numerous other instances of molecular combination might be adduced ; but it is only necessary here to point out that such molecular unions will be distinguished from atomic combinations by the use of the comma, as in the above and following examples : Tetramethylammonic tri-iodide, . NMe4I,T2. Tetramethylammonic pentiodide, . . NMe4I,2I2< Tetramethylammonic iodo-dichloride,. . NMe4I,Cl2, In all cases molecular combination seems to be of a much more feeble character than atomic union; for, in the first place, such bodies are generally decomposed with facility; and secondly, the properties of their constituent molecules are markedly perceptible in the compounds. Thus the above periodides of the organic bases greatly resemble iodine in appearance. CHAPTER X. CLASSIFICATION OF ELEMENTS. It has been already mentioned that the elements may be divided into two great classes, the metals and the non-metals or metalloids. A second division into positive and negative elements has also been ex- plained. A third and still more important classification is founded upon the atomicity of the elements. In the following classified table, all three methods are embodied, the names of the metalloids being printed in heavy type, and those of the metals in common type, whilst the names of the positive elements are printed in Roman characters, 89 CLASSIFICATION OF ELEMENTS. and those of the negative in italics. In addition, the different classes are also divided into sections, consisting of elements closely related in their chemical characters. / Monads. Dyads. Triads. Tetrads. Pentads. Hexads. Heptads. Octads. 1st Section. Hydrogen. 1st Section. Oxygen. 1st Section, lloron. 1st Section. Clarion. Silicon. Titanium. Zirconium. Tin. Thorium. 1st Section. Nitrogen. Phosphor’s Vanadium, Arsenic. Niobium, Antimony. Tantalum. Bismuth. 1st Section. Sulphur. Selenium. Tellurium. Chlorine 1 Bromine v (?) Iodine. J * Ruthenium. Osmium. 2d Section. Fluorine. Chlorine. 1 Bromine. V (?) Iodine, j * 2d Section. , Barium. Strontium. Calcium. Magnesium, Zinc. Beryllium. 2d Section. Gold. 2d Section. Uranium. Tungsten. Molybdenum. 3d Section. Thallium. Indium. 2d Section, Gallium. Aluminium. 2d Section. Didymium. 3d Section. Caesium. Rubidium. Potassium. Sodium. Lithium. 3d Section. Chromium. Manganese. Iron. Cobalt. Nickel. 3d Section. Cadmium. Mercury. Copper. 4th Section. Lanthanum. Yttrium. Erbium. Decipium.l Samarium. 1(?) Scandium. J 3d Section. Cerium. 4th Section. Platinum. Iridium. Palladium. Rhodium. 4th Section. Silver. 5th Section. Lead. * Chlorine, bromine, and iodine have been treated as monadic in the present work; but in the opinion of some chemists these elements are heptadic (see “Periodates”). 90 INORGANIC CHEMISTRY. Classification of the Elements according to their Atomic Weights.—The Periodic Law.—The idea of a possible connection between the atomic weights of the elements and their properties was first suggested by the observation that in many cases similar elements could be arranged in groups of three, in which the atomic weight of the intermediate element was approximately the arithmetical mean of the atomic weights of the highest and lowest. Examples of such groups, which were termed “ triads,” are P = 31, As = 75, Sb = 120 :{1 ; = 75.5. Cl = 35.5, Br = 80, I = 127 355 + 127 = 81.25. Ca = 40, Sr = 87.6, Ba = 137 137 = 88.5. The most complete expression of these relations that has yet been pro- posed is to be found in the “ periodic law of the elements.” The fact that the properties of the elements vary periodically with their atomic weights was first shown by Newlands in 1864.* More complete and systematic expressions of the same law were published a few years later by Mendeleef and by Lothar Meyer. The most precise of these systems is that of Mendeleef, which has lately attracted much attention on account of the number of new facts which it has enabled its author to predict. The following is a brief outline of the method followed by Mendeleef. If all the elements whose atomic weights lie between 7 and 35.5 be arranged in the arithmetical order of their atomic weights, thus: Li= 7; Be = 9.4; B= 11; C = 12; N = 14; O = 16; F = 19 ; Na= 23; Mg = 24; A 1 = 27.3; Si= 28; P = 31;5=32;C1 = 36.5, certain definite relations may be perceived. The character of the ele- ments is here seen to be subject to regular modification, so that, step by * Newlands was the first to point out that the elements, when arranged in the arithmetical order of their atomic weights, exhibit a periodic recurrence of similar properties. He stated that each such period consists of seven -elements, and that, with the eighth element, properties resembling those of the first recur. To this relation he gave the name of the Law of Octaves, comparing the periods of recurrence with the oc- taves of the musical scale, and the elements within the period with the notes included in the octave. Newlands’s system is therefore in all essential points identical with that of Mendeleef, which was published in 1869; except that Newlands failed to recognize the existence of the “transitional elements”—Mendeleef’s eighth group (see table, p. 92)—which divide the other elements into groups of two octaves each. The fact that Mendeleef’s table, published five years later than the first table given by Newlands, is undoubtedly more perfect in its details, has led some chemists to as- cribe the discovery of the periodic law to the former investigator. This is manifestly unjust. The credit of originating an idea is due solely to him who first formulates it, and this is irrespective of any subsequent development which the idea may undergoat the hands of others, provided that the central idea itself remains unaltered. No one, for example, has ever suggested that the authorship of the modern atomic theory is to be ascribed to Cannizzaro instead of to Dalton, because the rectification of the atomic weights was the work of the former chemist. CLASSIFICATION OF ELEMENTS. 91 step, as the atomic weights vary, the characters of the elements also vary, and by comparing the series of elements from Li to F with the series from Na to Cl, it is manifest that this variation is a periodic one, the same changes of character which are met with in traversing the first series, being again found in the second series: thus Li corresponds to Na, Be to Mg, Bto Al,* etc. The regularity of the change in trav- ersing a period may be seen by comparing with each other the oxides of one such series of elements, writing these so as to show the relative quantities of oxygen with which the same number of atoms of the va- rious elements combine, instead of employing the molecular formulae of the oxides: Na20; Mg202; A1203; Si204; P2Og; S2Og; C]207.f (MgO) (Sio2) (SO3) Here the proportion of oxygen in the various oxides throughout the period is as 1:2:3:4 : 5 : 6 : 7, At the same time there is a regu- lar gradation from left to right from the most electropositive element, through the various intermediate stages, to the most electronegative element. This periodic recurrence of the same properties with the gradual increase of the atomic weight has been formulated by Mende- leef thus : The properties of the elements are a periodic function of their atomic weights. Following out this principle, Mendeleef has tabulated the whole of the elements on the same plan (see table, page 92). The Roman numerals indicate the groups or families of similar ele- ments, which are thus arranged in vertical columns; the Arabic nu- merals refer to the series or periods, which are arranged horizontally. As regards the latter, it is to be noted that there are two kinds of pe- riods—the one following the even Arabic numerals, the other the odd. If we confine our attention to a single group, we find that the elements of the even periods correspond with each other in their properties, and that the elements of the odd periods likewise correspond with each other, but that there is less correspondence of the members of one of these classes with those of the other. Thus, in Group 11., the corre- sponding elements of the even series are Be, Ca, Sr, and Ba; of the odd series, Mg, Zn, Cd, and Hg. The series 2 and 3 are termed by Mendeleef “short periods”; the remaining series are grouped together in pairs—thus, 4 and 5,6 and 7, 8 and 9, etc.—the two series of such a pair together constituting a “ long period.” That is to say, if we traverse the series 3we find a periodic repetition of the chemical characteristics already met with in series 2; but in order to meet with a similar periodic change of char- acteristics—e.g., in order to pass from a highly electropositive to a highly electronegative element—it is necessary to traverse the entire double- series 4 and 5, and again the double-series 6 and 7, and so on. The full significance of this arrangement—at first sight, perhaps a somewhat arbitrary one—will be shown further on. * On this supposition Al would have to be regarded as triadic. This would he in harmony with the observed vapor-density of aluminic methide, AI(CPI3)3, at 240°. t Perchloric anhydride is not known; but the corresponding acid has been pre- pared. INORGANIC CHEMISTRY. Groups: I. II. III. IV. V. VI. VI I. VIII. rh4 eh3 eh2 EH (B2H) E20 E202 E203 E204 ba B206 e2o7 (BA) 1 1 H _ 2 Li 7 Be 9 B 11 C 12 N 14 0 16 F 19 3 23 Na 24.4 Mg 27 A1 28.2 Si 31 P 32 S 35.5 Cl 4 K 39 Ca 40 Sc 44 Ti 48 V 51.3 Cr 52 Mn 55 Fe 56, Co 58.6, Ni 58.6 5 63.2 Cu 65.3 Zn 68.8 Ga 72? 75 As 79 Se 80 Br 6 Rb 85.3 Sr 87.5 Y 89.8 ? ' Zr 90 Nb 94 Mo 95.5 ?100 Eu 104, Eh 104, Pd 105.7 7 107.7 Ag 112 Cd 113.4 In 118 Sn 120 Sb 125 Te 127 I 8 Cs 133 Ba 137 La 138.5 Ce 140.5 Di 146 Tb? 148.8 Sm? 150 ? 152, ? 153, ? 154 9 156? 158? 159 Dp? 162? 165.9 Er? 167 ? 169? 10 ? 170 ? 172 Yb 172.8 ? 177 Ta 182 W 184 ? 190 Os 198.6?, Ir 192.5, Pt 194.4 11 196 Au 200 ITg 204 TI 206.5 Pb 208.2 Bi 214 Ng? 219? 12 ?221 ? 225 ? 230 Th 233.4 ? 237 U 238.5 ? 244 The Periodic System of the Elements. CLASSIFICATION OF ELEMENTS. 93 In passing from the left to the right there is in every series, taking each group in that series in succession, a gradual increase in the quan- tity of oxygen with which the elements can unite. The members of the different groups taken in order exhibit a regular change (generally an increase) of atomicity, odd and even atomicities alternating. Group VIII. is anomalous. In this group there are always three elements in each series, instead of, as in the other groups, only one element. These elements of Group VIII. do not, when taken in any series in the order of their atomic weights, exhibit the above alternation of odd and even atomicity: they are all even; but their atomicity decreases with a rise of atomic weight. They are termed by Mendeleef “ trans- itional elements,” and their place is between the even and the odd se- ries of a long period. This transitional group will be referred to again later on. The grouping together of sodium, silver, and copper as similar ele- ments is justified by the isomorphism of some of the cuprous and ar- gentic compounds, and of some of the latter again with the corresponding sodium compounds. Mendeleef has employed this periodic law in the correction of doubt- ful atomic weights, and in the prediction of undiscovered elements. Thus, indium was formerly believed to be a dyad with the atomic weight 76, and its oxide was therefore supposed to possess the formula InO. With this atomic weight, it would take its jDlace between arsenic and selenium. But there is no vacant space for it in this part of the table, and it would, moreover, have no analogy with the elements with which it would have to be grouped. Mendeleef pointed out that by assuming indie oxide to possess the formula ln203, with an atomic weight for the metal of 114, indium would take its place in series 7 between cadmium and tin, and as an analogue of aluminium. The cor- rectness of this view has been demonstrated by the determination of the specific heat of indium by Bunsen. Again, chemists were uncertain whether uranium had the atomic weight 60 or 120. Mendeleef showed that no element of either of these atomic weights and of the properties of uranium would find a fitting place in the table, but that by assigning to it the atomic weight 240 (238.5), it would take its place as an analogue of chromium, mo- lybdenum, and tungsten. This change has been justified by the results of the determination of the specific heat of uranium and by the vapor- density of various uranium compounds. Again, the determinations of the atomic weight of molydenum left it uncertain whether this element possessed the atomic weight 92 or 96. The former of these weights would place it before niobium, and in a group of elements with which it presents no analogy. In order that it might take its place in Group VI. as an analogue of chromium, its atomic weight must be higher than 94, the atomic weight of niobium. A careful determination has in fact shown that the atomic weight of molybdenum is 95.5. Again, tellurium was supposed to have the atomic weight 128. In order that it might take its place in the same group as its chemical analogues sulphur and selenium, it was necessary that its atomic weight 94 INORGANIC CHEMISTRY. should be lower than 127, the atomic weight of iodine. A recent de- termination by improved methods has shown that the atomic weight of tellurium is 125. It will be noticed that in the foregoing table one element, osmium, has been placed in a position different from that indicated by its atomic weight as at present determined. Osmium from its properties ought to have an atomic weight lower than that of iridium, instead of higher than that of gold. It remains to be seen whether experiment will, as in the preceding cases, verify this prediction. Mendeleef has shown that the properties, both chemical and physical, of an element may be to a certain extent predicted from the properties of what he terras its “ atomic analogues.” By this term he understands not its chemical analogues, but the two elements which stand on either side of it in the same series, together with the two elements which stand above and below it in the same group. Thus As, Br, S, and Te are the atomic analogues of Se. It will be observed that there are in the table a number of gaps. These correspond, according to Mendeleef, with elements which have not yet been discovered. If such a gap is surrounded by the requisite atomic analogues, it is possible to predict the properties of the unknown element. Thus in the positions 111. 4, 111. 5, and IV. 5, Mendeleef placed three unknown elements to which he gave the names elcaboron, ekaluminium, and ehasilicon—following a system of nomenclature which he has devised for the designation of such unknown elements and which, while referring these to known elements of the same group, distinguishes them by prefixing the Sanscrit numerals elect, dvi, tri, etc., according to their position in the group. Concerning ekaluminium, he states that it has an atomic weight of about 68, and a specific gravity of about 6.0, and that it forms a sesquioxide. These predictions were verified by the discovery of gallium, which has an atomic weight of 68.8, a specific gravity of 5.9, and forms an oxide of the formula Ga203. The new metal scandium is possibly Mendeleef’s ekaborou. The above prediction of the specific gravity of ekaluminium (gallium) is rendered possible by the fact that the physical as well as the chemical properties of the elements are periodic functions of the atomic weight. This may be illustrated by reference to the magnetic properties of the elements. Faraday divided all substances into two classes: those which are attracted by a magnet, or paramagnetic bodies, and those which are repelled by a magnet or diamagnetic bodies. In the case of the elements, the magnetism of the following has been determined; Paramagnetic Elements. K, C, Ti, Ce, N, O, Cr, U, Mn, Fe, Co, Ni, Rh, Pd, Os, Ir, Pt. Diamagnetic Elements. H, ISTa, Cu, Ag, An, Zn, Cd, Hg, Tl, Si, Sn, Pb, P, As, Sb, Bi, S, Se, Cl, Br, I, An inspection of these two classes does not reveal any apparent con- nection between the chemical and the magnetic properties of the ele- CLASSIFICATION OF ELEMENTS. 95 merits. Thus we find that elements, chemically so closely related as potassium and sodium, oxygen and sulphur, nitrogen and phosphorus, titanium and silicon, are separated in the two classes. Camel ley has, however, pointed out that the paramagnetic elements are, without ex- ception, to be found in the even series of Mendeleef’s table and the dia- magnetic elements without exception in the odd series. Further, the paramagnetic power of the members of a paramagnetic group of elements (thus Fe, Co, Ni) diminishes, and the diamagnetic power of the mem- bers of a diamagnetic group of elements (thus P, Sb, Bi, or H, Cu, Ag, Au) increases, with increasing atomic weight. The fact that the physical properties of the elements are a periodic function of their atomic weights is, however, most strikingly shown by the curve given in the annexed diagram. This curve, which is in reality a graphic expression of the periodic law, was first constructed by Lothar Meyer. It is given here as supplementing in a remarkable manner Mendeleef’s table. In this curve the abscissae represent the atomic weights, and the ordinates the atomic volumes of the various elements in the solid state.* The curve is therefore primarily a graphic representation of the varia- tion of the atomic volume with the atomic weight. But a brief inspec- tion shows that it is much more than this. In the first place then, as regards the atomic volume, the curve shows m the plainest manner that this varies periodically with the atomic weight: at one point it reaches a maximum, then gradually decreases with increasing atomic weight till it falls to a minimum, again rising to a maximum, and so on. Each of these compound periods of decrease and increase corresponds with one hollow of the wave of the curve ex- tending from crest to crest. A comparison of this curve with Mende- leef’s table is highly instructive, especially when we consider that the two were constructed quite independently of each other. In the curve the periods of change of atomic volume—the hollows—are distinguished by Roman numerals. Periods 11. and 111. of the curve correspond with Mendeleef’s two “ short periods,” series 2 and 3 of the table. The large hollows of the curve, IV., V., etc., correspond with Mendeleef’s “long periods thus Period IV. of the curve is the “ long period ” made up of series 4 and sof the table; Period V. is the “ long period ” made up °f series 6 and 7of the table, and so on. (The latter part of the curve has not been finished for want of data.) The alkali metals with which Meudeleef’s periods commence are always found at the maxima of the curve. Mendeleef’s “ transitional elements ” of Group VIII., the metals which lie between the even and odd series of a “ long period,” are always bound at the minima of the large hollows. Osmium cannot, with its X V 1 atomic volumes of the elements are the relative volumes occupied by atomic quantities, i.e., quantities taken in the proportion of the atomic weights. These atomic volumes may be found by dividing the atomic weights of the elements by their specific gravities (see following chapter). In the diagram, wherever the elements are not Known in the solid state, the hypothetical course of the curve is represented by a dotted jue. As regards the rather irregular course of the curve in some parts, it is to be noted hat the specific gravities of the elements have not always been determined under strictly comparable conditions. Thus the specific gravity of potassium is determined a tew degrees below its fusing-point; that of platinum about 2000° below the fusing- 96 INORGANIC CHEMISTRY. present atomic weight, be made to fit into this curve, any more than into MendeleeLs table. Various other periodic relations between the atomic weights and the physical properties of the elements have been indicated on the diagram by appending to each part of the curve a list of the physical properties of the elements to which that part refers. Thus, elements possessing the same physical properties are to be found in corresponding parts of the curve. It is to be noted, however, that the alternation of “ electro- positive—electronegative,” which occurs only once in Period 11. and only once in Period 111. of the curve, occurs twice in Period IV. and twice in Period V. This is in harmony with the fact already referred to that Periods 11. and 111. correspond each with one series of Mende- leef’s table; Periods IV. and V. each with two series. It is quite inconceivable that the remarkable relationships expressed by the periodic law should be a work of chance. No explanation of the periodic law has yet been offered. At present it is an empirical law, established by careful experiment and compari- son. It stands in the same relation to chemistry as did the laws of Kepler to astronomy before the time of Newton. Its explanation will in all probability constitute the chemical theory of the future. CHAPTER XL RELATIONS BETWEEN CHEMICAL COMPOSITION AND SPECIFIC GRAVITY. ATOMIC VOLUME. The relative volumes which atomic or molecular quantities (quanti- ties taken in the proportion of the atomic or molecular weights) of sub- stances occupy, may be found by dividing the atomic or molecular weights of these substances by their specific gravities. The quotients thus obtained are termed atomic volumes and molecular volumes respect- ively. It must not be supposed that these quotients express the relative volumes occupied by the atoms or molecules. In the gaseous state the molecules are separated from each other by distances which are enor- mously great compared with the diameters of the molecules themselves. In the solid and liquid states, the atomic volumes could only represent the relative volumes of the atoms, provided that the spaces between the atoms were in every case proportional to the size of the atoms—an assumption for which there is not the slightest ground. The atomic volumes, therefore, represent the relative volumes of the atoms, plus the relative volumes of their interstitial spaces. The molecular volumes of gases have already been treated of (p. 54), and may be dismissed in a few words. As the specific gravities or vapor-densities of gaseous bodies are proportional to ~ . , , .i , ... molecular weight „ their molecular weights, the quotient = A5— will in all vapor-density RELATIONS BETWEEN CHEMICAL COMPOSITION, ETC. 97 cases possess the same value. The value of this quotient is either 28.9 or 2, according as the vapor-density is referred to air or to hydrogen (see p. 53). The laws which govern the relations between composition and specific gravity are less simple in the case of solids and liquids; but here also very striking regularities are manifested. As the specific gravity of a solid or liquid denotes the weight in grams of one cubic centimetre of the substance, so the atomic or mole- cular volume, if the atomic or molecular weight be expressed in grams, will represent cubic centimetres. The atomic weight of sulphur is 32, its specific gravity 2. The atomic weight of lead is 206.5, its specific gravity 11.37. The atomic volume of sulphur is therefore 16, that of lead 18.2. In grams and cubic centimetres this may be expressed as follows: If 2 grams of sulphur occupy the volume of 1 c.c,, 32 grams will occupy 16 c.c. If 11.37 grams of lead occupy the volume of 1 c.c., 206.5 grams will occupy 18.2 c.c. Among the elements, the various members of an isomorphous group frequently exhibit approximate equality of atomic volume. Atomic weight. Specific gravity. Atomic volume. Iron, 56 7.79 7.2 Cobalt, 58.6 8.60 6.8 Copper, 63.2 8.95 7.1 Manganese, 55 8.00 6.9 Nickel, 58.6 8.90 6.6 Again : Iridium, 192.5 22.38 8.6 Palladium, 105.7 11.40 9.2 Platinum, 194.4 21.53 9.0 Rhodium, 104 12.10 8.6 The members of an isoraorphous group of compounds generally have approximately the same equivalent volume. In the group of the spi- uelles, which crystallize in forms of the regular system, these relations are as follows: Molecular ■weight. Specific gravity. Molecular volume. Mg0,Al808, 142.4 3.45 41.3 ZnO,Al2G3 183.3 4.58 40.0 MnO,Cr203, 224 4.87 46.0 Zn0,Cr203, . 233.3 5.31 43.9 Zn0,Fe203 241.3 6.13 47.0 FeO,Fe20 232 5.09 45.6 98 INORGANIC CHEMISTRY. The subject of atomic and equivalent volumes of solids has been in- vestigated by H. Kopp, Schroder, and others. The molecular volumes of liquids, when compared at the same tem- perature, display no regularities. If, however, these volumes be deter- mined at temperatures at which the tensions of the vapors of the liquids are equal, that is to say, at temperatures at which the energy of the molecules which fly off from the surface of each liquid is equal, and at which temperatures consequently the liquids are in the same condition as regards the weakening of the force of cohesion, important laws be- come manifest. Under such conditions, it seems that each element has one or more fixed atomic volumes, and that the molecular volume of a compound in the liquid state is the sum of the atomic volumes of its elements. As the vapor-tensions of most liquids have not been deter- mined for a variety of temperatures, it is usual to compare the mole- cular volumes at the boiling-points of the liquids, at which tempera- tures the tensions of their vapors are equal to the normal atmospheric pressure (see p. 120). These laws may be deduced and expressed as follows: 1. A difference of n.CH2 in the formula of liquid compounds corre- sponds to a difference of n.22 in the molecular volume. Thus, ruethy lie formate (C21T402), methylic acetate (C3H602), ethylic acetate (04H8O2), and methylic butyrate (C51110O2), whose formulae differ by CH2, differ in molecular volume by nearly 22. (For a comparison of the experi- mental with the calculated results, see table, p. 101.) 2. Isomeric liquids, belonging to the same chemical type, such as acids and ethereal salts, alcohols and ethers, ketones and aldehydes, have the same molecular volume. Thus, the molecular volumes of propionic acid, ethylic formate and methylic acetate, all of which have the formula C 3H604, closely approximate to 86. 3. The substitution of one atom of oxygen for two of hydrogen causes a slight increase of molecular volume. The molecular volume of alco- hol (C2II0O) is between 61.8 and 62.5, that of acetic acid (C2H4G2) lies between 63.5 and 63.8. Cyraene (O10Hu) and cuminaldehyde (O10H12O) differ similarly in their molecular volumes. 4. In two liquids belonging to the same chemical type, the substitu- tion of one atom of carbon for two atoms of hydrogen produces no change of molecular volume. This may be seen in the case of ethylic benzoate (C9II10O2) and ethylic valerate (C7H1402); benzaldehyde (C7H6O) and valeraldehyde (C6II10O); cymene (C 1 and butvl (Ceuis). As the addition of CII2 to the formula of a compound produces an increase of 22 in the equivalent volume (Law 1), this number may be supposed to represent the equivalent volume of CII2. And since (Law 4) the exchange of C for II2 causes no change of molecular volume, the atomic volume of C may be taken to be equal to that of H2. Hence, 22 the atomic volume of Cis equal to =ll, and that of II2 is also equal to 11, or that of H = 5.5. From the increase in molecular volume which the substitution of O for H2 causes, the atomic volume of O may be calculated to be equal to 12.2. In this case when Ois substituted EELATIONS BETWEEN CHEMICAL COMPOSITION, ETC. 99 for II2, both its bonds are attached to the same atom of carbon, as for example when alcohol is converted into acetic acid. H H HO II I II H—C—C—O—H H—-C—C—O—II I I I H H H Alcohol. Acetic acid. It will be convenient, in discussing the subject of atomic volumes, to represent oxygen thus attached by the ordinary symbol O, whereas oxygen which serves to unite two elements or groups of elements, as in the case ot‘ hydroxylic oxygen, or of oxygen in ethylic oxide, will be distinguished by the symbol lt is found that the atomic volume of ois different from that of O. The value of the former may be de- duced from the molecular volume of water. Molecular > j ol u me of 0H2 = 18.8 Atomic U h2 = 11 u u 0 7.8 From these four atomic volumes, Atomic volume of C = = 11 U U H = = 5.5 C( u O = = 12.2 u u 0 - = 7.8 the molecular volumes of compounds containing only these four ele- ments may be calculated. The numbers so deduced approximate very closely to those obtained by experiment. It is evident that the value to be assigned to the atomic volume of oxygen will depend upon the constitution of the compound, and that, conversely, the molecular vol- ume of a compound containing oxygen will afford a means of ascer- taining the part which this element plays in its constitution. A few examples will suffice. The graphic formula of acetone is H H I I H—C—C—C—H. i II 1 H O H From this formula follows : Atomic volume of C, = 33 a a H == XI Q 33 “ “ o = 12.2 Molecular volume of acetone = 78.2 100 INORGANIC CHEMISTRY. The molecular volume of acetone as determined by experiment is between 77.3 and 77.6. The graphic formula of alcohol has been given on p. 99. The molecular volume would be calculated thus: Atomic volume of C2 = 22 “ “ h6 = 33 “ “ © = 7.8 Molecular volume of alcohol = 62.8 The observed volume is between 61.8 and 62.5. The graphic formula of acetic acid has /been given on p. 99. Its molecular volume would be as follows: Atomic volume of C2 — 22 u u h4 = v 22 U U o = 12.2 U U © = 7.8 Molecular volume of acetic acid = 64.0 The experimental value is between 63.5 and 63.8. The subject of the molecular volumes of liquids has been investigated chiefly by H. Kopp, to whom the enunciation of the above laws is due,* The following table contains a list of his determinations of molecular volumes at the boiling-point for a number of liquids into the composi- tion of which only carbon, hydrogen, and oxygen enter. The third column contains the temperatures at which the determinations were made. * Recently the subject has been studied by Thorpe, Ramsay, and others. RELATIONS BETWEEN CHEMICAL COMPOSITION, ETC. Molecular Volumes of Liquids containing Carbon, Hydrogen, and Oxygen. Substance. Formula. Temperature. Molecular volume. Observed. Calculated Benzene, .... C6e6 80° C. 176°F. 96 0— 99.7 99.0 Cymene, .... C10HU 175 347 183.5-185.2 187.0 Naphthalene, . . c10h8 218 424 149.2 154.0 Aldehyde, . . . C,H,0 21 70 56.0— 56.9 56.2 Valeraldehyde,. . c5h10o 101 114 117.3—120.3 122 2 Benzaldehyde, . . c7b6o 179 354 118.4 122.2 Cuminaldehyde, . c10h12o 236 457 189.2 188.2 Butyl, c8h18 108 226 184.5—186.6 187.0 Acetone, .... c3h60 66 133 77.3— 77.6 78.2 Water, ©h2 100 212 18.8 18.8 Methylic alcohol, . ch4© 59 138 41.9— 42.2 4H.8 Ethylic “ c2h6© 78 172 61.8— 62.5 62.8 Amvlic “ C5HJ2© 135 275 123.6—124.4 128.8 Phenol C6H6© 194 381 103.6—104.0 106.8 Benzvlic alcohol, . . C7H8® 213 415 123.7 128.8 Formic acid, . . ch2o® 99 210 40.9— 41.8 42.0 Acetic “ . . c2h4o© 118 244 63.5— 63.8 64.0 Propionic “ . . C3H6( )© 137 279 85.4 86.0 Butyric “ . . c4h8o© 156 313 106.4—107.8 108.0 Valeric “ . . c5h10o® 175 347 130,2—131.2 130.0 Benzoic “ . . C7Hfin© 253 487 126.9 130.0 Ethvlic oxide, . . c4H]0© 34 93 105.6—106.4 106.8 Acetic anhydride, . C4Hfi(J2© 138 280 109.9—110.1 109.2 Methylic formate, . C2H4G© 36 97 63.4 64.0 Methylic acetate, . c.,n6o© 55 131 83.7— 85.8 86 0 Ethylic formate, , c3h6o© 55 131 84;9— 85.7 86.0 Ethylic acetate, c4n8o© 74 165 107.4—107.8 108 0 Methylic butyrate,. c5H10o® 93 199 125.7—127.3 130.0 Ethvlic propionate, c5h10o® 93 199 125 8 130.0 Methylic valerate,. c6h120© 112 234 148.7—149 6 152.0 Ethylic butyrate, . c6h12o© 112 234 149.1—149.4 152.0 Butvlic acetate, . . c6h120© 112 234 149.3 152.0 Amvlic formate, c6h12o® 112 234 149.4—150 2 1520 Ethvlic valerate, . c7h14o© 131 268 173.5—173.6 174.0 Amylic acetate,. . c7h14o® 131 268 173.3—175.5 174.0 Amvlic valerate, . c10h20o© 188 370 244.1 240.0 Methylic benzoate, c8it8o® 190 374 148,5—150.3 152 0 Ethvlic “ c9h10o© 209 ' 408 172.4—174.8 1740 Amvlic “ c12h16o© 266 511 247.7 240.0 Ethvlic cinnamate, c„h120© 260 500 211.3 207.0 Methylic salicvlate, 223 433 156.2—157.0 159.8 Ethylic carbonate, c5h10o©2 126 259 138.8—139.4 137.8 Methvlic oxalate, . c4h6o2©2 162 324 116.3 117.0 Ethvlic “ c6h10o2©2 186 367 166.8—167.1 161.0 Ethylic succinate, . c8h14o2©2 217 423 209.0 205.0 In like manner, from the molecular volumes of the liquid chlorides, bromides, and iodides, the atomic volume of Cl has been determined to be equal to 22.8, that of Br 27.8, and that of I = 37.5. Elements of varying atomicity like nitrogen and sulphur seem to follow some less simple law. It is possible that the atomic volumes of these elements may vary in some way with their atomicity; but the pre- cise nature of this variation has not been ascertained. The subject requires thorough investigation by the light of modern constitutional formulas. 102 INORGANIC CHEMISTRY. CHAPTER XII. CHEMICAL AFFINITY. Chemical affinity has been referred to at some length in the opening pages of this introduction. It may be measured as regards its extent and as regards its intensify. A measure of the relative extent of the chemical affinity of two or more elements for some other element is af- forded by the number of atoms of this element with which each can combine. Extent of affinity is thus directly connected with atomicity. Relative intensity of affinity of two or more elements for any given ele- ment refers to the resistance which their compounds with this element offer to decomposition. The measure of this intensity is the quantity of heat evolved in combination or required for decomposition. Extent and intensity of affinity are quite independent of each other. Thus copper and mercury in the compounds CuO and HgO have the same extent of affinity for oxygen ; but since mercuric oxide breaks up at a relatively low temperature into its constituents, whereas cupric oxide does not undergo decomposition until a temperature above 1000° C. has been reached, and then yields up only a portion of its oxygen, the intensity of affinity for oxygen is much greater in the case of copper. Again, the extent of affinity of carbon towards hydrogen is four times as great as that of chlorine. This may be seen in raethylic hydride (OHJand hydrochloric acid (HCI). But whereas carbon and hydrogen cannot be made to combine directly at all, chlorine and hydrogen unite with evolution of great heat. Here the element of greatest extent of affinity lias least intensity of affinity. One atom of phosphorus can unite with three atoms of chlorine, giving off* much heat, and forming a compound which may be distilled without decomposition; one atom of silver can unite with only one atom of chlorine, and the resulting com- pound is decomposed by the action of daylight. Here extent and intensity of affinity go together.* Modes of Chemical Action.—Matter undergoes chemical change in five different ways, viz.: Ist. By the direct combination of elements or compounds with each other. 2d. By the displacement of one element or group of elements in a body by another element or group of elements. * The above can be regarded only as an approximately correct statement. In nearly every so-called direct combination of elements there is a preliminary decomposition of elementary molecules: H2-j-Cl2= 2HCI. Here the affinity of hydrogen for chlorine is the force which strives to bring about the reaction, and in this it is opposed by the two affinities of hydrogen for hydrogen, and of chlorine for chlorine, which have to be overcome before the reaction can occur. Thus, the apparently lower affinity of carbon for hydrogen may in reality consist in a higher affinity of carbon for carbon—the affinity of hydrogen for hydrogen remaining, of course, the same in both reactions. For the same reason the heat of combination is a complex quantity, and cannot be regarded as an infallible measure of the intensity of affinity (see Thermochemistry). CHEMICAL AFFINITY. 103 3d. By a mutual exchange of elements or groups of elements in two or more bodies. 4th. By the rearrangement of the elements or groups of elements already contained in a body. sth. By the resolution of a compound into its elements, or into two or more Jess complex compounds. 11l ustrations of these five modes of chemical action have already been given (p. 76). Combination.—The part of this subject which refers to the fixed proportions in which the elements combine, has been fully treated of under Laws of Combination (Chap. IV.). But not only are the pro- portions by weight in which every combination takes place perfectly definite, but the amount of heat liberated or absorbed in each com.bi- Bation is also a fixed quantity (see Heat of Chemical Combination, Chap. Decomposition.—The forces which accomplish the resolution of a compound, either into simpler compounds, or into its elements, have been referred to on pp. 36 and 49. The chief of these forces are heat and electricity. The action of heat has frequently been described in the course of this introduction. In the decomposition of compounds by heat two cases may be distin- guished, according as the products of decomposition have, or have not, a tendency to re-combine and form the original compound. Decompo- sition in which this regenerative tendency exists is known as dissocia- tion. The phenomena of dissociation have been very carefully studied, and, in regard to these, definite laws have been deduced ; whereas in the case of the more complex phenomena of ordinary decomposition by heat general principles have yet to be discovered. Decomposition by means of the electric current is termed electrolysis, and the compound which is thus decomposed is termed an electrolyte. The electrolyte must be in the liquid condition—either in solution or in a state of fusion. The current from a voltaic battery, when passed through the electrolyte, decomposes it into two constituents known as ions. The terminals of the battery, which are immersed in the elec- trolyte and on the surfaces of which the separation of the ions occurs, are termed electrodes. The material of the electrodes may vary accord- ing to circumstances, but plates of platinum are generally employed in the case of solutions. Dissociation.—Examples of dissociation have already been given (see Apparent Exceptions to Avogadro’s Law, p. 63). Further examples of dissociable compounds are—the aquates of some salts, which by heating give off* their water of crystallization ; and the carbonates, most of which at a sufficiently elevated temperature evolve carbonic anhydride. A very important law of dissociation is, that the volatile products given °ff by a substance undergoing dissociation have a constant tension for each temperature. This tension corresponds exactly in character to the tension of the vapor of a liquid, and its amount may be measured in the same way (see Chap. XVII.). The tension of dissociation depends en- tirely on the temperature, being higher for higher temperatures; and 18 quite independent both of the space filled by the volatile products 104 INORGANIC CHEMISTRY. and of the quantity of substance which has already undergone decom- position. Thus Debray {Cornpt. Rend., 61, 603) has shown that the tension of dissociation of calcic carbonate is not altered by the addition of an excess of quicklime—the solid product of decomposition. Decomposition may also be effected by means of the electric spark, which may be applied either in the form of the voltaic arc or as the induction spark. In both cases the electric discharge acts solely by its heating effect, and its action must therefore not be confounded with electrolyds. It differs from other sources of heat in being at the same time local and more intense. If a series of induction sparks be passed through carbonic anhydride, those molecules which lie in the path of the spark are broken up by the heat into carbonic oxide and oxygen. The moment the molecules of the two latter gases pass beyond the immediate sphere of the spark, they reach a relatively cold region, the temperature of which lies far below their temperature of combina- tion, so that they can continue to exist in the free state. If, in the above experiment, the proportion of decomposed car- bonic anhydride be allowed to pass beyond a certain limit, re-combina- tion of the oxygen and carbonic oxide will take place with explosion. This occurs as soon as a sufficient number of molecules of the two latter gases are present to propagate the heat of combination through the body of the gas. This propagation is impossible as long as their molecules are separated by a large number of indifferent molecules of carbonic anhydride. Electrolysis.—The following are the laws of electrolysis : 1. The liquid condition is necessary to electrolysis. 2. Electrolytes must be compounds and conductors of the electric current. These compounds generally consist of a conductor and a non-conductor of electricitv. 3. Compounds which suffer electrolysis when dissolved in water do so also when fused. 4. The electrolyte is resolved into two constituents, which, impelled in opposite directions, are eliminated at the opposing surfaces of the two electrodes, and never in the intervening liquid. 5. Oxygen, chlorine, bromine, iodine, and acids appear at the positive electrode, and are, therefore, electro-negative; whilst hydrogen, metals, and alkalies are evolved at the negative electrode, and are, therefore, electro-positive. 6. The quantity of electricity which passes through the electrolyte is always directly proportional to the quantity of the electrolyte which is decomposed. 7. All compound molecules possessing the same active atomicity to be overcome, require, if decomposable, the same quantity of electricity to decompose them. Therefore, if the same electric current be passed through a number of metallic solutions in succession, the metals will be reduced in the ratio of their atomic weights divided by their active atomicities. 8. The quantity of electricity which a compound molecule requires to decompose it, is equal to the quantity which that molecule evolves when it is formed in the generating cell of the battery. CHEMICAL AFFINITY. 105 9. The quantity of electricity evolved by the union of two or more bonds, is capable of effecting the disruption of the same number of bonds in any compound susceptible of electrolysis. The following is a list of weights of various chemical compounds requiring for their decomposition equal quantities of electricity; Water, J(0//H2) 9.0 grams. Hydrochloric acid, .... HC1 30.5 U Argentic chloride, .... AgCl 143.2 (( Cupric chloride, i(Cu"Cl2) 67.1 u Cuprous chloride i('C.i'aCl2) 98.7 u Plumbic chloride, .... i(Pb"Cl2) 138.7 u Antimonions chloride, . . . i(Sb"'Cl8) 75.5 u Plumbic iodide, J(Pb"I2) 230.2 u Plumbic acetate, §(Pb//Ac2) 1622 u Cupric sidphate, J(S04Cu") 79 6 “ Zincic sulphate, J(S04Zn") 80.6 u Stannous chloride, .... 4(Sn"CI2) 94 5 u Ferrous chloride, .... *(Fe"CI2) 63.5 (C Ferric chloride, *(W"8C18) 54.2 u Thus if the electric current were passed through argentic chloride, cupric chloride, and cuprous chloride, included in the same circuit; by the time 143.2 grams of argentic chloride had been decomposed, the quantities of cupric and cuprous chlorides which had undergone decom- position would be 67.1 grams and 98.7 grams respectively. The weight of silver deposited from the first salt would be 107.7 grams; that of copper from the other two 31.6 grams and 63.2 grams, the quantity being in every case in the proportion of the atomic weight of the metal, divided by its active atomicity. What is termed secondary action in electrolysis takes place when the primary products of decomposition exert a chemical action, either on the solvent, or on other substances which are present, or on the elec- trolyte itself. Thus when a solution of sodic chloride is electrolyzed, the salt is broken up into sodium and chlorine. The sodium, however, does not make its appearance as such, but decomposes the water with evolution of hydrogen and formation of sodic hydrate: Na2 + 20 H2 = H2 + 20NaH. Water. Sodic hydrate. Hydrogen and chlorine are thus obtained in the electrolysis of a solu- tion of sodic chloride, but the hydrogen is a secondary product. Again, ff a mixed solution of hydrochloric and hydriodic acids is electrolyzed, no chlorine is evolved, since chlorine instantaneously liberates iodine from the hydriodic acid, regenerating hydrochloric acid. Again, if the positive electrode consists of an oxidizable metal, the electronegative element or group will combine with it. Thus, if acidulated water be electrolyzed with copper as the positive electrode the copper will go into solution, and form a copper salt with the acid. The electrolysis of sulphuric acid and plumbic sulphate has of late 106 INORGANIC CHEMISTRY, acquired great importance in connection with secondary batteries or accumulators as an economic means of storing energy. Various forms of storage battery have been suggested, but all are modifications of the original invention of Plante. They consist essentially of plates c6ra- posed of or coated with plumbic sulphate, these plates being arranged as in primary batteries and immersed in dilute sulphuric acid. When an electric current, either from a primary battery or a dynamo- electric machine, is passed through the cells of a secondary battery, em- ploying the plates of plumbic sulphate as electrodes, the intervening hexabasic sulphuric acid is electrolyzed according to the following equation: SHo6 = S03 + 30 + 3H2. On -j- plate. On plate. The sulphuric anhydride thus liberated is immediately reconverted into hexabasic sulphuric acid : S03 + 30H2 = SHo6. The nascent oxygen in contact with the plumbic sulphate on the positive plate converts the lead salt into plumbic peroxide (Pbo2), liberating sulphuric anhydride, which in contact with water regenerates hexabasic sulphuric acid as just described. The nascent hydrogen on the negative plate converts the plumbic sulphate into lead and hexabasic sulphuric acid : S02Pbo" + H2 + 20H2 = Pb + STIog. Plumbic Hexabasic sulphate. sulphuric acid. Under the influence of an electric current, therefore, the opposing plates of the secondary battery become coated, the one with plumbic peroxide, and the other with metallic lead, the latter being in a spongy state; and they are in a highly electro-polar condition. On joining them by a conductor, a powerful electric current, with an electromotive force of about 2.4 volts for each cell, flows through the conductor from the plate coated with peroxide of lead to that coated with spongy lead, whilst within the cell the current passes through the dilute sulphuric acid in the opposite direction, viz., from lead to peroxide of lead, de- composing the acid as in charging. As, however, the current now flows in a direction opposite to that during charging, the ions are liberated on the opposite plates. On the positive plate, which was formerly the negative electrode, the chemical change is as follows: Pb02 + H3 = PbO + 0H2. Plumbic Plumbic peroxide. oxide. The plumbic oxide, which is thus formed in contact with sulphuric acid, is converted into plumbic sulphate. On the negative plate, which was the positive electrode, the follow- ing action takes place: CHEMICAL AFFINITY. 107 Pb + O + SHo6 = S02Pbo" + 30H2. Hexabasic Plumbic sulphuric acid. sulphate. During the discharge, therefore, both plates return to their original condition. Instead of discharging the plates immediately, however, the energy invested in them may, with but inconsiderable loss, be allowed to re- main stored for weeks, or even months, ready at any moment to yield a powerful electric current available for the production of light, heat, or mechanical power. Electrochemical Equivalents.—For some time after the revival of the atomic theory in its chemical form by Dalton, chemists were at a loss which of several possible atomic weights of an element to accept as the true one. The laws of vapor-density, of specific heat, of isomorphism, were enunciated not very long after; but as their significance was not generally perceived, their application as a means of checking the atomic weights was out of the question. In the midst of the uncertainty which prevailed, the law of electrolysis as stated by Faraday was eagerly wel- comed. According to this law the quantities of various electrolytes decomposable by the same current are chemically equivalent, and the quantities of the several elements eliminated in such decompositions are also chemically equivalent. On this principle chemists constructed tables of equivalents of the elements, representing the relative weights which are eliminated in electrolysis, that of hydrogen being taken as unity. Such equivalents would be, for example: H = 1 o=B Cl = 35.5 8 = 16 Pb = 103.25 etc. This mode of procedure was thus far strictly legitimate, inasmuch as the above weights can replace each other in chemical combination, and are therefore equivalent. But most chemists went further than this, and assumed that these equivalents were identical with the atomic weights of the elements. By this means the significance of the above- mentioned three important laws was effectually obscured, and a true chemical classification was for many years rendered impossible. Furthermore, the system of equivalents was not logically carried out. The electrolytic equivalent of antimony is 40; but instead of this the number 120, its present atomic weight, was adopted. The same hap- pened with several other elements. Another objection to this system is that the equivalent of an element does not, like its atomic weight, represent a constant quantity, but varies with the active atomicity. This may be seen in the case of cop- per in its cuprous and cupric salts. A knowledge of the so-called equivalent notation is necessary for the study of many important works on chemistry in which it is employed. 108 INORGANIC CHEMISTRY. The old equivalent formulae may be converted into modern atomic formulae, either by doubling the number of the perissad, or by halving that of the artiad atoms (see p. 79). Thus: Old so-called equiva- Atomic lent formula. formula. Water, . . . . ... HO h2o Sulphuric acid, . . . . . nso4 h2so4 Nitric acid, . . . . bno6 HN(>3 Ferric chloride, . . . . . Fe2Cl, Fe2Cl6 In modern works equivalent formulae, when quoted, are generally written as above, in italics. The fact that a single atom of one element may be equivalent to two or more atoms of another, sufficiently explains the discrepancies between atomic and equivalent proportions noticed in treating of the law of equivalent proportions (see p. 47). CHAPTER XIII. CHEMICAL HOMOGENEITY. A CHEMICALLY homogeneous substance is one in which all the mole- cules are exactly alike. It is evident from this definition that such a substance will exhibit constant composition : if it is a simple body, it will yield on analysis no other body; if it is a compound, it will con- tain the same ingredients in unvarying proportion. But in the case of compounds, analysis alone cannot furnish complete evidence of the homogeneous nature of a substance: for example, it is plain that a mix- ture of molecular proportions of manganous oxide (MnO) and manganic peroxide (Mno2), would yield analytical results corresponding to man- ganic oxide (Mn2o3). Hence, other means of identification are necessary, and these are frequently to be found in the physical properties of the substance. Thus, all substances, in whatever physical state they exist—gaseous, liquid, or solid—possess a definite specific gravity at a given temper- ature. The specific gravities of the more important chemically homo- geneous substances have been determined, and it is thus frequently possible to identify a substance, as it is not probable that a mixture accidentally possessing a percentage composition the same as that of a true chemical compound, would also have the same specific gravity. This characteristic is least certain in the case of solids, where a slight alteration in physical condition, such as that produced in metals by hammering, is sufficient to cause a change in the specific gravity. Such variations, however, occur within narrow limits. The number of the characteristics available for establishing the chem- ical homogeneity of substances varies with the complexity of the phys- CHEMICAL HOMOGENEITY. 109 ical state, being greatest in the case of solids, and smallest in the case of liquids. Gases.—A mixture of equal volumes of methylic hydride (CH.t) and propylic hydride (C3HB) would yield not only the same analytical re- sults as ethylic hydride (C2H6), but would also possess the same specific gravity. In this case the best method of determining whether the gas is single or a mixture, is to submit it to diffusion. For this purpose, it is transferred to a tube over mercury, closed at the upper extremity by a porous diaphragm (graphite or gypsum). By the law of diffusion {q.v.) the lighter molecules will pass through the diaphragm more rap- idly than the heavier molecules. If, therefore, in the above case, on examining the residual gas, the proportions of carbon and hydrogen be found to have changed, it may be concluded that the original gas was a mixture; if these proportions remain the same, then either the gas is single or it is a mixture in which each gas is present in the ratio of its coefficient of diffusion, a case which must necessarily be of very rare occurrence. Sometimes the gas is submitted to the action of various absorbents—caustic potash, potassic pyrogallate, fuming sulphuric acid. If part be absorbed by any of these reagents, whilst part remains unacted upon, it is at once proved that the gas is a mixture. Liquids.—When a liquid can be distilled without decomposition, its boiling-point affords one of the best tests of its homogeneity. Every chemical compound which is capable of volatilizing without decompo- sition, has, at a given barometric pressure, a fixed boiling-point, at which it must distil from the first to the last drop. As a rule, a mix- ture of two liquids of different boiling-points will begin to boil about the boiling-point of the lowest, and a thermometer placed in the vapor will in turn indicate all temperatures up to the boiling-point of the highest. Mixed liquids may be separated by fractional distillation; the fractions of the distillate passing over at different temperatures are collected separately, and these fractions are redistilled until liquids of constant boiling-point are obtained. Some liquids cause the plane of a ray of polarized light which passes through them to rotate to the right or to the left, and, as this rotation is constant for a given stratum of a given liquid, the action on polarized light may be frequently employed in the case of such liquids as a test of their purity. Solids.— When a solid possesses the property of crystallizing, its crys- talline form offers the surest means of identifying it. If the crystals are so well developed that their angles may be measured, the values of these angles, coupled with the analytical results, suffice to place the identity of any substance for which such determinations have previously been made, beyond all possibility of doubt. Even when the crystals are too small to admit of measurement, a microscopic examination will generally be sufficient to decide whether they are homogeneous or mixed. Heterogeneity of crystalline form does not necessarily involve chemical difference; a substance may be dimorphous. Thus the sublimate of arsenious anhydride frequently contains, side by side, rhombic prisms and regular octahedra. When solids are fusible, they possess a con- stant fusing-point. This property is of great value in identifying organic substances, of which the greater number fuse within the limits 110 INORGANIC CHEMISTRY. of the mercurial thermometer. As mixtures fuse at a lower tempera- ture than the mean fusing-point of their constituents, impurities gener- ally tend to lower the fusing-point. Every soluble solid, when pure, has a fixed solubility for each of its solvents at a given temperature. This solubility generally increases with the temperature (see Solubility). If the various ingredients of a mixture possess very different solubilities, this property may be taken advantage of in order to effect their separa- tion, as the least soluble will crystallize out first, and, by repeated recrystallization, may generally be obtained pure. What is known as fractional crystallization consists in evaporating the solution of a sub- stance until sufficiently concentrated to crystallize. The liquid is then separated from the crystals and evaporated until a fresh crop of crystals is obtained. This process is repeated until the solution is exhausted. If the last crop of crystals is exactly like the first, as regards composi- tion, form, fusing-point, etc., it may be concluded that the substance was homogeneous. The reverse of fractional crystallization is fractional solution. The solid substance is successively extracted with small por- tions of the solvent. In this way the more soluble ingredients, if such are present, will be removed. Sometimes various solvents are employed in succession, according to the nature of the substances suspected to be present in a mixture; and in this way a separation may frequently be effected. Fractional precipitation consists in adding to a solution a precipitant in quantity insufficient to precipitate the whole of the sub- stance present. In a mixture, the various ingredients will probably be affected in varying degrees by the precipitant—that, for example, which has the greatest affinity for the precipitant will be found chiefly in the first fraction. By redissolving this fraction and partially precipitating it, and repeating this operation each time with the partial precipitates, one of the ingredients of the mixture may usually be obtained pure. This process is seldom necessary in the case of inorganic compounds, as with these a sharp separation by means of precipitants is generally at once possible. Fractional saturation is analogous to fractional precipitation, and depends on the varying degrees of affinity which the ingredients of a mixture exhibit towards the saturant. A mixture of bases, for example, is imperfectly saturated with an acid; a mixture of acids, with a base. These fractional methods are chiefly of use in the case of organic compounds, which very seldom possess properties such as to render them separable from each other by a single operation. In the case of single substances such methods afford a guarantee of purity by the correspondence of the different fractions; and, in the case of mixtures, they yield, by systematic repetition, a means of separating the various ingredients. CHAPTER XIV. ISOMERISM, METAMERISM, POLYMERISM, ALLOTROPY. Compounds which, while possessing the same percentage composi- tion, exhibit differences of chemical and physical character, are termed isomeric. Metamerism and polymerism are special cases of isomerism ; HEAT OF CHEMICAL COMBINATION. THERMOCHEMISTRY. metameric compounds have the same molecular weight, the difference m properties depending on difference of arrangement of the atoms within the molecule; in polymeric compounds the molecular weights are different, one being a multiple of the other. Examples of meta- merism and polymerisra are most common among the compounds of carbon, where the frequency of high molecular weights and the prop- erty which carbon possesses of repeatedly combining with itself, favor variety of atomic arrangement. The compounds propionaldehyde, acetone, ally lie alcohol, propylenic oxide, and trimethylenic oxide, all possess the molecular formula C3II60, and are, therefore, metameric. The hydrocarbons of the ethylenic or CuH2n series, ethylene (C2H4), propy- lene (C3H6), butylene (C4HB), etc., are polymeric. The single members of this group may possess metamers; thus, there are three butylenes of the formula C4lIB—butylene, isobutylene, and pseudobutylene. Allotropy stands in the same relation to elements that isomerism does to compounds. Many of the elements exist in several different modi- fications, possessing entirely distinct properties. Carbon is known in three forms: as charcoal, as graphite, and as diamond. Sulphur and phosphorus also possess allotropic modifications. One of the most striking and instructive instances of this phenomenon is found in the case of oxygen in its two modifications of common oxygen and ozone. It is probable that allotropy is to be explained by reference rather to polymerisra than to metamerism. It is certainly conceivable that molecules containing equal numbers of only one kind of atom should differ through the arrangement of these atoms within the molecule; but a difference of properties can more easily be accounted for by supposing that the molecules of the allotropic modification contain different numbers of atoms, and in the only case of true allotropy in which the molecular weights of the allotropic modifications are known, this is found to be the case. Common oxygen contains two atoms in the molecule, whereas ozone contains three. It is to be noted that allotropy has been observed only in the case of polyad elements. The atoms of a monad element can only combine with each other in pairs, thus H—H, and in this way all variety, either in the number of atoms in the molecule, or in their arrangement, is excluded. CHAPTER XY. HEAT OP CHEMICAL COMBINATION. THERMOCHEMISTRY. Thermochemistry, that branch of the science which deals with the heat liberated or absorbed in chemical action, has been studied in great detail by Berthelot, Thomsen, and others. The first-named chemist has published {Ann. Chim. Phys. [4], VI., and [s], IY.) a summary of the results obtained in this field, and from this source the annexed account is extracted. He enunciates as the three fundamental laws of thermochemistry the following: 1. Law of Molecular Work.—The quantity of heat liberated in any reaction is a measure of the sum of the chemical and physical work performed in that reaction. 112 INORGANIC CHEMISTRY. 2. Law of the Equivalence of Heat and Chemical Change.—When a system of bodies, simple or compound, taken in definite conditions, undergoes physical or chemical changes which are capable of bringing the system into a new state without producing any mechanical effect external to the system, the quantity of heat liberated or absorbed during these changes depends solely on the initial and final states of the system, and remains the same, whatever be the nature and order of the intermediate states. 8, Law of Maximum Work.—Every chemical change, accomplished without the intervention of foreign energy, tends to the production of that body, or system of bodies, in the formation of which most heat is liberated.* The first two laws are corollaries of the law of the conservation of energy; the third must be developed more in detail. It is possible to conceive the necessity of tin’s law by considering that the system which has given off most heat no longer possesses the energy necessary to accomplish a fresh transformation. Every fresh change involves the performance of work, and this work cannot be performed without the intervention of foreign energy. On the other hand, a system capable of liberating heat by a fresh change, still possesses the energy requisite to produce this change without foreign aid. It is in the same way that a system of heavy bodies tends to that arrangement of its parts in which the centre of gravity is as low as possible; but the system will only attain to this arrangement should no foreign obstacle intervene. This is, however, rather an illustration than a demonstration. lit the equations which will now be employed in proof of this law, the atomic and molecular weights are to be understood in grams. The units of heat will then be calories (see p. 68). The latter are written to the right of the equation, and denote the heat liberated by the com- bination represented in the equation, supposing the combining quantities to be taken, as stated above, in the proportion of grams. Combination.—According to the law of maximum work, oxygen, in combining with other bodies, will form a higher oxide or a lower oxide, according as the one or the other stage corresponds to the greater lib- eration of heat. In the formation of nitrons anhydride from two molecules of nitric oxide and one atom of oxygen, the thermal effect is as follows; 2'N"0 + O = N203, .... 20,000 cal. Nitric Nitrous oxide. anhydride. But when two molecules of nitric oxide combine with two atoms of oxygen to form nitric peroxide, the calorimetric equation is: 2'N"0 + 02 = 'N204, .... 34,000 cal.; Nitric Nitric oxide. peroxide. * It ought to be mentioned that the universal validity of the law of maximum work has been called in question. Some of the objections urged against the law have been successfully met by its author; but there are anomalies connected with the phenomena of heat of neutralization which do not appear capable of explanation on Berthelot’s theory. (See more fully p. 115.) HEAT OF CHEMICAL COMBINATION. THERMOCHEMISTRY. 113 or, the quantity of heat liberated is greater by 14,000 calories. There- fore, whenever an excess of oxygen is present, nitric peroxide ought to be formed. Not only is this the case, but nitrous anhydride combines directly with oxygen to form nitric peroxide : N203 + O = 'N204, .... 14,000 cal., Nitrous Nitric anhydride. peroxide. On the other hand, hydrogen in combining with oxygen to form water yields : H2 + O = 0H2, 69,000 cal., Water. whereas, when these two elements unite to form hydroxyl, the effect is: H2 + 02 = '0'2Ha, 47,000 cal., Hydroxyl. giving a difference of 22,000 calories in favor of the lower oxide. When oxygen and hydrogen combine, water ought, therefore, to be formed, whilst hydroxyl ought to have a tendency to decompose into water and oxygen. Furthermore, the formation of hydroxyl, starting from water and oxygen, ought to be accompanied by an absorption of heat. This com- pound cannot, therefore, be formed without the intervention of some foreign energy—for instance, that of a simultaneous chemical action. There are several compounds, in the formation of which, starting from their elements, heat is absorbed. Such, for example, are the ox- ides of nitrogen, the oxides of chlorine, chloride of nitrogen, acetylene, cyanogen, etc.; and none of these can be produced by the mere inter- action of their elements, acting by their intrinsic energy. Acetylene, for example, is formed by the direct union of carbon and hydrogen ; but this combination does not take place under the influence of chemical affinity alone : it requires the aid of the electric arc. The oxides of nitrogen are all derived from nitric peroxide, which can be formed from its elements only under the influence of intense heat (elec- tric discharge, simultaneous combustion of hydrogen). The oxides of chlorine are produced by the action of chlorine on the alkaline oxides; but this is because their formation is accompanied by that of an alkaline chloride, the production of which is attended with liberation of much heat. Decomposition.—A body that lias been formed directly from its ele- ments with liberation of heat will not spontaneously decompose ; the intervention of external energy is necessary to separate its elements. Such forms of external energy are heat, light, electricity, a simultane- ous chemical action and the energy of disaggregation developed by so- lution. The action of this last agent is displayed in the case of salts of weak acids, and those of certain feebly basic metallic oxides. 114 INORGANIC CHEMISTRY. If, however, a compound be formed with absorption of heat, it will be capable of effecting its own decomposition. This is the case with the oxides of chlorine, which explode under the slightest disturbing in- fluence; to this class belong chloride of nitrogen, ammonic nitrite, etc., bodies which decompose spontaneously at ordinary temperatures. When bodies formed with absorption of heat do not readily undergo sponta- neous decomposition, they show a marked tendency to enter into direct combination or to undergo fresh chemical changes—such as polymeric condensation, breaking up into groups, complex decomposition—all of which changes are accompanied by liberation of heat. Bodies formed with absorption of heat are, moreover, particularly sensitive to the ac- tion of so-called catalytic or contact agents. Such agents do not in these cases usually introduce any special energy into a reaction ; they merely serve to liberate a store of pre-existent potential energy. Substitution.—Substitutions also take place according to the law of maximum work. Chlorine, in combining with hydrogen or the metals, liberates more heat than bromine, and bromine liberates more than iodine. Therefore bromine decomposes the iodides, expelling iodine, and forming bromides; chlorine decomposes both bromides and iodides, expelling bromine and iodine, and forming chlorides. In the same manner, whenever one metal displaces another from its salts, the forma- tion of the new salt is attended with a greater liberation of heat. From this follows the well-known direct relation between the electromotive force of the metals and their heat of oxidation. Double Decomposition.—ln general one hydrated ba.se displaces another from its salts, when it liberates more heat in combining with the same acids.* This is the case when the hydrates of the metals are precipitated by alkaline solutions. Thus : fN02 { Pho" + 2KHo = 2N02Ko + PbHo2 Ino2 Plumbic Potassic Potassic nitrate. Plumbic hydrate. nitrate. hydrate. This reaction liberates 12,200 cal. if all the compounds are in solu- tion, and 45,G00 cal. if they are in the solid state. In the same way, one acid expels another from its salts, when it liberates more heat in combining with the same base; at least, this is so in all cases whereeach of the acids forms only one salt with the base. But all these relations are only then strictly true, when the heat liberated by the acids, bases, and salts is calculated for these bodies in the same physical condition, namely, the solid state. The following example will show how a change of physical condition and the special combinations formed with the solvent may affect the result. Gaseous hydrochloric acid acts upon dry mercuric cyanide, forming mercuric chloride and hydrocyanic acid : 2HC1 + HgCy2 = 2HCy + HgCl2 . . + 10,600 cal. Hydrochloric Mercuric Hydrocyanic Mercuric acid. cyanide. acid. chloride. * See, however, p. 115. HEAT OF CHEMICAL COMBINATION. THERMOCHEMISTRY. 115 But hydrocyanic acid in solution acts upon mercuric chloride in solu- tion, forming mercuric cyanide and hydrochloric acid. This reversal of the reaction is explained by the fact that two molecules of hydro- cyanic acid in solution liberate in acting upon mercuric oxide 31,000 cal., whilst a solution of hydrochloric acid liberates only 19,000 cal. There are therefore -j-12,000 cal. liberated in the reaction in the wet way, a result which experiment completely confirms. Theory, there- fore, predicts this reversal of the reaction corresponding to the change in the thermal sign. This change is due to the intervention of a new chemical reaction attended by liberation of heat, the combination of gaseous hydrochloric acid with water, by which the hydrochloric acid has yielded up a portion of its energy. The same principle of maximum work enables us to produce a num- ber of compounds which could not be obtained directly, because their formation is attended with absorption, and their decomposition with liberation of heat. This end is attained by the device of a double de- composition bringing about the simultaneous formation of some other compound, the production of which is attended with a liberation of heat greater than the absorption first mentioned. For example, in the for- mation of hydroxyl from oxygen and water, 0H2 + O = 'O'.H, —21,800 cal. there is absorption of heat, and the reaction cannot therefore take place directly. In order to accomplish it, baric oxide is made to combine with oxygen, thereby liberating 11,800 cal.; and the baric dioxide thus obtained is acted on with dilute hydrochloric acid, forming baric chlo- ride and hydroxyl, with liberation of 22,000 cal. more. The formation of baric chloride furnishes the supplementary energy which is employed in producing hydroxyl. The rules given by Berthelot for the relation between the heat of neutralization of acids and bases, on the one hand, and their mutual affinity on the other, do not hold good in the case of solutions. In fact, the very reverse is frequently the case. Thomsen has made a series of careful determinations of the heat of neutralization of various acids and bases, and he shows that in mixed solutions of equal equivalents of two acids with a quantity of a base only sufficient for the neutralization of one, the larger portion of the base is frequently appropriated by that acid with which it evolves least heat in neutralization. This is in direct opposition to Berthelot’s law of maximum work. Ostwald, by measuring the contraction or expansion which occurs on mixing solutions of acids and bases, has arrived at the same conclusion. It appears, therefore, that the heat of neutralization cannot be regarded as a measure of affinity. Thomsen shows that every base and every acid has a fixed heat-equivalent, which is liberated in its neutralization, and that the heat of neutralization in any given case is the sum of the heat- equivalents of acid and base. This follows from the fact that, if any two acids be neutralized with a given base, the difference between their heats of neutralization will be the same for their neutralization with 116 INORGANIC CHEMISTRY. any other base, provided always that acids, bases, and salts in every case remain in solution. The same holds for the neutralization of bases with various acids: the difference between the heats of neutral- ization of any two bases with a given acid is the same for their neutralization with any other acid. It follows from this that the heat of neutralization is independent of the degree of affinity between acid and base. Ostwald has shown that a precisely similar law regulates the contraction or expansion which occurs when solutions containing equivalent quantities of acid and base are mixed : the difference in the degree of change of volume for any two acids with any given base is the same with any other base; each acid and each base produces its own definite and invariable change of volume, and the change of volume in any given case of neutralization is the sum of the changes for acid and base. The heat of neutralization appears to be greater the greater the contraction, or the smaller the expansion. Taking these facts together, the conclusion seems unavoidable that the heat of neutralization is directly connected, not with chemical affinity, but with the changes which occur in the aggregation of the solution—expansion and contraction. The great obstacle to the interpretation of thermocheraical data lies in the fact that, under the conditions of temperature at which calorimetric determinations are possible, there is no such thing as mere direct com- bination of elements. The thermal equation, H + Cl = HCI 22,000 cal. is a fiction. This equation ought to be written H2 + Cl2 = 2HCI 44,000 cal. and the thermal effect 44,000 cal. is in reality the algebraic sum of three distinct thermal effects—the heat absorbed by the separation of hydrogen from hydrogen, the heat absorbed by the separation of chlo- rine from chlorine, and the heat liberated by the union of hydrogen with chlorine. If the first of these be denoted by x, the second by y, and the third by z, we should have— 2z (x f y) = 44,000 cal. Every thermal equation (except such as contain elements with mona- tomic molecules) is therefore a single equation with three unknown quantities, which are consequently undeterminable. If hydrochloric acid could exist at a temperature at which the mole- cules of hydrogen and chlorine dissociate into single atoms, then the conditions of the first of the above thermal equations would be realized and x and y would be eliminated. But if there are such conditions, they lie far above the range of temperature at which such determina- tions are at present possible. FUSION AND FUSING-POINTS. 117 CHAPTER XYI. FUSION AND FUSING-POINTS, The molecular changes which correspond to the passage of a body from the solid to the liquid state have already been discussed. As these changes depend on the energy of the molecules, and as this energy will be constant for any given body at a given temperature, it is evident that every substance which is fusible at all ought to have a fixed fusing- point, and such is, with few exceptions, the case. The use of the fusing-point as a means of identifying substances and testing their pu- rity has also been described. Change of Volume Accompanying Fusion. Most substances in passing from the solid to the liquid state expand : the melted substance is the specifically lighter. With water and bismuth the reverse is the case; these bodies expand in solidifying. Thus, ice floats on the sur- face of water; and closed vessels, in which water is frozen, burst with the internal pressure. Effect of Pressure in Altering the Fusing-point.—lf a body expands in fusing, increase of pressure will tend to raise the fusing-point. In this case, the pressure acts counter to the energy of the molecules. The efl'ect is very slight: according to Bunsen, a pressure of 156 atmos- pheres is necessary to raise the fusing-point of spermaceti from 47.7° C. to 50.9° C. If, on the contrary, fusion is accompanied by contraction, an increase of pressure will lower the fusing-point, the pressure in this case aiding the energy of the molecules. The effect in the case of water is a lowering of the fusing-point by .0075° C. for each atmos- phere. Moussou succeeded, by means of very great pressure, in melt- ing ice at —lB° C. Latent Heat of Fusion.—lf a given weight of water at 100° C. be mixed with an equal weight of water at 0° C., the temperature of the mixture will be 50° C. If a given weight of water at 100° C. be mixed with an equal weight of powdered ice at 0° C., the temperature of the mixture will be only 10.4° C. If we suppose that, in this last case, a gram of each wras taken (though in practice the experiment could not be accurately performed with such small quantities), the gram of water at 100° C. in cooling to 10.4° C. will have given off* 100 10.4 = 89.6 calories. But in giving off this quantity of heat, it has melted one gram of ice and raised the temperature of the resulting gram of water 10.4° C. This last rise of temperature will represent 10.4 ca- lories. Therefore, as the heat given off is equal to the heat taken up: Melting of 1 gram of ice -f- 10.4 cal. = 89.6 cal.; or Melting of 1 gram of ice = 79.2 cal. In other words, when one gram of ice at 0° C. is converted into one gram of water of the same temperature, 79.2 calories—a quantity of heat sufficient to raise the temperature of an equal weight of water 118 INORGANIC CHEMISTRY. 79.2° C.—disappears. This quantity of heat is known as the latent heat of fusion of ice, or, as it is sometimes termed, the latent heoi of water. The energy of motion represented by this latent heat is taken up by the molecules in some form which does not affect the thermometer: it occa- sions no rise of temperature, but only brings about a difference in the condition of the molecules in regard to each other, each molecule being enabled to overcome the attraction of its immediate neighbors, and to wander through the liquid. All substances capable of assuming the liquid state possess latent heat of fusion. Water has the highest latent heat of all known liquids. The disappearance of heat in the liquefaction of ice may be roughly shown by heating over a flame a vessel containing pieces of ice. As long as any ice remains unmelted, the temperature will rise very little above 0° C., all the heat which is taken up by the water being instantly employed in melting the ice. By first pounding the ice so as to increase the surface, and stirring continually so as thoroughly to mix the ice and water, the temperature of the whole may be kept af 0° C. As soon as the ice is melted, the temperature of the water will begin to rise as usual until the boiling-point is reached, when the temperature will again re- main constant. The heat which disappears when a body passes from the solid into the liquid state, is again evolved in the passage from the liquid to the solid state. (See suspended solidification.) The cold which is produced by the solution of solids is attributable to the same cause. (See solubility.) Iti the process of solution, a solid in contact with its solvent may become liquid without the application of heat. Hence, when the latent heat of liquefaction of the solid dis- appears, the temperature of the whole is lowered, the heat of liquefac- tion being taken from the mass itself. This is the principle involved in freezing-mixtures. In such mixtures, the more rapid the process of solution or liquefaction without application of external heat, the greater is, cceteris paribus, the degree of cold attainable, there being less time for heat to be taken up from without. A mixture of 5 parts of ammonic chloride, 5 of potassic nitrate, and 19 of water, produces a reduction of temperature from -f 10° to 12° C. A solution of common salt in water freezes at a much lower temperature than pure water; if, there- fore, salt be added to snow, the latter will melt. In this case there is simultaneous liquefaction of the snow and solution of the salt; but owing to the great latent heat of water, the cold is derived chiefly from the former source. A mixture of three parts of snow with one of com- mon salt produces a cold of—22° C. If equal weights of snow and dilute sulphuric acid, previously cooled to —7° C., be mixed, the tempera- ture will sink as low as 51° C. The researches of Guthrie into the nature of the solid compounds which various salts form with water, have thrown great light upon the mode of action of freezing-mixtures and upon the degree of cold attain- able by their means. Guthrie shows that all salts which are capable of dissolving in water form definite solid compounds with this solvent, and that every such compound has a fixed fusing-point. To the com- pounds of this class which are solid only at temperatures below 0° C., EBULLITION AND BOILING-POINTS. he has given the name cryohydrates. The same salt frequently forms more than one cryohydrate. Thus sodic chloride, which at —lo° C. crystallizes with 20H2, combines at a still lower temperature with 10.5 OH2, yielding a compound fusing at 22° C. The important law holds good that the fusing-point of that cryohydrate which is formed at the lowest temperature is the limit to the degree of cold attainable with a given freezing mixture, since any further abstraction of heat from the mixture occasions, not depression of temperature, but separation of the cryohydrate. Thus the greatest degree of cold which can be pro- duced with a mixture of ice and sodic chloride is 22° C. Further, the maximum effect from a freezing mixture is obtained when the in- gredients are employed in the proportions requisite for the formation of the cryohydrate. Suspended Solidification.—Although it is not possible (at least at ordinary pressures) to heat a substance a single degree above its fusing- point without producing liquefaction, yet many substances, when fused, may be cooled many degrees below their fusing-point without solidify- ing. This state, which is known as suspended solidification, is most readily produced in bodies from which air is excluded. Water inclosed in a small glass vessel from which the air has been removed may be cooled as low as —B° or —lo° C. without solidifying. The fusing- point of phosphorus is 54° C.; but if melted under water, it may be cooled to 32° C. without becoming solid. If a liquid body, thus cooled below its fusing-point, be touched with a portion of the same body in the solid state, solidification instantly ensues, and the temperature of the mass rises to the fusing-point. The cause of this rise in temperature is the latent heat of fusion, which is again evolved when the body passes back into the solid state. Solidi- fication may also frequently be induced in such cases by agitation. CHAPTER XVII. EBULLITION AND BOILING-POINTS. When the molecules of a liquid, in the course of their wanderings, reach the free surface of the liquid, they are carried by the force of their motion, should this happen to be in an upward direction, into the air. Here they behave like the molecules of a gas, striking against other molecules—either of the air or of their own kind—sometimes proceed- ing further upwards, sometimes being thrown back into the liquid. If the space above the liquid is unlimited, the molecules above the liquid will gradually wander away from it and no longer be exposed to the risk of falling into it again, whilst their place will be constantly taken by fresh molecules from the surface. This is the phenomenon of spon- taneous evaporation at ordinary temperatures. If the space above the liquid is limited, the diffusion of molecules into it from the liquid will go on as before; but a point will be reached at which the number of 120 INORGANIC CHEMISTRY. molecules which fall back into the liquid is as great as that of the molecules which leave its surface, upon which the evaporation will appear to cease, though in reality it is going on as before. The space is then said to be saturated with vapor. The quantity of vapor which will thus diffuse into a given space is constant for a given temperature and independent of the pressure. Thus at a given temperature the same quantity of vapor will diffuse into a vacuum and into an equal space containing air, the only difference being that the vacuum will fill more rapidly with vapor, as there are no molecules of air to oppose the passage of the molecules of vapor. This vapor exerts a pressure, and as this pressure must be proportional to the quantity of vapor present in the unit of space, it will also be constant for any given tem- perature. This pressure is known as the tension of the vapor of the liquid. Its action may be illustrated, and its amount measured, as fol- lows: Two barometer-tubes are filled with mercury and inverted over a mercury trough. The mercury will stand equally high in both, and the height of the column will represent the pressure of the atmosphere. A few drops of water are now introduced into one of the tubes by al- lowing the water to rise through the mercury in the tube. In a very short time this column of mercury will show a marked depression, corresponding to the tension of the vapor of water for that temperature. If this barometer-tube be surrounded with a second wider tube, which can be filled with water of various temperatures, it will be noticed that as the temperature rises, the mercury in the barometer-tube sinks, cor- responding to the increased vapor tension. The difference in height between the columns of mercury in the two barometer-tubes at any given temperature, will give the vapor tension of water for that tem- perature, When the temperature reaches 100° C., the boiling-point of water, the mercury inside and outside the tube with the water will stand at the same level—in other words, the tension of the vapor inside the tube exactly balances the pressure of the atmosphere. Hence the im- portant law: The temperature at which a liquid boils is that at which the tension of its vapor is equal to the atmospheric pressure. The mo- ment this point of equality is passed, the molecules from the surface of the liquid stream forth freely into space, carrying before them the layer of air which presses upon them. Bubbles of vapor are formed in the interior of the liquid, rise through it, and are discharged at its sur- face. From the above law it follows, that by lowering the pressure, the boiling-point of a liquid may also be lowered. Water will boil in a vacuum at ordinary temperatures, if means be taken to absorb the aqueous vapor as quickly as it is formed. In like manner, by raising the pressure, the boiling-point may be raised. By heating water in a strong closed vessel, by which means the liquid is subjected to the pres- sure of its own vapor, the temperature may be raised far above 100° C. without causing ebullition. There is, however, for every liquid a fixed temperature beyond which no degree of pressure will suffice to restrain the liquid from passing into the gaseous state. This temperature is known as the critical point. If the liquid be heated in a very strong glass tube, the surface of the liquid, when the critical point is reached, EBULLITION AND BOILINQ-POINT3. 121 will be seen to disappear, and the whole tube will be filled with trans- parent vapor, almost of the same density as the liquid itself.* The law that the tension of a vapor is constant for a given tempera- ture aud independent of the pressure, holds only for what are known as saturated vapors—vapors in contact with an excess of their liquids. When the space is not saturated with the vapor, and there is none of the liquid present from which a greater supply may be derived, the vapor behaves, in regard to temperature and pressure, like a true gas: for example, a forcible diminution of the volume would cause a corre- sponding increase in the pressure. In the case of a saturated vapor such a diminution of volume would only occasion a partial condensation of the vapor, the pressure remaining as before. Non-saturated vapors are also termed superheated. When a liquid assumes the gaseous form, its molecules have to over- come, besides the pressure resting on the liquid, the force of cohesion, that is, of their mutual attraction. Hence anything which tends to in- crease the force of cohesion will raise the boiling-point of the liquid. As the attraction between the molecules of a substance and those of the liquid in which it is dissolved is greater than that of the molecules of the liquid for each other, it is clear that the presence of any solid substance in solution will increase the force of cohesion, and conse- quently raise the boiling-point of the liquid. Hence it is that aqueous solutions of salts boil above 100° C. The boiling-point of such solu- tions rises with the concentration. The boiling-point of a liquid is best ascertained by means of a ther- mometer immersed in the vapor of the liquid. The temperature at which the liquid enters into ebullition varies with the nature of the vessel in which it is contained; but the temperature of its vapor or steam is constant. Water boils in a glass vessel at a higher temperature than in a vessel of iron, owing to the greater adhesion between water and glass, which hinders the formation of bubbles of steam at the points of contact of the liquid and the vessel. By heating in a glass vessel water from which the air had been previously expelled by boiling, the tem- perature may be raised several degrees above 100° C. without ebulli- tion supervening. When this state of molecular inertia is from any cause disturbed, ebullition suddenly commences with explosive violence, and the temperature sinks to 100° C. Liquids thus heated above their boiling-points are said to be superheated, and the phenomenon of sudden percussive ebullition is commonly known as humping. Various attempts have been made to discover some law connect- ing the boiling-point of a liquid with its constitution or molecular weight. Such laws as have been deduced hold only for compounds belonging to the same group, and generally only for a few members of * According to Ramsay, however, the critical point is merely the temperature at which the liquid in the tube has the same specific gravity as its vapor, and a gas may be liquefied at any temperature, provided sufficient pressure be applied. 122 INORGANIC CHEMISTRY. such a group. Moreover, the correspondence between experiment and theory is seldom more than approximate. A very few examples will suf- fice. The normal alcohols of the general formula CnH2n+iHo dis- play among their lower members a difference of boiling-point amount- ing to about 19.5° C. for every difference of CH2 in the molecular formula. For a similar difference of CH2 in the normal fatty acids of the general formula CnH2n+i (COHo), the difference of boiling-point is about 22° C. The difference becomes, in the case of the acids, rap- idly less for the higher members. Normal alcohols. Boiling-point. Difference. Ethyl ic alcohol, 0.2H5Ho 78° 19 4 Propylic “ C3H7Ho 97.4 1 l/.T 19.5 Butylic “ 04H9Ho 116.9 • 20.1 Araylic “ 05HuHo 137 20.5 Hexylic “ C6E13Ho 157.5 i q ft Heptylic “ CrH15Ho 176 1 i/.O Normal fatty acids. Boiling-point. Difference. Acetic acid { COHo 118° ,, 7 ( 0 IT Propionic acid < CO Flo 140.7 Butyric acid -f 163 J 1 COHo 216 Valeric acid 1 new 184.5 C H 20.5 Caproic acid < COHo 205 (Enanthylic acid COHo 223.5 Caprylic acid | COHo 236.5 Pelargonic acid / qqj^0 253.5 Latent Heat of Vapors.—lt lias already been mentioned that bodies, in passing from the solid to the liquid state, take up heat without ex- hibiting any rise of temperature, the heat which thus disappears being employed in producing a change in the molecular condition. The same phenomenon is observed in a still more marked degree during the pas- sage from the liquid to the gaseous state. If two thermometers be introduced into a flask of water boiling over a flame, one being plunged in the liquid, the other suspended in the steam, both will register the same temperature, 100° C. (Thethemometer in the liquid may happen to be a fraction of a degree higher; see Boiling-points.) This temper- ature will be preserved by both thermometers, as long as there is any liquid left, though all the time heat is being communicated to the water. The heat which thus disappears in causing a change of mole- cular condition, is known as the latent heat of steam, and is evolved EBULLITION AND BOILING-POINTS. 123 again in exactly the same quantity when the steam is condensed. This last fact is turned to account in the determination of the latent heat of steam. If steam be passed into a kilogram of water at 0° C. till the temperature of the latter reaches 100° C., it will be found that the weight of the water has increased to 1.186 kilograms ; in other words, 0.186 kilogram of steam at 100° C., in being converted into water at 100° C., gives off heat sufficient to raise the temperature of 1 kilogram of water through 100° C.; therefore, 1 kilogram of steam will raise 5.37 kilograms of water through 100° C. or 537 kilograms through 1° C.; or 1 gram of steam, will raise 537 grams of water through 1° C. The latent heat of steam is therefore 537 calories. Steam has the highest latent heat of all known vapors. It is this which renders it such a valuable heating agent when the heat has to be carried to a distance from its source. The phenomena of latent heat, both of liquids and vapors, were first observed and studied by Black.* Liquefaction of Gases.—The fact that the non-saturated or super- heated vapors of liquids behave like true gases leads naturally to the converse idea that the gases may be nothing more that the superheated vapors of liquids unknown under ordinary conditions of temperature and pressure. There are two methods of condensing a vapor to a liquid, one being refrigeration, and the other pressure; pressure hav- ing, as we have already seen, the effect of raising the boiling-point of the liquid. This last method was that chiefly employed by the earlier experimenters in this field, of whom Faraday may be mentioned as the chief. Faraday’s earlier method consisted in generating the gas to be liquefied from some suitable substance contained in one of the limbs of a bent sealed glass tube. The other limb was immersed in cold water, and in this extremity of the tube the gas, liquified by its own pressure, condensed. In this way Faraday succeeded in liquifying chlorine, cyanogen, ammonia, and some other gases. In his later experiments, however, he combined cold with pressure, and thus liquefied carbonic anhydride, nitrous oxide, and other gases. There were, however, a number of gases—oxygen, hydrogen, nitrogen, carbonic oxide, nitric oxide, and marsh-gas—which till quite lately defied all efforts to re- duce them to the liquid state. The reason of this was, that the earlier experimenters relied chiefly on pressure to produce liquefaction, and it was not till the discovery of the phenomenon of the critical point by Andrews, that it became evident that at ordinary temperatures no amount of pressure could liquefy these gases.f Now, however, by the united agency of intense cold and enormous pressure, the problem has * The expression “ latent heat,” though still in very general use, must be regarded as a survival, as it no longer expresses the views of physicists regarding this phenome- non. The heat which has disappeared as such in the above process is no longer heat, and ought not, properly speaking, to be called by this name. It has performed the work of overcoming cohesion; it is no longer present in that form of molecular vibra- tion recognizable as heat, and possibly exists only as the potential energy of position of the molecules. It would be just as admissible to apply the epithet “latent” to the heat which disappears when a steam-engine is employed to raise a weight, because the potential energy of the raised weight can be reconverted into heat, f See, however, p. 121, footnote. 124 INORGANIC CHEMISTRY. been solved simultaneously by two workers in this field, MM. Pictet and Gailletet. (See Hydrogen.) To give an idea of the difficulties to be surmounted in these experiments, it will suffice to mention that oxygen required a pressure of 300 atmospheres and a temperature of —llo° C. (—l66° F.), for its liquefaction,* and that hydrogen did not succumb till a pressure of 650 atmospheres, coupled with a temperature of —l4o° C. (—22o° F.), had been reached. In the descriptions of the various gases the temperatures and pressures of liquefaction will be given. CHAPTER XVIII. SOLUTION. Solubility is the property which many substances—gaseous, liquid, and solid, possess of mixing homogeneously with some liquid employed as a solvent. Gaseous and solid bodies, when in solution, assume for the time being the liquid state. Solubility of Gases.—The solubility of gases is known as absorption. Some gases, such as hydrogen and nitrogen, are soluble in water to a very slight degree only ; others, like carbonic anhydride, chlorine, and sulphuretted hydrogen, are dissolved in moderate quantity; whilst others again, like hydrochloric acid and ammonia, are extremely soluble, the volume absorbed being in the case of the last-mentioned gas at 0° more than a thousand times that of the water employed. In the case of gases slightly or only moderately soluble, the quantity absorbed is approximately proportional to the pressure. This fact may be accounted for by the assumption that the gas occupies the spaces between the molecules of the liquid as it would any other empty space: the quan- tity which can be pressed into this space will then be proportional to the pressure. The solubility generally decreases as the temperature rises. Hence this law may be expressed by saying that the volume of these gases absorbed is constant for a given temperature, being less for higher temperatures, and independent of the pressure. For those gases which are very soluble, this law does not hold. In these cases, the solubility is the result of a powerful affinity between the molecules of the gas and those of the solvent. Such absorptions are accompanied by great evolution of heat—partly the latent heat of the gas, partly the heat of chemical combination. Solubility of Liquids- -Miscibility.—The following views on solubility have been enunciated by Dossios: Let there be two liquids A and B, and let the single molecules of each be represented by a and b respect- ively, and let the attraction of similar molecules be expressed by an, bb, * According to the still more recent results of Wroblewski and Olzewski, oxygen liquefies at the somewhat lower temperature of —l36° under a pressure of only 22.5 at- mospheres. SOLUTION. 125 and that of dissimilar molecules by ab. Then if ah be greater than aa -j- hb, the liquids will obviously be miscible in all proportions. But if ab be less than aa -j- bb, the attraction ab can effect the mixture of the two liquids only with the aid of the energy of their molecules. At the surface of separation of the two liquids, single molecules of A will sometimes be carried, by the force of their own motion, among the molecules of B, where they will wander about until they happen again to reach the surface of separation, when they will for the most part be retained by the other molecules of A. At length a condition will be reached in which as many molecules a return to A as leave it, and as this is the case, JB is saturated with A. The same holds in regard to the saturation of A with B. Two such liquids will dissolve in each other only up to a certain point. An example of this is afforded by the behaviour of ether and water towards each other. If equal volumes of these liquids be agitated together, the ether dissolves about of its bulk of water, whilst the water takes up | of its bulk of ether. When two liquids are miscible in all proportions, the force which comes into play is the preponderating attraction of dissimilar molecules. The heat which is liberated by the approximation of these dissimilar molecules will therefore be greater than that absorbed in the separation of similar molecules. Hence, in most cases where two liquids are miscible in all proportions, heat is evolved by their mixture. A remarkable exception to this rule is presented by a mixture of equiv- alent proportions of ethylic oxalate and araylic iodide, a depression of temperature amounting to 9-3° occurring when the liquids are suddenly blended. Solubility of Solids.—Let A be a solid body, and B a liquid, and let the single molecules and their attractions be designated as above. Then the forces which strive to prevent solution will be aa and bb, those which tend to induce it, ab, and the energy of the molecules. The attraction ab must be less than aa, otherwise the liquid and the solid would form a solid compound. The molecules a, are carried away from A by their energy, plus the attraction ab, wander through the liquid and sometimes return to A. When as many molecules return to Ain unit of time as leave it, the solution is saturated. As the projection of the molecules of A among those of B is dependent in part on the molecular energy, it is evident that the solubility will in- crease with the temperature. This is generally found to be the case; the cause of some apparent exceptions to this rule will be mentioned later. The diagram (p. 126) is a graphic representation of the relations between temperature and solubility in the case of various salts, the sol- vent being water. The abscissae express the temperatures; the ordi- nates, the number of parts of anhydrous salt soluble in 100 parts of water. The method of using this diagram will be evident on inspection. Thus at 0° C., 100 parts of water dissolve 26 parts of magnesic sul- phate; at 40° C., 45 parts; at 100°, 74 parts. As the increase of solu- bility of magnesic sulphate is proportional to the increase of tempera- ture, the line representing its solubility will be straight. The more 126 INORGANIC CHEMISTRY. rapid the increase of solubility in a salt, the more its curve will ap- proach the vertical; the slower this increase, the more nearly horizontal the curve will be. In the ease of sodic chloride, which is almost equally soluble at all temperatures, the curve is nearly horizontal. If the solubility increases more rapidly than the temperature, the curve will show this by bending upwards. In the case of potassic nitrate Fig. 1.—Solubility of Salts in 100 Parts of Water. Parts dissolved. Temperature. and plumbic nitrate the solubility at 0° ot these two salts in 100 parts of water is 13 and 40 parts respectively; at 45° C., both salts are equally soluble, 100 parts of water dissolving 85 parts of each ; whilst at 73° C., the solubility of potassic nitrate is 150 parts against 108 parts of plumbic nitrate. Thus, by rise of temperature, the relative SOLUTION'. 127 solubilities of these two salts have been reversed, the more soluble be- coming the less soluble. This is shown in the diagram by the inter- section of the curves. The point of intersection indicates the tempera- ture of equal solubility. The solubility of sodic sulphate presents a singular anomaly. At 0° C., the solubility in 100 parts of water is 5 parts; it increases more rapidly than the temperature, till at 33° C., it is 51 parts; then it sud- denly decreases, and goes on decreasing the higher the temperature rises. This anomaly would be quite inexplicable, if we were forced to assume that it is the same body which is contained in the solution above and below 33° C.; but closer examination shows this assumption to be un- necessary. Below 33° C., the solution deposits crystals of the formula SO2lSao2,100H2; above this temperature the salt which separates out possesses the formula SO2Nao2,0H2.* The latter salt is less soluble than the former, hence the change in the solubility. The higher the temperature, the greater the quantity of S02Nao2,100H2 which dissoci- ates into S02Nao2,0H2 and water. There is no difficulty in conceiving that a salt may exist in different states in its solutions, at one time with more, at another time with less water of crystallization. Anhydrous cobaltous chloride is blue, as is also the aquate CoC12,20H2; whilst the aquate CoC12,60H2 is pink, and dissolves in water with this color. If to a concentrated aqueous solution of the pink salt a dehydrating agent— strong hydrochloric acid, or absolute alcohol—be added, the solution becomes blue. If less alcohol be added, the solution remains pink in the cold; but on heating, the color changes to blue, and on cooling returns to pink again. Here we have a dissociation perfectly analo- gous to that of the higher aquate of sodic sulphate, the presence of the anhydrous cobaltous chloride (or of the lower aquate) being denoted by the change of color in the solution. Solution is almost invariably attended with contraction, the volume of the substance dissolved, together with that of the solvent, being greater than that of the resulting solution. The only known exception among anhydrous salts occurs in the case of amnionic chloride, the so- lution of which is accompanied by expansion. The most marked con- traction is displayed by dehydrated salts which form definite compounds with water. Contraction also takes place when asolution of a substance is further diluted with the solvent. Solution is attended with absorption of heat. In thosecases in which heat appears to be liberated, the substance enters into definite chemical combination with the solvent, in which process heat is evolved. The compound thus formed dissolves with absorption of heat. The excess of thermal effect due to chemical combination produces the rise of tem- perature. Caustic potash (KHo) dissolves in water with liberation of great heat. But the crystalline aquate KH0,20H2, which is obtained by cooling a concentrated solution of caustic potash, dissolves in water with absorption of heat. The absorption of heat which attends solution is for the most part attributable to the latent heat of liquefaction of the substance (see Latent * Generally st:,(ed to be anhydrous. See, however, Thomsen, Deut. chum. Ges. Her., 11, 2042. ' 128 INORGANIC CHEMISTRY. Heat of Fusion). It is difficult to give an exact account of the various thermal items which go to make up the total thermal effect of solution, as the process is of a complex nature. The explanation formerly in vogue, according to which the fall of temperature during solution is entirely due to the latent heat of liquefaction of the substance, solution itself being caused by the excess of affinity of solvent for substance over that of substance for substance plus that of solvent for solvent, is mani- festly untenable. According to this explanation, solution itself would always be accompanied with liberation of heat, the absorption of heat which is observed being attributable to the excess of heat which becomes latent in the liquefaction of the substance. The absorption of heat during solution would therefore be less than the latent heat of fusion. But very often the reverse is the case. The latent heat of fusion of 1 gram of potassic nitrate is 49 calories ; but by dissolving the same weight of this salt in 20 grams of water at 20° C., 81 calories are absorbed. Supersaturation or Suspended Crystallization.—When a solution con- tains at a given temperature more salt than the coefficient of solubility of that salt indicates, the solution is said to be supersaturated, or the crystallization is said to be suspended. The phenomenon is analogous to that of suspended solidification, observed in the case of fused solids. It occurs most readily with salts which form more than one aquate, and is unknown in the case of anhydrous salts. It may be induced by dis- solving, with the aid of heat, a salt which has a tendency to form a supersaturated solution, and allowing the clear liquid, which must be free from undissolved substance, to cool, excluding dust. On dropping into such a solution a crystal of the aquate which would be formed at that temperature, crystallization immediately ensues with elevation of temperature, the latent heat of liquefaction being evolved. A salt well suited for this experiment is sodic sulphate. No other aquate or modi- fication of a salt than the one which is formed at the given tempera- ture will induce crystallization ; thus sodic sulphate of the formula SO2Nao2,0H2, crystallized above 33° C., may be added to a supersatu- rated solution of sodic sulphate at ordinary temperatures without effect; whilsttheadditiou of the smallest fragmentof theaquate SO2Nao2,100H2 causes instantaneous crystallization. CHAPTER XIX. DIFFUSION. If water be carefully poured on a concentrated solution of a salt con- tained in a tall glass vessel, the liquids will be seen to form two distinct layers, the specifically heavier solution of the salt remaining at the bottom. After standing for some time, however, the salt will be found to be equally distributed throughout the liquid. If a solution of a colored salt, such as cupric sulphate or potassic dichromate, be employed, the progress of this distribution or diffusion, as it is termed, will be rendered DIFFUSION. 129 visible to the eye by a gradation of shades, extending from the bottom to the surface of the liquid, and ranging through every intermediate tint from the color of the concentrated solution to absolute colorlessness. At last, when the process of diffusion is complete, the liquid will exhibit a uniform tint throughout. In like manner, if two tall glass vessels be placed mouth to mouth, one over the other, and separated by a glass plate, the upper being filled with air and the lower with chlorine, then, if the glass plate be carefully withdrawn, the lower vessel will be seen to be filled with the yellowish- green chlorine, whilst the gas in the upper vessel is colorless. But after a short time, the yellowish-green color will begin to extend into the upper vessel, and this will continue until the entire gas presents one uniform tint. The upward progress of the chlorine may further be made visible by the gradual bleaching of a strip of moist carmine-paper attached to the inside of the upper vessel and extending from top to bottom. In both these cases, the force of diffusion is sufficient to overcome the counteracting force of gravity. The heavier molecules of the salt find their way upwards through the lighter molecules of the water; the lat- ter penetrates downwards, diluting the concentrated solution. Chlorine is nearly two and a half times heavier than air; yet its molecules grad- ually rise through those of the oxygen and nitrogen of the air, whilst the latter find their way into the lowest parts of the vessel. In both experiments the ultimate result is uniform mixture. This diffusion has its source in the independent motions of the mole- cules. These motions have already been referred to on various occa- sions in this introduction, while discussing the gaseous and liquid states of matter. The phenomena of diffusion were first thoroughly investigated by Graham, to whom is due the deduction of various important laws in regard to this subject. Diffusion of Liquids.—The quantities of a salt which pass in equal times from a solution into the adjacent water are proportional to the weight of salt originalJy in solution. (This law does not hold for very concentrated solutions.) Rise of temperature increases the velocity of diffusion. This must evidently be the case, as the velocity with which the molecules move increases with the temperature. Different substances have different velocities of diffusion. Isomor- phous salts frequently possess equal velocities of diffusion. Mixed solutions of salts, which do not act chemically on each other, do not diffuse at the same rates as when separate, the difference in their rates of diffusion being increased by mixture. Double salts may fre- quently be decomposed by means of the unequal velocity of diffusion of their component single salts. Dialysis.—ln the course of his investigations on the diffusion of liquids, Graham made the remarkable discovery that certain substances when in solution diffuse through porous membranes, such as bladder or parchment, whereas others do not possess this property. He found further, that the substances which thus diffuse are always crystal lizable, 130 INORGANIC CHEMISTRY. whereas those which are unable to pass through the membrane are amorphous. He thus divided all substances into crystalloids and colloids (from xoUa, glue), and founded upon the above observations a method of separating these two classes of substances. This method, to which he gave the name of dialysis, is carried out as follows: A piece of blad- der or parchment paper is tied tightly over the bottom of a glass cylin- der open at both ends. The liquid to be dialyzed is poured into the cylinder, so as to rest on the membrane, the lower surface of which is kept in contact with water. The crystallizable substance diffuses freely through the membrane and mixes with the water, whilst the colloid remains in the cylinder. By constantly changing the external water, a pure solution of the colloid may be ultimately obtained. The explanation of the phenomenon is as follows: The porous mem- brane, although itself insoluble, takes up water. This may be shown by the great increase in bulk which a piece of bladder undergoes when placed in water. Through the medium of this absorbed water the mole- cules of the crystalloid are enabled to diffuse. It is possible that the molecules of colloids, on the other hand, are much larger, or are aggre- gated into small masses, so that they are unable to pass through the pores of the membrane. The membrane must itself be a colloid. Dialysis has been performed with an artificial membrane of amorphous silicic acid. Diffusion of Gases.—Gases may diffuse either freely into each other, as in the experiment already mentioned, or through very fine openings. A porous diaphragm of gypsum or compressed graphite constitutes a system of such fine openings. Owing to the exceedingly small dimen- sions of the molecules of a gas, they pass through the pores of such a diaphragm almost unimpeded. The law of free diffusion, and of dif- fusion through diaphragms, is the same, and may be stated to be as follows: The velocities of diffusion of any two gases are inversely as the square roots of their densities. Thus the densities of hydrogen and oxygen are as 1 :16, and their velocities of diffusion are as 4: 1. The kinetic theory of gases informs us that the mean velocities of the mole- cules of any two gases are inversely proportional to the square roots of their densities. The above law may therefore also be expressed : The velocities of diffusion of any two gases are directly as the mean velo- cities of their molecules. The extreme velocity with which hydrogen diffuses may be well shown by the following experiment: A tube, closed at the upper end with a thin plate of gypsum, is filled with hy- drogen, and the lower end is plunged into water. Since the hydrogen passes out through the pores of the gypsum much more rapidly than the air can enter, the water rises in the tube. The degree of agreement between theory and experiment for the above law will be seen from the following table, which contains deter- minations of the velocities of diffusion of some of the commoner gases. In these experiments the gas to be examined was contained in a tube, closed at one end with a porous plug of gypsum, and at the other with mercury or water, according to the nature of the gas. The quantity of the gas which escaped through the porous diaphragm, and the quantity of air which entered, were carefully determined. In this way it was CRYSTALLOGRAPHY. 131 found that if the density of a given gas, referred to air as unity, be d, then the volume of this gas which diffuses in the same time as one volume of air, is equal to \/l, as expressed in the foregoing law. This calculated value is given in the third column, and the observed volume in the fourth column of the table : Name of gas. Density of gas = chloric acid on borax : B405Nao2 + 2HCI + 50H2 = 4BHo3 + 2NaCI. Borax. Hydrochloric acid. Water. Boric acid. Sodic chloride. One part of borax is dissolved in 2J parts of boiling water and an excess of concentrated hydrochloric acid is added. On cooling, the boric acid crystallizes out in thin plates. For laboratory purposes boric acid is best prepared by recrystallization of the commercial acid. Properties.—Boric acid, as crystallized from water, forms lustrous laminae, unctuous to the touch. One hundred parts of water at 10° C. dissolve 2 parts, at 100° C., 8 parts, of boric acid. The solution turns blue litmus wine-red, and turmeric paper, even in presence of hydro- chloric acid, brown. When the aqueous solution is boiled, the boric acid volatilizes with the steam, as in the soffioni. Boric acid is also slightly soluble in alcohol, and communicates to the flame of the alcohol a characteristic green coloration. At a temperature of 100° C., boric acid parts with the elements of water, and is converted into metaboric acid, BOHo: BHo3 = BOHo + 0H2. Boric acid. Metaboric acid. Water. Metaboric acid forms stable salts, such as sodic metaborate (BOISTao) I fBO \ and magnesic metaborate j< Mgo" . UBO / W hen boric acid is heated for a long time to 140° C. (284° F.) tetraboric acid is formed as a brittle vitreous mass: 4BHo3 = BjOgUoj -f- 50H2. Boric acid. Tetraboric acid. Water. The tetraborates are also stable compounds. Anhydrous borax (B4OgJSTao2) is sodic tetraborate. The normal borates or orthoborates, derived from the tribasic acid (BHo3 ), are the least stable of the compounds of boric acid. C A EBON. 193 BORIC SULPHIDE. Molecular weight = 118. B2S"3. Preparation.—Boric sulphide is formed when the vapor of sulphur is passed over heated boron; but it is best prepared by heating to bright redness a mixture of lamp- black and boric anhydride in a current of carbonic disulphide vapor; 2B203 + 80S", + 3C = 2B2S//3 + 6C"O. Boric Carbonic Boric Carbonic anhydride. disulphide. sulphide. oxide. Properties.—Boric sulphide is thus obtained as a solid, yellowish-white, fusible, vit- reous mass, which may be volatilized in a current of sulphuretted hydrogen, and then forms silky needles. It has a pungent odor, and its vapor irritates the eyes. Water at once decomposes it into sulphuretted hydrogen and boric acid: BgS", + 60 H2 = 3SH2 + 2BHo3. Boric Water. Sulphuretted Boric sulphide. hydrogen. acid. CHAPTER XXV. TETEAD ELEMENTS. Section I. CARBON, C. Atomic weight = 12. Atomicity " and iv. Evidence of atomicity: Carbonic oxide, . . . . 0"O. Carbonic tetrachloride, . . . . c-ci,. Marsh-gas, . . . . CivH4. Chloroform, . . . . CivHCl3 Occurrence.—Carbon exists in the free state in three distinct allo- tropic modifications, as amorphous carbon, as graphite, and as diamond, all of which are found in nature. In combination with oxygen as car- bonic anhydride, it occurs in the air. It is a constituent of all organic substances, and upon its varied combining powers the infinite mani- foldness of the animal and vegetable kingdoms ultimately depends. General Properties.—The following properties are common to carbon in all its modifications: It is solid, infusible, probably non-volatile at the highest temperatures that can be artificially produced, and insoluble in all known solvents at ordinary temperatures. a. Amorphous Carbon.—The chief varieties of amorphous carbon are: Charcoal, lamp-black, gas-carbon, and coke. Occurrence.—Amorphous carbon is found in nature as mineral char- coal. 194 INORGANIC CHEMISTRY. Charcoal. Preparation.—When wood is heated to redness in dosed vessels, the cellulose (C6HIOOS)x gives off its oxygen and hydrogen, partly as water, partly along with a portion of the carbon in the form of oxides of car- bon and of more or less complex organic compounds. When these various gaseous and liquid products of destructive distillation have ceased to be evolved, the charcoal remains behind in the retort as a black, non-lustrous substance, preserving the form of the wood from which it was prepared. The liquid products of distillation constitute wood-tar, and their nature will be described under Organic Chemistry. In order to obtain the greatest possible yield of charcoal, care must be taken to expel all moisture from the wood before raising the tem- perature to redness, otherwise the charcoal at a red heat will decompose the water, forming carbonic anhydride or carbonic oxide and liberating hydrogen. The distillation is performed in cast-iron retorts. The wood to be carbonized is placed in a perforated iron case F (Fig. 31), known as a slip, which is then introduced into the retort A. The volatile products Fig. 31. of decomposition are led by the pipe L into the furnace B, where they are burned, a saving of fuel thus being effected. In well-arranged works at the present day, these products are condensed, acetic acid and wood-naphtha being obtained from them. One hundred parts of wood yield on an average 27 parts of charcoal. In countries where wood is plentiful, a method of carbonizing in heaps is employed, the heat being produced by the combustion of a part of the wood. This is the oldest process of charcoal-burning. The logs are piled on end in a heap (Fig. 32), and a space is left in the middle to serve as a flue. The whole is covered with turf and earth, small apertures being made at the base of the heap to admit air. Fire is applied from below, and the action of the heat is carefully regulated by opening or closing the air-holes in different parts of the heap. The charcoal obtained by this method is inferior in quality to that produced by carbonizing in retorts. A very pure charcoal for special laboratory purposes is obtained by carbonizing sugar in a closed platinum vessel. If it is necessary to get CARBON. 195 rid of the last traces of hydrogen, the product must be strongly ignited in a current of chlorine. This charcoal possesses the advantage of contain- ing no silica, and may therefore be employed in the preparation of volatile chlorides (see boric chloride, p, 188), which would otherwise be contaminated with silicic chloride (SICI4). Another variety of charcoal is animal charcoal or bone-hlach, pro- duced by the carbonization of bones in closed vessels. A fetid oil of very complex character distils over during the process. The charred mass which remains in the retort is afterwards coarsely granulated, in Fig. 32. which form it is employed to decolorize liquids. Animal charcoal which has lost its decolorizing properties by repeated use may have these restored by again heating it in closed vessels. A very pure ani- mal charcoal is obtained by carbonizing dried blood which has been mixed with potassic carbonate, in order to render the product more porous. The potash is afterwards extracted with hydrochloric acid. Properties.—The qualities of the product vary with the temperature employed. The best wood-charcoal for laboratory and metallurgical purposes is prepared at a high temperature, and is a hard brittle sub- Fig. 33. stance with a lustrous fracture. When struck, it emits a metallic sound. Common charcoal is a bad conductor of heat and electricity; but by exposing it for a long time in closed vessels to a very high temperature, it becomes an excellent conductor. The elimination of the oxygen and hydrogen from the wood in the formation of charcoal leaves the mass in an extremely porous condition, 196 INORGANIC CHEMISTRY. and the infusibility of the charcoal causes it to retain this porosity. A very small piece of charcoal may thus expose an enormous surface, and hence all phenomena dependent upon surface action are displayed in a high degree by this substance. To this class belong the condensation of gases and decolorizing of liquids. The absorbent power of wood-charcoal for gases may be shown by cooling a fragment of freshly ignited charcoal under mercury, and then passing it into a tube filled with gaseous ammonia over the mer- curial trough (Fig. 33). The mercury will rapidly rise in the tube as the ammonia is absorbed. The following list gives the volumes of some of the principal gases absorbed by one volume of boxwood-char- coal at 0° C., and under a pressure of 760 mm., as determined by Hunter: Absorption of gases by charcoal— Yols. Hydrogen, ...... . . . 4.4 Nitrogen, . . . 15.2 Oxygen, . _ Carbonic oxide, . . . 17.9 . . . 21.2 Carbonic anhydride, . . . . 67.7 Nitric oxide, . . . 70.5 Nitrous oxide, . . . 86.3 Ammonia, ...... . . . 171.7 As a rule the most easily liquefiable gases are absorbed in greatest quantity by charcoal. Noxious effluvia are in like manner absorbed by charcoal, and at the same time undergo oxidation at the expense of the oxygen condensed in its pores, a property which has led to the use of charcoal for disin- fect! ng purposes. The property of decolorizing liquids depends upon the absorption of the coloring matter in the pores of the charcoal. Animal charcoal is best suited for this purpose, inasmuch as the inorganic matter contained in the bones increases the porosity of the product. If a red wine be warmed with freshly ignited animal charcoal and then filtered, the filtrate will be colorless. In the process of sugar refining the raw syrup is decolorized by filtration through animal charcoal. Charcoal filters are also employed for the purification of water for drinking purposes, but they are not to be recommended, owing to the stimulus which animal charcoal gives to the development of animalcularlife. Lamp-black.—When certain organic substances rich in carbon, such as resins, essential oils, and heavy hydrocarbons, are burned in air, the supply of oxygen is insufficient for complete combustion, and the flame smokes. A porcelain dish or any cold object held in the flame is quickly covered with a finely divided black deposit. This is the substance known as lamp-black. On a large scale, the tar, resin, or other highly carbonaceous substance is burnt with a limited supply of air, and the heavy smoke is made to pass through chambers, where the lamp-black settles. Lamp-black, after strong ignition in a stream of chlorine in order to CARBON. 197 free it from the hydrogen which the ordinary product always contains, is one of the purest forms of amorphous carbon. Lamp-black is employed in the manufacture of printing ink and China ink, and also as a common black paint. Coke.—When coal is subjected to destructive distillation in the manufacture of coal-gas, a number of volatile products are expelled, and an impure amorphous carbon, known as cohe, remains in the retort. Coke is also prepared by burning coal in heaps, as in the conversion of wood into charcoal; but in the coking-heap the central flue is built of fire-bricks. The coking is thus effected by the combustion of a portion of the coal. As soon as smoke ceases to be given off, the air-holes at the bottom of the heap are closed with wet sand, or, more frequently, the fire is quenched with water. At the present day most of the coke is obtained by partially burning the coal in specially constructed coking ovens. The coke prepared in ovens is denser and of better quality than that obtained by other means. Coke does not ignite readily, nor is its combustion well maintained, except in large masses and with the aid of a rapid current of air ; but its combustion produces a very high temperature, and is unattended with the production of smoke. It is largely used in iron smelting and other metallurgical operations. Gas Carbon.—This substance is also produced in the manufacture of coal-gas. When the heavier hydrocarbons formed from the coal pass over the red-hot walls of the retort, they deposit a portion of their carbon in an exceedingly dense and coherent form. The gas carbon so obtained forms a gray, very hard mass, possessing a metallic lustre. It is an excellent conductor of heat and electricity. The carbon-plates of the Bunsen battery, and sometimes the carbon-rods for the electric arc-light, are made from this material, A very pure form of amorphous carbon is obtained by the action of potassium at a high temperature on carbonic anhydride or a carbonate: 3C02 + 2K2 = C + 2COKo2. Carbonic anhydride. Potassic carbonate. The carbon must be carefully washed with hydrochloric acid to free it from the last traces of alkali. Reaction.—By treatment with a mixture of potassic chlorate and fuming nitric acid, amorphous carbon is converted into brown com- pounds soluble in water. Potassic permanganate, in alkaline solution, or nascent electrolytic oxygen, converts it into raellitic acid, and other products. Coal.—This substance consists of the remains of a former flora. It is the result of a decomposition which woody fibre has undergone during long geological periods under varying conditions of temperature and moisture, and with exclusion of air. Under these circumstances the hydrogen and oxygen of the wood have been gradually reduced in quantity by elimination, partly as water and partly in combination with a portion of the carbon as methylic hydride (the fire-damp of the miner) and carbonic anhydride. The process is thus very similar to that which occurs when wood is converted into charcoal by heating in closed 198 INORGANIC CHEMISTRY. vessels. The degree of change which the woody fibre has undergone varies with the age of the coal: thus lignite, a more recent forma- tion, preserves its fibrous structure and contains a large percentage of oxygen and hydrogen; whereas anthracite, which is found in the oldest carboniferous deposits, is dense and amorphous, and contains a very high percentage of carbon. The following is a list of some of the chief varieties of coal: Lignite or Brown Coal is generally of more recent date than the chalk formation; whilst true coal is older than the chalk. Its specific gravity is also lower than that of true coal. It yields a powdery coke and burns with a comparatively smokeless flame. Bituminous or Caking Goal.—The greater number of English coals belong to this class. Bituminous coal fuses and cakes together on heat- ing, giving off much smoke and gas, and yielding a lustrous coke. Cannel coal is a variety of bituminous coal. It contains a large per- centage of hydrogen, and is much in request for purposes of gas manu- facture. Anthracite.—This is a very hard coal with a conchoidal fracture. It is of an iron-black color, with a semi-metallic lustre, and its smooth surface frequently displays iridescence. It splinters when heated, and ignites with difficulty, burning with very little flame and no smoke, and giving out an intense heat. It is much used as a steam coal and also for smelting purposes. The following table shows the average composition of coals from different localities in Great Britain. The last column contains the thermal effect as measured by the number of pounds of water at 100° C. which were found to be converted into steam in a Cornish boiler by 1 pound of the coal : Table showing the Average Composition of Coals from different Localities. Locality. Sp. gr. Carbon. Hydro- gen. Nitro- gen. Sul- phur. Oxygen. Ash. Percentage of coke left by each coal. Evap- orating power Wales, . . Durham, . Lancashire, Scotland, . Derbyshire, 1.315 1.256 1.273 1.259 1.292 83.78 82.12 77.90 78.53 79.68 4.79 5.31 5.32 5.61 4.94 0.98 1.35 1.30 1.00 1.41 1.43 1.24 1.44 1.11 1.01 4.15 5.69 9.53 9.69 10.28 4.91 3.77 4.88 4.03 2.65 72.60 60.67 60.22 54.22 59.32 9.05 8.37 7.94 7.70 7.58 /?. Graphite. Occurrence.—This variety of carbon constitutes the mineral plumbago or black-lead. It is found in various crystalline rocks, such as granite, gneiss, and piorite. It is possibly of vegetable origin, and in this case corresponds to the most complete transformation of vegetable substance, inasmuch as it never contains more than traces of hydrogen. The geological formations in which it occurs are likewise much older than the carboniferous strata. Preparation.—-1. When the diamond is exposed to the heat of the electric arc in an atmosphere devoid of oxygen, it swells up and is CARBON. 199 converted into a black mass of graphite. The various forms of amor- phous carbon are also converted into graphite under these conditions. 2. Cast iron is a compound of carbon and iron. In the molten state the iron dissolves more carbon than is required for combination, and, on cooling, this excess separates out as crystalline scales of graphite. When gray pig-iron is dissolved in an acid, the graphite remains behind. Properties.—Graphite crystallizes in six-sided plates, in which form it sometimes occurs in nature; but it is more frequently found in granular, foliated, or fibrous masses. The natural variety is grayish- black, with a metallic lustre (hence the name black-lead), and is unctu- ous to the touch. Its specific gravity varies from 1.8 to 2.4. It is soft enough to leave a mark on paper, a property which is turned to account in the manufacture of black-lead pencils. It conducts heat and electricity well. Reaction.—lf 1 part of pure graphite be heated for some days on a water-bath to 60° with 3 parts of potassic chlorate and sufficient con- centrated nitric acid to render the whole fluid, a portion of the graphite is converted into graphitic acid (CnH4Og), and by the repetition of this treatment pure graphitic acid may be obtained in thin yellowish trans- parent crystals (Brodie). When heated, graphitic acid decomposes with violence, evolving gas, and yielding a very bulky finely-divided black powder of pyrographitic oxide (C22H204) which is dissolved by a mix- ture of potassic chlorate and nitric acid. Baric graphitate detonates violently when heated. Applications.—Graphite is chiefly employed in the manufacture of black-lead pencils. Other uses are: the coating of iron-work as a pre- servative against rust, the polishing of gunpowder, the lubrication of machinery, and the preparation of plumbago crucibles. y. Diamond. Occurrence.—This gem is found in alluvial deposits produced by the disintegration of a particular micaceous rock known as itacolumite. It has also been found in matrix in the rock itself. The principal diamond fields are those of Brazil and the Cape of Good Hope. Among all the known allotropic modifications of the elements, the diamond is remarkable as the only one which has not been produced artificially. Properties.—The diamond crystallizes in forms derived from the regular octahedron. The faces of the crystals are very frequently con- vex. The finer specimens are transparent and colorless. Colored varieties are not uncommon. It possesses a characteristic and brilliant lustre, known as the adamantine lustre, which is due to its very high refractive and dispersive power. This lustre is artificially intensified by cutting. The diamond is the hardest of known substances, and can be cut only by means of its own dust, the gem being pressed against a revolving steel plate covered with diamond-dust and oil. Its specific gravity is 3.55. It is a non-conductor of electricity. In closed vessels, it may be heated to very high temperatures with- out undergoing change, but when subjected to the heat of the electric arc it is converted into graphite. When intensely heated in air or 200 INORGANIC CHEMISTRY. oxygen it burns, forming carbonic anhydride, and leaving a small quantity of ash. The diamond contains neither hydrogen nor oxygen. The diamond is not attacked by a mixture of potassic chlorate and nitric acid. Unlike boron and silicon, the diamond does not dissolve in molten aluminium. Applications.—Besides its well-known use as an ornament, the dia- mond is employed in the arts. Diamonds are used for cutting glass, for which purpose only the natural curved edge of the crystal is able, as the cut or broken diamond merely scratches the glass super- ficially. The rock-boring apparatus employed in tunnelling and well- sinking is frequently fitted with diamonds set in the edge of a steel ring. Diamond-dust is the best grinding and polishing material for hard substances. For these purposes, inferior varieties of diamond may be employed. The optical properties of the diamond have caused it to be used for microscopic objectives; but the great difficulty of grinding lenses of so refractory a material has limited this application. COMPOUNDS OF CARBON WITH OXYGEN CARBONIC ANHYDRIDE. 002. Molecular weight = 44. Molecular volume 11 I. 1 litre weighs 22 criths. Fuses at —s7° C. (—70.6° F.) Boils below its fusing point. History.—This gas was discovered by Van Helraont in the seven- teenth century. It was further studied by Black, but its true chemical nature was first demonstrated by Lavoisier. Occurrence.—Carbonic anhydride occurs in small quantity in the atmosphere, to the extent of about 3 volumes in 10,000 volumes of air. All spring-water contains it in solution, and in the case of some springs arising in volcanic districts, the quantity of carbonic anhydride dis- solved is so great as to cause the water to effervesce strongly. In such volcanic districts, the gas is often given off from fissures in the earth, and this continues for thousands of years after the cessation of active volcanic phenomena. Preparation.—l. When carbon is burned in an excess of oxygen or air carbonic anhydride is formed : c + 02 = C02. Carbonic anhydride. Unless an excess of oxygen or air is employed, carbonic oxide is also formed. This method is sometimes employed when carbonic anhydride is required in very large quantities for manufacturing purposes, as in the preparation of white lead. Coke is burnt in atmospheric air for such applications. CARBONIC ANHYDRIDE. 201 2. The method usually employed in the laboratory, for the prepara- tion of this gas in a state approximating to purity, depends on the fact that carbonates are easily decomposed by stronger acids, and that the carbonic acid thus produced instantly breaks up into carbonic anhy- dride and water. Calcic carbonate in its naturally occurring varieties, as chalk or marble, is the salt usually employed for this purpose. The marble, broken into coarse fragments, is introduced into a flask fitted with a funnel and delivery tube as in the apparatus for the preparation of hydrogen (Fig. 16, p. 143), and the flask is half-filled with water. Hydrochloric acid is then poured through the funnel until the gas is evolved in a sufficiently rapid stream ; COCao" + 2HCI = C02 + OII2 + CaCJ2. Calcic Hydrochloric Carbonic Water. Calcic carbonate. acid. anhydride. chloride Other carbonates may be substituted for calcic carbonate and other acids for hydrochloric acid in the above reaction : COKo2 + S02Ho2 = C02 + 0H2 + S02Ko2. Potassic Sulphuric Carbonic Water. Potassic carbonate. acid. anhydride. sulphate. COHoKo + N02Ho = C02 -j- OU2 + N02Ko. Hydric potassic Nitric acid. Carbonic Water. Potassic carbonate. anhydride. nitrate. 3. Sulphuric acid cannot be employed, in the foregoing way, with marble in the preparation of carbonic anhydride, as the insoluble calcic sulphate coats the marble and prevents further action. But if concen- trated sulphuric acid be poured upon chalk and then a little water be added, the gas is evolved in a steady current, as the acid under these conditions produces a disintegration of the chalk. 4. Most carbonates, when strongly heated, evolve carbonic anhy- dride, as for example when chalk or marble is calcined to form quick- lime : COCao" = CaO + C02. Calcic Calcic Carbonic carbonate. oxide. anhydride. The carbonates of the alkali metals are the only exceptions to this rule. Formation.—When any substance containing carbon is burned in air, the carbon is converted into carbonic anhydride, the hydrogen with which the carbon is generally associated forming water. In this way immense quantities of carbonic anhydride are continually discharged into the atmosphere in the combustion of coal and wood. Active combustion is a rapid oxidation. But combined carbon may also undergo slow oxidation with production of carbonic anhydride. Thus the slow oxidation of the animal tissues of the living body produces the carbonic anhydride which is given off from the lungs during respira- 202 INORGANIC CHEMISTRY. tion. This may be shown by breathing through lime-water, which is thus rendered turbid. In fermentation, decay, and putrefaction, processes in which complex chemical changes take place in organic matter under the influence of minute living organisms, part of the carbon of the substance is often evolved along with a portion of its oxygen as carbonic anhydride. Thus in the fermentation of grape-sugar with yeast at a temperature of about 22° : C 6H1206 = 202HSHo + 2G02. Grape-sugar, Ethylic alcohol. Carbonic anhydride. A similar evolution of carbonic anhydride occurs during the forma- tion of coal. Circulation of Carbon in Nature.—All the carbon present, in every form of combination, in the bodies of plants and animals is derived ultimately from the carbonic anhydride of the air. Plants, by means of the chlorophyll, or green coloring matter of their leaves, and with the aid of sunlight, decompose this carbonic anhydride, evolving the oxygen, and retaining the carbon for the purpose of building up their tissues. Animals—the herbivora directly, the carnivora indirectly— derive their entire nourishment from plants. The carbon is thus trans- ferred to the bodies of animals, where it serves, by its oxidation, as a source of vital heat and of energy of motion. The oxygen necessary for this oxidation is absorbed during respiration by the haemoglobin or red coloring matter of the blood, which thus serves as a carrier of oxygen to the tissues; and the carbonic anhydride formed in the oxidation is, as already stated, expelled with the breath and thus finds its way back into the atmosphere. A similar cycle of operations occurs with hydrogen. The plant de- composes, under the same conditions, either the aqueous vapor of the atmosphere or the water contained in its own juices, evolving the oxygen and assimilating the hydrogen, A portion of the oxygen, either from the carbonic anhydride or from the water, or from both, is at the same time retained by the plant. During the oxidation of the animal tissues the hydrogen is for the most part re-oxidized to water, and in this form is exhaled or otherwise expelled from the body. The plant thus inhales carbonic anhydride and aqueous vapor, and exhales oxygen. Animals inhale oxygen and exhale carbonic anhydride and aqueous vapor. In this way the action of the one tends to balance that of the other. Broadly speaking, the functions of the plant may be said to be syn- thetical, those of the animal analytical. Properties.—Carbonic anhydride is a colorless gas, with a slightly pungent odor and an acidulous taste. It does not support either com- bustion or respiration: the flame of a taper, plunged into the gas, is extinguished, and animals are rapidly asphyxiated by it. Its physio- logical action is that of a narcotic poison. In small quantities it may be breathed with impunity; but air containing 0.5 per cent, produces headache and oppression, and the presence of even 0.2 per cent, is suf- ficient to render air unwholesome. 203 CARBONIC ANHYDRIDE. The specific gravity of carbonic anhydride is, according to Regnault, 1.5241 (air = 1). It is thus rather more than one and a half times heavier than air. Owing to its great density it may be collected by displacement, and may be poured from one vessel into another like a liquid. On lowering a taper into the vessel into which the gas has been poured, the flame will be extinguished as soon as it is immersed in the carbonic anhydride. In like manner, if a counterpoised beaker be sus- pended from one arm of a balance (as in the experiment for demon- strating the lightness of hydrogen, with the exception that in the case of carbonic anhydride the beaker is suspended mouth upwards), then, on pouring the heavy gas from another vessel into the beaker, the arm of the balance supporting the beaker will be depressed by the weight of the gas. This property, which causes carbonic anhydride to collect at the lowest level, is sometimes the source of fatal accidents, as in cases where wells or beer-vats containing this gas have been incautiously entered. The phenomena of the Grotto del Cane and of the Poison Valley in Java are due to the same cause. Carbonic anhydride is formed in coal-mine explosions by the combustion of the fire-damp Fig. 34. (methylic hydride, CH4), and it frequently happens that miners who escape the violence of the explosion are asphyxiated by the after-damp. It has been shown, however, that the after-damp generally also contains the much more deadly carbonic oxide. When carbonic anhydride is subjected to a pressure of 36 atmo- spheres at a temperature of 0° C., it condenses to a colorless liquid. The liquefaction of the gas may be conveniently effected by means of the apparatus shown in Fig. 34, devised by Thilorier. Into the strong wrought-iron generator g, hydric sodic carbonate, stirred up with a little over twice its weight of water, is introduced. Sulphuric acid is poured into the inner tube (represented by dotted lines in the figure) 204 INORGANIC CHEMISTRY. and the head of the generator is screwed on. The generator, which swings upon trunnions on the stand s, is then turned over so as to allow the sulphuric acid to flow out of the tube and mix with the hydric sodic carbonate. Carbonic anhydride is liberated according to the equation— 2COHoNao + S02Ho2 = 2002 + S02Nao2 + 20H2, Hydric sodic Sulphuric Carbonic Sodic Water, carbonate. acid. anhydride. sulphate. On bringing the apparatus back into its former position, the carbonic anhydride, liquefied by pressure, rises to the surface and floats as a layer on the solution of sodic sulphate. The generator is then connected by the copper tube t with the wrought-iron receiver r. On opening the screw-taps w and v, applying a gentle warmth to the generator, and cool- ing the receiver, the liquefied anhydride distils over into the latter vessel. The screw-tap v is then closed ; the generator is disconnected, emptied, recharged, and the above operations repeated. Six or seven charges suffice to fill the receiver. The nozzle nis then attached to the receiver in place of the tube t. An improved form of Thilorier’s apparatus has been constructed, in which the liquefied anhydride, instead of being distilled, is forced over in the liquid state into the receiver by means of water pumped in at the base of the generator. Liquid carbonic anhydride is colorless and very mobile. Under the influence of heat it expands more rapidly than any known substance, surpassing even the gases in this respect. The following table shows this rapid alteration of density : Temperature. Sp. gr. —lo° C. (14° F.) .9951 + 0° C. (32° F.) .9470 +2o° C. (68° F.) .8266 Carbonic anhydride at —7B° exerts a pressure of 760 mm. When the liquid is exposed to the air the heat rendered latent by its evapora- tion causes it to solidify. The following apparatus (Fig. 35) is well adapted for procuring solid carbonic anhydride. It consists of a circular brass box in two halves, one of which fits over the other as a lid, each half being furnished with a hollow handle covered with wood or some other bad conductor of heat. Through a small tubular opening in the circumference the nozzle of the screw-tap of the wrought-iron cylinder con- taining the liquefied carbonic anhydride is inserted. On opening the screw-tap a jet of liquid carbonic anhydride is projected with great violence into the brass box, and striking at a tangent to its internal circumference, flows round it, solidifying in the process, and filling the interior with a snow-like Fig. 35. mass. On opening the box, the snowball of solid carbonic anhydride may be removed. Solid carbonic anhydride thus prepared is a coherent white powder, CARBONIC ANHYDRIDE. 205 resembling snow in appearance. It may be exposed for a short time to the air ; but eventually disappears as gas, without previously melting. Though its temperature is so low, it may be touched without incon- venience, as the gas which it evolves forms a non-conducting layer around it; but if it be pressed upon the skin, it produces a blister like that caused by a burn. It is soluble in ether, and in this condition its evaporation can be conveniently employed as a source of cold. When the solution of carbonic anhydride in ether is evaporated in vacuo, the temperature sinks so low as —llo° C. (—l66° F.). By means of the depression of temperature thus produced, liquid carbonic anhydride contained in a tube may be frozen into a transparent ice-like solid. Water at 15° C, (59° F.) dissolves its own volume of carbonic anhydride under a pressure of 760 mm. The quantity of gas absorbed is approximately proportional to the pressure. (See Introduction, p. 124.) If water be saturated with the gas at a higher pressure, and the pressure be suddenly removed, evolution of gas ensues. The solubility of carbonic anhydride decreases rapidly at higher temperatures, and the whole of the dissolved gas may be expelled by boiling. Composition.—l. When carbon is burned in oxygen, it is found that the volume of the carbonic anhydride formed is exactly equal to that of the oxygen employed. It is thus evident that carbonic anhydride contains its own volume of oxygen. In this way the composition by weight of carbonic anhydride may be deduced. Suppose the volume of oxygen employed to have been 1 litre— 1 litre of C02 formed weighs . . 22 criths, Deduct the weight of 1 litre of 0 . . . 16 “ There remain : carbon, . . 6 “ Therefore 6 parts by weight of carbon combine with 16 of oxygen to form 22 parts of carbonic anhydride. Expressed in atomic weights, this gives— Proportion of carbon is to oxygen as 12 : 32, corresponding to the formula C02. 2. The composition by weight of carbonic anhydride can be directly determined by ascertaining the weight of this gas which is formed when a known weight of pure carbon (diamond or purified graphite) is burnt in a current of oxygen. This was the method employed by Du mas and Stas, The oxygen is contained in the Woulff’s bottle a (Fig, 36), from which it is expelled during the operation by dilute caustic potash, this liquid being employed in order to prevent the gas from being contaminated by the carbonic anhydride contained in ordi- nary water. It then passes through three U-tubes b, c, and d, the first containing pumice moistened with strong potash, the second fragments of solid potash, and the third pumice moistened with concentrated sulphuric acid. The oxygen, thus thoroughly freed from carbonic anhydride and moisture, passes on through the glazed porcelain tube ef, which contains the weighed portion of carbon placed in a platinum boat. This tube is 206 IJSTORGANIC CHEMISTRY. heated to redness in the furnace F, and the carbon in the boat thus burns in the current of purified oxgen. As carbonic oxide may be formed in this combustion, the gases are passed through a second tube gh of refractory glass, containing granu- lated cupric oxide, and heated to redness by means of charcoal placed in the iron trough q. In this way any carbonic oxide is converted into carbonic anhydride at the expense of the oxygen of the cupric oxide. The mixture of carbonic anhydride and oxygen passes on through the U-tube k, containing pumice and sulphuric acid ; then through the Liebig’s bulbs m, containing a strong solution of potash, by which the greater part of the carbonic anhydride is absorbed ; then through the tube n filled with pumice moistened with strong potash, in order to absorb the last traces of carbonic anhydride. The tube o, containing fragments of solid potash, serves to arrest any moisture which may be given olf from the tube n. The last tube, p, also containing fragments of solid patash, is intro- duced in order to prevent access of carbonic anhydride and moisture from the air to the tube o. The tubes k, m, n, and o are accurately weighed both before and after the combustion of the carbon. If the experiment has been properly conducted so as to exclude every trace of moisture, and if the carbon employed has been per- fectly free from hydrogen, the tube k ought to show no increase in weight. The increase in weight of the tubes m, n, and o gives the weight of car- bonic anhydride formed. The weight of carbonic anhydride, minus the Fig. 36. weight of carbon employed, gives the weight of oxygen consumed. In this way it has been found that 1 gram of carbon yields 3.666 grams of carbonic anhydride. The weight of oxygen consumed is therefore 2.666 grams, from which it follows that 32 parts by weight of oxygen combine with 12 of carbon to form 44 of carbonic anhydride, a result which exactly coincides with that obtained by the foregoing method. The platinum boat ought to be weighed both before and after the ex- periment in order to determine the weight of ash, which is present in CARBONIC ANHYDRIDE. 207 even the purest forms of carbon. This weight is then deducted from the weight of carbon originally taken. Reactions.—l. Carbonic anhydride is decomposed by the action of intense heat, such as that of the electric spark, into carbonic oxide and oxygen: co2 = 00 + 0. Carbonic Carbonic anhydride. oxide. Only a small portion of the carbonic anhydride is thus decomposed, inasmuch as, when the proportion of the products of decomposition passes a certain limit, they again combine with formation of carbonic anhydride (see Introduction, p. 104). 2, When potassium is heated in an atmosphere of carbonic anhydride, the gas is decomposed with liberation of carbon : 3C02 + 2K2 == 2COKo2 + C. Carbonic Potass ic anhydride. carbonate. 3. Carbonic anhydride acts upon metallic hydrates, forming carbo- nates : . C02 + 2KHo = COKo2 + 0H2. Carbonic Potassic Potassic Water, anhydride. hydrate. carbonate. 0O2 + KHo = COKoHo. Carbonic Potassic Hydric potassic anhydride. hydrate. carbonate. 0O2 + CaHo2 = COCao" + 0H2. Carbonic Calcic Calcic Water, anhydride. hydrate. carbonate. The carbonates are very stable compounds. The alkaline carbonates may be exposed to a white heat without undergoing decomposition; all other carbonates are decomposed at higher temperatures into metallic oxide and carbonic anhydride: COCao" = CaO + 0O2. Calcic Calcic Carbonic carbonate. oxide. anhydride. The alkaline carbonates are soluble in water; all other carbonates are insoluble. Free carbonic acid, OOHo2, is not known in a state of purity, but the solution of carbonic anhydride in water contains this acid. This is shown by the fact that the solution reddens litmus, a property not possessed by carbonic anhydride. Moreover, a solution of carbonic anhydride saturated under pressure loses its gas much more rapidly when freshly prepared than when the saturated solution has been preserved under pressure for some time, as in the case of artificial aerated waters. 208 INORGANIC CHEMISTRY. Tliis seems to denote that at first mere solution takes place, but that in course of time the carbonic anhydride combines chemically with the water: C02 ■+ OBT2 = COHo2. Carbonic Water. Carbonic anhydride. acid. With inorganic bases carbonic acid almost always acts as a dibasic acid, forming acid and normal salts. The acid carbonates of the alka- lies are the only acid carbonates known in the solid state. Ethereal salts of the tetrabasic acid, CHo4, have been prepared (see Organic Chem- istry). Dicupric carbonate, CCuo"2, which occurs as the mineral mysorine, may be regarded as a salt of the tetrabasic acid. Owing to the insolubility of the carbonates of the alkaline earths, lime water or baryta water is rendered turbid by carbonic anhydride, or by a solution of a carbonate. An excess of carbonic anhydride dis- solves the precipitate, owing to the formation of an acid carbonate, which, however, can exist only in solution. CARBONIC OXIDE Molecular weight = 28. Molecular volume I I I. 1 litre weighs 14 criths. Liquefiable by great pressure and cold. History.—Carbonic oxide was discovered by Lasonne in 1776. Preparation.—l. When carbonic anhydride is passed over red-hot charcoal, it gives up half of its oxygen to the charcoal, and carbonic oxide is formed ; COz + C = 2CO. Carbonic Carbonic anhydride. oxide. 2. In like manner red-hot iron lower stage of oxidation : reduces carbonic anhydride to the 4C02 + 3Fe = iv(Fe3)viii04 + 4CO. Carbonic Triferric Carbonic anhydride. tetroxide. oxide. The reaction may be carried out by passing carbonic anhydride over iron turnings contained in a tube of porcelain or iron heated to redness in a furnace. 3. Instead of acting on free carbonic anhydride, this gas may be em- ployed in the nascent state. Thus, if any of the carbonates which evolve carbonic anhydride at higher temperatures be heated to redness with charcoal or iron filings, carbonic oxide will be produced: OOCao" + C = CaO + 2GO. Calcic Lime. Carbonic carbonate. oxide. 209 CARBONIC OXIDE. 4. Carbonic oxide is also formed when ferric or zincic oxide is heated to redness with charcoal: ZnO + C = Zn + CO. Zincic Carbonic oxide. oxide. 5. Concentrated sulphuric acid, from its strong affinity for water, has the power of abstracting the elements of water from a number of organic substances. Thus, when oxalic acid is heated with concentrated sul- phuric acid, water is removed, and a mixture of equal volumes of car- bonic anhydride and carbonic oxide is evolved : {cOHo = 0H° + °°2 + CO- Oxalic acid. Water. Carbonic Carbonic anhydride. oxide. The carbonic anhydride may be absorbed by passing the mixed gases through a strong solution of sodic hydrate. The carbonic oxide is thus obtained in a state of purity. 6. In like manner, when formic acid ora formate is heated with con- centrated sulphuric acid, pure carbonic oxide is evolved : {cOHo = 0H° + CO- - add. Water. Carbonic oxide. 7. The most convenient method of obtaining carbonic oxide for lab- oratory purposes consists in heating potassic ferrocyanide with from eight to ten times its weight of concentrated sulphuric acid (Fownes). The flask containing the mixture must be gently heated in order to start the reaction, which afterwards continues of itself. The evolution of gas is apt to be somewhat violent. The reaction takes place accord- ing to the following equation: Fe"C6N6K4 + 60H2 + 6S02Ho2 = 6CO Potassic Water. Sulphuric Carbonic ferrocyanide. acid. oxide. + 2S02Ko2 + S02Feo" + 3S02(NH40)2. Potassic Ferrous Ammonic sulphate. sulphate. sulphate. The water necessary for the reaction is derived partly from the water of crystallization of the potassic ferrocyanide, which is an aquate of the formula Fe//C6N6K4,30H2, and partly from the commercial sulphuric acid, which never possesses the concentration corresponding to the pure dibasic acid S02Ho2. Formation.—When air enters a coal fire at the lower part of a grate or stove, the carbon combines with the oxygen of the air, forming car- 210 INORGANIC CHEMISTRY. bonic anhydride. The carbonic anhydride passes upwards through the glowing carbon, and is in this way (see Reaction 1, p. 208) reduced to carbonic oxide, which may frequently be seen burning with a peculiar bluish flame where it escapes into the air at the upper part of the fire. Sometimes this carbonic oxide passes off unburnt, involving great waste of fuel. The same formation of carbonic oxide occurs on a large scale in blast furnaces. Carbonic oxide is also formed in the destructive distillation of many organic substances containing oxygen. For this reason, it is a never- failing constituent of coal-gas. Properties.—Carbonic oxide is a colorless gas, devoid of taste, but possessing a faint odor. It is only very slightly soluble in water. Neither the gas nor its aqueous solution has any action on litmus. It is readily inflammable, and burns in air or oxygen with a pale blue flame, forming carbonic anhydride: CO + O = C02. Carbonic Carbonic oxide. anhydride. Mixed with half its volume of oxygen, as expressed in the above equa- tion, it explodes on the approach of a burning body. It is perfectly stable at all known temperatures. In its physiological action it displays the characteristics of a violent narcotic poison. Traces of it, if present in air, are sufficient to cause giddiness and headache when inhaled; in larger doses it produces in- sensibility, and even death. Small animals die quickly in an atmos- phere containing 1 per cent, of this gas. Its action seems to depend on the formation of a compound of carbonic oxide with the haemoglobin of the blood, by which the latter is prevented from exercising its function as an absorbent of oxygen. Carbouic-oxide-hsemoglobin pos- sesses a characteristic absorption spectrum, by means of which the presence of carbonic oxide in the blood, in cases of poisoning by this gas, may be recognized. Owing to the readiness with which carbonic oxide is formed, such cases of poisoning, both accidental and inten- tional, occur not infrequently when the products of combustion from stoves or braziers are allowed to escape into dwelling-rooms. Reactions.—The following reactions of carbonic oxide all depend upon its peculiar character as a compound containing dyad carbon. The carbon passes readily into its normal tetradic condition, and in this way carbonic oxide is enabled to form additive compounds. 1. At high temperatures carbonic oxide acts as a reducing agent, taking up oxygen and forming carbonic anhydride. Many of the oxides of the metals are reduced to the metallic state when heated in the gas, which in this way plays an important part in many metallur- gical operations. 2. At a temperature of 80° 0.(176° F.) carbonic oxide is readily ab- sorbed by potassium, forming a compound of the formula 3. Carbonic oxide and chlorine in equal volumes unite under the in- fluence of sunlight to form carbonic oxydichloride ox: phosgene gas: .NITROGEN. 211 CO + C12 = OOC12. Carbonic Carbonic oxide. oxydichloride. Carbonic oxydichloride has a suffocating odor. At lower temperatures it condenses to a colorless liquid, boiling at 8.2° C. 4. Carbonic oxide is readily absorbed by a solution of cuprous chloride in hydrochloric acid, or by solutions of cuprous salts in ammo- nia. The compound with cuprous chloride crystallizes in fatty scales possessing the formula CO(CuCI)2,2OH2. Composition.—The composition of carbonic oxide is most readily as- certained by exploding the gas with oxygen in a eudiometer. 100 c.c. of carbonic oxide and 100 c.c. of oxygen are introduced into the eudio- meter, making a total of 200 c.c. After the passage of the electric spark, it is found that the volume has been reduced to 150 c.c. Of these, 100 c.c. are absorbed by caustic potash, proving them to be carbonic anhydride. The remaining 50 c.c. are found to consist of pure oxygen. Therefore the carbonic oxide has yielded its own volume of carbonic anhydride, taking up half its vol- ume of oxygen in the process. But it has already been proved (p. 205) that carbonic anhydride contains its own volume of oxygen ; carbonic oxide therefore contains half its volume of oxygen. Expressing the volumes in litres: 1 litre of carbonic oxide weighs . . . J litre of oxygen weighs . . , 14 criths. . . . 8 “ U Difference, . . . 6 “ The difference is the weight of carbon. In carbonic oxide the pro- portion of carbon to oxygen is, therefore, as 6 : 8, or, in atomic weights, as 12 : 16, and the formula of this compound is therefore CO. The compounds of carbon with chlorine, nitrogen, and hydrogen, will be described under Organic Chemistry. CHAPTER XXVI. PENTAD ELEMENTS. NITROGEN, Azote, X 2. Section I. Atomic weight = 14. Molecular weight = 28. Molecular volume I I 1. 1 litre weighs 14 criths. Liquefiable by great pressure and cold. Atomicity T, which, by the mutual saturation of pairs of bonds, becomes reduced to or to ' (see p. 80). Evidence of atomicity: Nitrous oxide, . on2 Ammonia, . N'"HS. Ammonic chloride, . NvH4C1 Phosphoric fluoride (analogy), ■ P"F,. History.—Nitrogen was discovered by Rutherford in 1772. He found that when an animal was allowed to breathe the air confined 212 INORGANIC CHEMISTRY. under a bell-jar, and the impure air thus obtained was treated with a caustic alkali, a gas remained behind, incapable of supporting combus- tion or respiration. The name nitrogen signifies “ the nitre-producer” (from nitrum, nitre, and pww, I bring forth), and refers to the fact that this element is a constituent of nitre. Occurrence.—Nitrogen occurs in the free state in the atmosphere, of which it forms about four-fifths by volume. Recently its presence in the sun and in some nebulae has been rendered probable by spectrum analysis. In combination it is found in minute quantity as ammonia in the atmosphere, and it is also a constituent of numerous animal and vegetable substances. Preparation.—1. Nitrogen is most readily obtained from atmospheric air by the removal of the oxygen. For this purpose the combustion of phosphorus is usually employed. The phosphorus is placed in a small porcelain crucible, supported by a cork floating on water, and, after setting fire to the phosphorus, a bell-jar is placed over it. The phos- phorus burns, combining with the oxygen, and forming dense white clouds of phosphoric anhydride, which are speedily absorbed by the water. The nitrogen thus obtained is never quite pure, inasmuch as the phosphorus ceases to burn before the last traces of oxygen have been removed. It may be purified by leaving it in contact with moist phos- phorus, which by its slow oxidation completely removes the remaining oxygen. Moist alkaline sulphides, moist ferrous sulphide, and a number of other easily oxidizable substances, act in a similar manner in remov- ing oxygen from gaseous mixtures, 2. Very pure nitrogen may be obtained by passing a current of air, freed from carbonic anhydride and moisture, over metallic copper con- tained in a tube of hard glass, and heated to redness in a furnace. The oxygen of the air combines with the copper, forming cupric oxide, whilst the nitrogen passes on unchanged and may be collected. 3. On heating a concentrated solution of ammonic nitrite or a mixture of ammonic chloride and potassic or sodic nitrite, nitrogen is evolved ; N'"0(NvH4O) = N2 + 20H2. Ammonic nitrite. Water. NH401 + NONao = NaCl -f N2 + 20H2. Amraonic Sodic Sodic Water, chloride. nitrite. chloride. 4. Nitrogen is given off when ammonic dichromate, or a mixture of potassic dichromate with amnionic dichloride, is heated: f Cr02(NvHtO) < O = N2 + Cr203 40 H2. (Cro2(NvH4O) Ammonic chromate. Chromic oxide. Water. 5. When chlorine is passed through an excess of an aqueous solution of ammonia, the chlorine combines with the hydrogen of the ammonia, forming hydrochloric acid, which unites with the excess of ammonia, and nitrogen is liberated : COMPOUNDS OF NITROGEN. 213 8NH3 + 3C]2 = 6NH4CI + N2. Ammonia Ammonic chloride. The entrance of each bubble of chlorine into the solution is attended with a flash of light. Great care must be taken that the ammonia is always in excess, otherwise the very dangerously explosive compound, nitrous chloride, will be formed. Properties.—Nitrogen is a colorless, tasteless, and inodorous gas, slightly lighter than air. It is not capable of supporting either com- bustion or respiration. A lighted taper is extinguished, and small ani- mals die when plunged into this gas. It is not, however, poisonous, as is evident from the fact that it is contained in such large quantities in atmospheric air. Water dissolves only 0.025 of its bulk of the gas. Nitrogen is neither acid nor alkaline. It is one of the most indifferent bodies known, combining directly with only very few of the elements. COMPOUNDS OF NITROGEN WITH OXYGEN AND HYDROXYL. Nitrous oxide {hyponitrous anhydride), on2 N—O—N or, "N'20 II >0 NX Nitric oxide, * . ;N"0 —N=0 Nitrous anhydride, . f NO {° (NO 0 o tl 1! N—O—N Nitric peroxide, j no2 \no2 0=N=0 i 0=N=0 and, no2 0=N=0 1 Nitric anhydride, ....... p Ino2 o o II II N—O—N II If o o n Hyponitrous acid, / NHo \NHo N—0—H 11 N—0—H Nitrous acid, NOHo G=N—0—II Nitric acid, . N02Ho O 1! N—O—H il O 214 INORGANIC CHEMISTRY. The most important member of the above group, and the starting- point for the preparation of all the others, is nitric acid. This com- pound will be described first. NITRIC ACID, Aquafortis. N02Ho. Molecular weight = 63.' Fuses at —so° C. (—sB° F.). Boils at 86° C. (186.8° F.). History.—Nitric acid was known to the alchemists. Lavoisier showed that it contained oxygen, but its exact composition was first ascertained by Cavendish. Production.—1. When a series of electric sparks is passed between platinum points in a glass globe containing air, red fumes of nitric per- oxide (q.v.) are formed. On shaking the contents of the globe with water, the red fumes disappear, and the water acquires an acid reaction, arising from the presence of nitric acid in solution. It was in this way that the formation of nitric acid was studied by Cavendish. The production of red fumes is enormously increased by passing the sparks through compressed air. In like manner, nitric acid is formed when hydrogen is burned in oxygen containing a small proportion of nitrogen, or when an excess of the gases obtained by the electrolytic decomposition of water is mixed with air and exploded in a eudiometer. Nitric acid is also produced in the combustion of ammonia in oxygen. 2. When nitrogenous animal matter is slowly oxidized by the action of the air, at a temperature between 20° and 30° C. (68°-86° F.), in presence of water and powerful bases, nitric acid is formed, and com- bines with the bases to form nitrates. In this way the nitrites and nitrates which are found in the shallow well waters of towns have been formed from the nitrogenous matter contained in the soil. In hot climates, particularly in districts where there is little rain, the nitrates make their appearance as an efflorescence on the surface of the soil, as in India and in Chili. This natural formation of potassic nitrate is imitated artificially in the so-called nitre plantations. In these, animal matters mixed with lime and ashes, are placed in loose heaps, exposed to the air but shel- tered from rain. From time to time the heaps are watered with urine and stable runnings. The nitre bed is usually lixiviated every three years, and the product, consisting chiefly of calcic nitrate, is converted by treatment with potassic carbonate into potassic nitrate, which is purified by crystallization. In this way a cubic metre of earth may, under favorable conditions, be made to yield as much as 20 kilos, of nitre. Nitrification appears to depend upon the presence of an organized ferment. Manufacture.—Nitric acid is prepared by distilling potassic nitrate NITRIC ACID. 215 (nitre) or sodic nitrate (cubic nitre or Chili saltpetre) with concentrated sulphuric acid: N02Ko + S02Ho2 = S02HoKo + N02Ho Potassic Sulphuric Hydric potassic Nitric acid, nitrate. acid. sulphate. By employing two molecules of potassic nitrate to one of sulphuric acid, a saving of sulphuric acid is effected, but a higher temperature is required, which destroys some of the nitric acid. In this case the reac- tion takes place in two stages, of which the first is expressed in the above equation, whilst in the second, the hydric potassic sulphate acts upon another molecule of potassic nitrate: S02HoKo + N02Ko = S02Ko2 + N02Ho Hydric potassic Potassic Potassic Nitric acid sulphate. nitrate. sulphate. A further disadvantage of the second method lies in the fact that the normal potassic sulphate can be removed from the retort only in the solid state, whereas the hydric potassic sulphate, from its greater fusi- bility, can be poured out. On a commercial scale the distillation is performed in cast-iron cylin- ders A (Fig. 37) lined with fire-clay and heated over a furnace. The distillate is condensed in large stoneware WoulfPs bottles, B, each con- Fig. 37. nected with the one following. The last of these leads into a coke tower, down which a stream of water trickles. Any fumes of nitric peroxide which have escaped condensation in the Woulff’s bottles are absorbed by the water in the coke tower. Chili saltpetre is generally employed in the manufacture of nitric acid. It is cheaper than nitre, and, owing to the lower atomic weight of sodium, yields a larger proportion of nitric acid. The acid thus obtained may be purified by distillation with its own volume of concentrated sulphuric acid. The distillate contains from 99.5 to 99.8 per cent, of N02Ho. 216 INORGANIC CHEMISTRY. Properties.—Pure nitric acid is a colorless fuming liquid of sp. gr. 1.53. It has an irritating odor, and is powerfully corrosive, cauterizing the skin and staining it yellow. It begins to boil at 86° C. (187° F.), but is partially decomposed into nitric peroxide, oxygen, and water, so that gradually the distillate becomes weaker, and the boiling point rises, till at last an acid containing 68 per cent, of N02Ho, and boiling at 120.5° C. (248.9° F.), distils over under ordinary pressure without further change. This acid has a sp. gr. of 1.414 at 15° C. (59° F.), and is the ordinary concentrated nitric acid of commerce. If a weaker acid be distilled, the liquid in the retort becomes gradually more con- centrated, till the acid containing 68 per cent, is obtained, which then distils unchanged. Notwithstanding the constancy of its boiling point, this acid is not a definite compound. By varying the pressure under which the distillation is performed, acids of varying strength may be obtained, but for each of these pressures, there is a fixed strength of acid with a constant boiling point. Under a pressure of 70 mm. an acid containing only 66.7 per cent, of N02Ho distils over between 65° and 70° C. (149°-158° F.). The higher pressure thus corresponds to the greater strength of acid, the reverse being the case with hydrochloric acid (see p. 158). When concentrated nitric acid is mixed with water, diminution of volume and elevation of temperature ensue. The following table con- tains the specific gravities of various strengths of aqueous acid at 0° and 15° C. (32°-59° F.), as determined by J. Kolb: Per cent. NO2H0. Bp. gr. at 0° C. (32° F.). Bp. gr. at 15° C. (59° F.). 100.00 1.559 1.530 90.0 1.522 1.495 80.0 1.484 1.460 70.0 1.444 1.423 60.0 1.393 1.374 50.0 1.334 1.317 40.0 1.267 1.251 30.0 1.200 1.185 20.0 1.132 1.120 15.0 1.099 1.089 10.0 1.070 1.060 5.0 1.031 1.029 The decomposition which concentrated nitric acid undergoes under the influence of heat is expressed by the following equation: 4N02Ho = 20H2 + 2'Niv204 +, 02. Nitric acid. Water. Nitric peroxide. This decomposition is very rapid at 100° C., and on this property the powerful oxidizing action of hot nitric acid depends. Concentrated nitric acid, when exposed to the action of light, turns yellow, owing to a decomposition similar to the above. Reactions.—1. With metallic oxides or hydrates nitric acid yields nitrates: NITRIC ACID. 217 OKH + N02Ho == N02Ko + OH2. Potassic Nitric acid. Potassic Water, hydrate. nitrate. PbO + 2N02Ho = + 0H2. Plumbic Nitric acid. Plumbic Water, oxide. nitrate. 2. The action of nitric acid upon metals is of a somewhat complicated character, varying not only with different metals, but also, for the same metal, with the strength of the acid employed and the temperature at which the reaction takes place. Nitrates of the metals are formed, but at the same time another portion of the nitric acid is reduced to some lower oxide of nitrogen. Thus, silver, copper, and mercury, are attacked by nitric acid acid with formation of nitrates and evolution of nitric oxide: (NO, 3Cu + 8N02Ho = S< Cuo" + 2'N"Q -f 40H2. (NO, Nitric acid. Cupric nitrate. Nitric oxide. Water. With very concentrated acid, nitric peroxide ('Niv2o4) is generally evolved, and when the reaction takes place at a high temperature, a portion of the nitric acid is completely reduced to nitrogen. When silver is slowly dissolved by weak nitric acid in the cold, nitrous acid is formed. When nitric acid acts upon copper in presence of much cupric nitrate, the gas evolved consists chiefly of nitrous oxide. When nitric acid acts upon a more electro-positive metal, such as zinc, nitrous oxide is evolved, and when a very concentrated acid is employed, ammonia is formed and combines with the excess of nitric acid: f no2 4Zn + 10NO2110 = 0N2 -j- 4 < Zno" -j- 50H2. Uo2 Nitric acid. Nitrous oxide. Zincic nitrate. Water. (NO, 4Zn + 9N02Ho = 4< Zno" + 30H2 + NH3. I NO, Nitric acid. Zincic nitrate. Water. Ammonia. By the action of zinc in an alkaline solution, the whole of the nitric acid present is reduced to ammonia by the nascent hydrogen. The ammonia may be distilled off and absorbed in a solution of hydrochloric acid. This method is employed in the quantitative estimation of nitric acid. 3. The general action of nitric acid is that of a powerful oxidizing agent. Sulphur, phosphorus, carbon, amorphous boron and silicon, 218 INORGANIC CHEMISTRY. arsenic, and iodine, are converted by treatment with nitric acid into sulphuric, phosphoric, carbonic, boric, silicic, arsenic, and iodic acids. In the case of phosphorus the oxidation takes place with explosive vio- lence, and if the concentrated acid be dropped upon hot sawdust or finely powdered charcoal, the latter inflames. It has been mentioned under the heading of hydrochloric acid that oxidizing agents liberate chlorine from this acid. In this way chlorine is evolved from a mixture of nitric and hydrochloric acids: N02Ho + 3HCI = NOCI + 20H2 + Cl2. Nitric acid. Hydrochloric Nitrons Water, acid. oxychloride. (Nitrosylic chloride.) This mixture was known to the alchemists, who gave to it the name aqua-regia, from its power of dissolving gold, the king of metals. It is employed in the laboratory as a solvent for gold, platinum, and various ores. The solvent action depends on the presence of the chlorine evolved in the above reaction. The action of nitric acid on organic compounds will be studied in connection with these (Organic Chemistry). Nitrates.—Nitric acid is generally monobasic. The numerous so- called basic nitrates may, however, be regarded as salts of tribasic and pentabasic nitric acid (NOHo3 and NHo5). Graham first pointed out that in basic salts the base frequently replaces the water of crystalliza- tion of the normal salt. This supposed water of crystallization must, therefore, in as far as it may be replaced by a base, be regarded as water of constitution. Thus cupric nitrate (^Q2Cu°r/,30H2) and basic cupric nitrate 2Cu0,0112j might be formulated j^q^2 Cuo",OH2 and //Cuo//,0112. The monobasic nitrates are ail soluble in water. At a high temperature the nitrates are all decomposed. They gen- erally evolve, first, pure oxygen, then nitric peroxide, or a mixture of nitrogen and oxygen, whilst an oxide of the metal is left. The presence of nitrates in solution may be recognized by the follow- ing characteristic reaction : The solution supposed to contain a nitrate is mixed in a test-tube with a solution of ferrous sulphate. Concen- trated sulphuric acid is then poured down the side of the sloping tube, so as to sink to the bottom of the liquid without mixing with it. If a nitrate is present, a characteristic brown coloration will be visible at the surface of contact of the two layers. The explanation of this is that the nitric acid, liberated by the sulphuric acid, is reduced by the fer- rous sulphate to nitric oxide, the latter dissolving in the excess of fer- rous sulphate with a brown color. NITRIC ANHYDRIDE. 219 NITRIC ANHYDRIDE. N2O, Probable molecular weight = 108. Fuses at 29.5° C. Boils at 45° C. History.—Nitric anhydride was discovered by Deville in 1849. Preparation.—l. Thiscomponnd isformed when dry chlorine is passed over dry argentic nitrate contained in a U-tube and heated in a water- bath. The reaction takes place in two stages. In the first of these nitric dioxychloride, a volatile liquid, is formed : N02Ago + Cl2 N02CI + AgCl '+ O. Argentic Nitric Argentic nitrate. dioxychloride. chloride. In the second the nitric dioxychloride acts on the unattached argentic nitrate: N02Ago + N02CI = N205 + AgCl. Argentic Nitric Nitric Argentic nitrate. dioxychloride. anhydride. chloride. The reaction begins at 95° C. (203° F.), and, when once started, con- tinues, even when the temperature is allowed to fall as low as 60° C. (140° F.). All unnecessary heating must be avoided, as the anhydride is totally decomposed at a temperature very slightly above that required for its formation. The anhydride distils over, and is condensed in a tube surrounded by a freezing mixture. 2. Nitric anhydride may also be obtained by abstracting the elements of water from nitric acid by means of phosphoric anhydride: 2N02Ho = N2Os + 0H2. Nitric acid. Nitric Water, anhydride. The phosphoric anhydride is added very gradually to the concentrated nitric acid, cooled by ice, and the pasty mass is afterwards distilled at a low temperature. The anhydride collects as a crystalline mass in the receiver. Properties.—Nitric anhydride forms large colorless prisms, which fuse at 29.5° C. (85.1° F.) It boils with decomposition and evolution of brown fumes about 45° C. (113° F.). When sealed in a glass tube, it may be preserved unaltered, if kept in a cool place; but, in a warm room, gradually undergoes decomposition into oxygen and nitric perox- ide, ultimately fracturing the tube with the internal pressure. When thrown into water the anhydride hisses violently, evolving great heat, and combining with the water to form nitric acid; N2G5 + 0H2 = 2N02Ho. Nitric Water. Nitric acid, anhydride- 220 INORGANIC CHEMISTRY. . NITROUS OXIDE, Hyponitrous Anhydride, Laughing Gas. ON2. Molecular weight = 44. ' Molecular volume LLJ. 1 litre weighs 22 criths. Fuses at —lol° C. (—149.8° F.). Boils at —BB° C. (—126.4° F.). History.—This compound was discovered by Priestley in 1772. Preparation.—l. Nitrous oxide is formed by the action of dilute nitric acid upon zinc : f no2 10N0.,H0 + 4Zn = ON, -f 4/ Zno" + 50H2. u°2 Nitric acid. Nitrous oxide. Zincic nitrate. Water. This method does not, however, yield the compound in a state of purity, and is never employed in its preparation. 2. Nitrous oxide may readily be obtained in large quantity by heat- ing ammonic nitrate. Under the influence of heat, the elements of water are removed from this salt and nitrous oxide is formed : N02(NtH40) = 20H2 + 0N2 Ammonic nitrate. Water. Nitrous oxide. The ammonic nitrate, previously dried, is heated in a flask to which a delivery tube is attached. The heat must not be applied too suddenly, otherwise the decomposition takes place with explosive violence, and nitric oxide is formed. The gas is purified by passing it first through a solution of ferrous sulphate, in order to absorb nitric oxide, and then through caustic potash, to free it from chlorine derived from ammonic chloride contained in the commercial nitrate. It may be collected over mercury, or over warm water, in which it is less soluble than in cold wrater. Properties.—Nitrous oxide is a colorless gas with a faint pleasant odor, and a sweetish taste. Its density is 1.527 (air = 1). Water dis- solves about four-fifths of its volume of the gas, and alcohol takes up a still larger quantity. Nitrous oxide supports the combustion of bodies which burn in oxy- gen. A glowing match is rekindled when plunged into the gas, and burns almost as brightly as in oxygen. Phosphorus burns with a flame of dazzling brightness. Feebly burning sulphur is extinguished by the gas, but, if burning strongly, the combustion continues with great vigor. All combustions in nitrous oxide are effected solely at the expense of the oxygen contained in the gas, the nitrogen taking no part in the re- action. In order that combustion may continue, it is necessary that the temperature of the burning body should be sufficiently high to decom- pose the nitrous oxide into nitrogen and oxygen. If this condition is HYPONITEOTJS ACID. 221 not fulfilled, combustion is impossible, as may be seen* in the case of feebly burning sulphur. Strictly speaking, therefore, nitrous oxide, as such, does not support combustion. It does so only by the agency of one of its products of decomposition—oxygen. Nitrous oxide was first liquefied by Faraday, by heating amnionic nitrate in a bent tube (see p. 155). It may be most conveniently lique- fied with the aid of a force-pump, cooling the wrought-iron receiver with ice. Liquid nitrous oxide is colorless, and very mobile. It boils at —BB° C. (—126.4° F.) under atmospheric pressure, whilst at 0° C. the tension of its vapor is 30 atmospheres. By means of the cold pro- duced by its own evaporation, or by plunging a tube containing it into a bath of solid carbonic anhydride in ether, and allowing this freezing mixture to evaporate in vacuo, liquid nitrous oxide may be frozen into colorless crystals resembling in appearance amnionic nitrate. By the evaporation in vacuo of a mixture of liquid nitrous oxide and carbonic disulphide, a degree of cold equal to —l4o° C. (—22o° F.) may be obtained. Liquid nitrous oxide, in spite of its low boiling point, may be preserved in open glass tubes for over half an hour. If mercury be poured into this liquid, the metal is instantly frozen. Nitrous oxide, when inhaled, acts as a narcotic poison. In smaller doses it produces temporary nervous exhilaration or intoxication ; hence the name laughing gas. It is employed in minor surgical operations as an anaesthetic. Composition.—The composition of nitrous oxide may be ascertained by heating sodium in a bent glass tube containing a measured volume of the gas over mercury (see p. 159). The sodium combines with the oxygen of the gas, forming solid sodic oxide, and liberating the nitrogen. After the action is finished, the gas remaining in the tube is found to possess exactly the same volume as the gas employed, and may be shown •to consist of pure nitrogen. Hence nitrous oxide contains its own vol- ume of nitrogen. Expressing the volumes in litres— 1 Jitre of nitrous oxide weighs . . . Deduct weight of litre of nitrogen, . . 14 “ There remain . . 8 “ which is the weight of J litre of oxygen. One litre of nitrous oxide therefore contains 1 litre of nitrogen and J litre of oxygen; or, 2 volumes of nitrogen combine with 1 volume of oxygen to form 2 vol- umes of nitrous oxide. Expressed in atomic weights, 28 parts by weight of nitrogen combine with 16 of oxygen to form 44 of nitrous oxide. HYPONITROUS ACID. //JNHo t NHo' Knoim only in its salts, or in aqueous solution. Preparation of Argentic Hyponitrite (//N/2Ago2).—When an aqueous solution of po- tassic nitrate is treated with sodium amalgam in the proportion of four atoms of sodium to one molecule of nitrate, a reduction of the nitrate takes place according to the fol- lowing equation; 222 INORGANIC CHEMISTRY. ■2N02Ko + 4H2 = //N/2Ko2 + 40 U2. Potassic Potassic Water, nitrate. hyponitrite. Potassic nitrite is formed as an intermediate product in this reaction, and a saving of sodium amalgam may be effected by starting from the nitrite : 2NOKo + 2H2 = //N/2K02 + 20H2. Potassic Potassic Water, nitrite. hyponitrite. The alkaline liquid obtained by either of these processes is then accurately neutral- ized with acetic acid, and argentic nitrate is_ added. Argentic hyponitrite is thus ob- tained as a greenish-yellow precipitate, which, by solution in dilute nitric acid and precipitation with ammonia, acquires a pure yellow color. Properties.—Argentic hyponitrite may be dissolved in weak acids without suffering immediate decomposition, but the solution is very unstable. A solution of potassic hyponitrite acidulated with acetic acid undergoes decomposi- tion on heating, the liberated hyponitrous acid breaking up into nitrous oxide and water: Hence nitrous oxide may be considered as the anhydride of hyponitrous acid. The acid salts of hyponitrous acid are known only in solution. Thus baric hypo- "H'2Ho# = "N',o + OH,. nitrite, " | //, which is insoluble in water, dissolves in aqueous hyponitrous acid with formation of an acid salt. The existence of this salt proves that hyponitrous acid must be at least dibasic. NITROUS ANHYDRIDE NA- Probable molecular weight = 76. Preparation.—1. When nitric acid is heated along with bodies ca-* pable of taking up oxygen, such as arsenious acid or starch, nitrous anhydride is formed; AsA + 2N02Ho As A + NA + 0H2. Arsenious Nitric acid. Arsenic Nitrous Water, anhydride. anhydride. anhydride. The nitrous anhydride thus obtained is mixed with nitric peroxide. 2 Nitrous anhydride may also be prepared by mixing 4 volumes of nitric oxide with 1 volume of oxygen. Direct combination takes place according to the equation : 2'N//0 + O = N2Q3. Nitric oxide. Nitrous anhydride. Properties.—Nitrous anhydride prepared by either of the above reac- tions is a reddish gas, which by passing through a U-tube immersed in a freezing mixture, may be condensed to a blue liquid. It is a very unstable compound, and undergoes gradual decomposition, even below 0° C., into nitric oxide and nitric peroxide : 223 NITROUS ACID N2Os = 'N"O + 'Niv02 Nitrous Nitric Nitric anhydride. oxide. peroxide. On warming, this decomposition is very rapid. The addition of a small quantity of water to nitrons anhydride con- verts it into nitrous acid : N203 + 0H2 = 2NOHo. Nitrous Water. Nitrous anhydride. acid. A larger quantity of water decomposes the compound with efferves- cence : nitric oxide is evolved, and nitric acid remains in solution; 3N203 + OH2 = 2N02Ho + 4'N"O. Nitrous Water. Nitric acid. Nitric oxide, anhydride. The two foregoing reactions illustrate strikingly the inadequacy of chemical equations as expressions of chemical change. In the first equation, the proportion of water to nitrous anhydride is three times as great as in the second; yet the first stands for a reaction in which only a small quantity of water is required, and the second for a reaction which occurs only in presence of an excess of water. The reason of this discrepancy is that ordinary equations take no account of the rela- tive masses of the reacting substances, and the mass of a substance is frequently an important factor, determining in some cases the direction of the chemical change. NITROUS ACID. NOHo. Molecular weight 47 Preparation.—Nitrous acid may be obtained by mixing liquefied nitrous anhydride with water as above described. It cannot be pre- pared in a state of purity, and is an exceedingly unstable compound. Decompositions.—1. In the presence of much water nitric acid and nitric oxide are formed : 3NOHo = N02Ho + 2'N"O + 0H2. Nitrous acid. Nitric acid. Nitric oxide. Water. 2. Under some circumstances nitrous acid acts as a reducing agent: 2NOHo + 02 = 2N02Ho. Nitrous acid. Nitric acid. In this way acidulated solutions of the nitrites decolorize potassic per- manganate, reduce soluble chromates to green chromic salts, and precipi- 224 INORGANIC CHEMISTRY. tate gold and mercury in the metallic state from solutions of their salts. 3. In many other cases nitrous acid displays oxidizing properties: 4NOHo = 4'N"O + 20H2 + 02. Nitrous acid. Nitric oxide. Water. Thus acid solutions of the nitrites liberate iodine from potassic iodide and bleach a solution of indigo. Nitrites.—With metallic oxides or hydrates, nitrous acid forms ni- trites : OKH + NOHo = NOKo + OH2. Potassic Nitrous Potassic Water, hydrate. acid. nitrite. The alkaline nitrites may be most readily obtained by cautiously heating the nitrates. An addition of copper or lead facilitates the re- action by aiding in the removal of the oxygen : 2N02Ko = 2NOKo + 02. Potassic Potassic nitrate. nitrite. The temperature must not be raised too high, otherwise the nitrite will be decomposed. The alkaline nitrites are soluble in alcohol, and may thus be separated from unaltered nitrate, which is insoluble. The nitrites evolve reddish vapors when treated with dilute acids, and may thus be distinguished from the nitrates, which do not possess this property. NITRIC OXIDE. 'N"O. Molecular weight = 30. Molecular volume I I I. 1 litre weighs 15 criths. Liquefiable by great pressure and cold. History.—Nitric oxide was discovered by Van Helraont, who, how- ever, failed to recognize its true character. It was first investigated by Priestley. Preparation.—l. Nitric oxide is formed when nitric acid acts upon mercury or copper: fNO2 3Cu + 8N02Ho = 3< Cuo" + 2'N//0 + 40H2. I NO, Nitric acid. Cupric nitrate. Nitric oxide. Water. The gas is purified by passing it through a solution of caustic soda. Nitric oxide thus prepared is apt to contain nitrous oxide and free nitrogen, particularly towards the end of the reaction. In order to purify the product, advantage is taken of the property which nitric 225 NITRIC OXIDE. oxide possesses of dissolving in a concentrated solution of ferrous sul- phate. The solution of this salt absorbs the gas in large quantity, forming a compound of the formula 2S02Feo'/,/N//G, which remains dissolved in the liquid, imparting to it a deep brown color. On heating this brown liquid, pure nitric oxide is evolved. 2. Nitric oxide may be readily obtained in a state of purity by acting upon nitric acid with ferrous sulphate. A convenient mode of apply- ing this reaction consists in introducing into a retort 30 grams of nitre with 240 grams of ferrous sulphate, and pouring in through a funnel 250 cubic centimetres of a mixture of sulphuric acid with three times its bulk of water: fso2_, 6S02Feo" + 2N02Ko + 5S02Ho2 = 2'N"O + S02—('Fe'"206)Ti lso2 1 Ferrous Potassic Sulphuric Nitric Ferric sulphate, sulphate. nitrate. acid. oxide. + 2S02HoKo + 40H2. Hydric potassic Water, sulphate. Properties.—Nitric oxide is a colorless gas of density 1.039 (air = 1). Water dissolves one-twentieth of its volume of the gas. Neither the gas nor its aqueous solution exerts any action upon litmus. The molecular formula NO, deduced from the vapor-density of this compound, is anomalous. This formula involves the assumption that the molecule contains an odd number of unsatisfied bonds (see Note, p. 179). Although nitric oxide contains, for the same volume of nitrogen, twice as much oxygen as nitrous oxide, it does not support combustion so readily, owing to its greater stability. Feebly ignited charcoal is extinguished when plunged into the gas, whereas strongly glowing charcoal burns in it with great brilliancy. Phosphorus may be melted in the gas without igniting, and the flame of feebly burning phosphorus is extinguished by it; but phosphorus already well ignited continues to burn in it, emitting an intense light. Sulphur, even when burning strongly, is extinguished by nitric oxide. A mixture of nitric oxide and the vapor of carbonic disulphide burns with a vivid blue flame, very rich in chemically active rays. Reactions.—l. When nitric oxide and oxygen are mixed, a reddish gas is formed, consisting of nitrous anhydride and nitric peroxide, both of which compounds are produced by the direct union of the nitric oxide with the oxygen: 4'N"O + 02 = 2N'"203. Nitric oxide. Nitrous anhydride. 2'N"O + 02 = 'Niv204. Nitric oxide. Nitric peroxide. These gases are absorbed by water, to which they impart an acid reaction. 226 INORGANIC CHEMISTRY. 2. Nitric oxide also combines directly with chlorine to form nitrous oxychloride (q.v.): 2/N"0 + Cl2 = 2NOCI. Nitric oxide. Nitrous oxychloride. (Nitrosylic chloride.) The direct union which occurs in the above cases is probably de- pendent on the presence of a free bond in the nitrogen atom of nitric oxide, and the reactions consist in the saturation of this free bond by some suitable element. Composition.—The composition of nitric oxide may be determined in the same manner as that of nitrous oxide (see p. 221), but potassium must be employed, as sodium merely melts in the gas without decom- posing it. After the reaction is finished, it is found that the original volume has decreased by one-half, and that the residual gas is pure nitrogen. 1 litre of nitric oxide weighs .... . 15 criths. Deduct weight of J litre of nitrogen, . . . 7 “ There remain . . 8 “ which is the weight of J litre of oxygen. One litre of nitric oxide con- tains therefore J litre of nitrogen and litre of oxygen; or 1 volume of nitrogen combines with 1 vol. of oxygen to form 2 vols. of nitric oxide. Expressed in atomic weights, 14 parts by weight of nitrogen combine with 16 of oxygen to form 30 of nitric oxide. NITRIC PEROXIDE. ;Niv02 at higher temperatures. | j^Q2, or 'N~A, at tower temperatures. Molecular weight == 46 and 92. Molecular volume I I I. 1 litre weighs 23 to 46 criths. Fuses at —9° C. (15.8° F.). Boils at 22° C. (71.6° F.). Preparation.—l. Nitric peroxide may be obtained by the union of 2 volumes of nitric oxide with lof oxygen (see preceding page). The red gas thus formed may be condensed in a U-tube immersed in a freezing mixture. 2. Nitric peroxide is most conveniently prepared by the action of nitric acid on arsenious anhydride: As203 + 4N02Ho == AsA + 2'NiyA + 20H2. Arsenious Nitric acid. Arsenic Nitric Water, anhydride. anhydride. peroxide. NITRIC PEROXIDE. 227 Small fragments of arsenious anhydride are introduced into a retort with sufficient nitric acid of sp. gr. 1.393 to cover them. The reaction takes place on gently heating, and a mixture of nitric peroxide and nitrous anhydride condenses in the receiver, which is cooled by a freezing mixture. By passing a slow current of oxygen through this mixed pro- duct, the whole of the nitrous anhydride is converted into peroxide. 3. Certain nitrates, when subjected to destructive distillation, are decomposed into nitric peroxide, oxygen, and an oxide of the metal. Plumbic nitrate is well suited for this purpose: fNQ2 1 Pbo" = + PbO + 'NivA + O. I no. Plumbic nitrate. Plumbic oxide. Nitric peroxide. The thoroughly dried plumbic nitrate is heated in a retort connected with a TJ-tube which is drawn out at its further extremity to a tine opening and surrounded by a freezing mixture. The liquefied nitric peroxide collects in the tube, whilst the oxygen escapes through the fine opening. 4. It is also formed by the action of nitric acid on tin: Sn5 + 20NO2Ho = Sn.O5Ho10 + 50H2 + 10'N%04. Nitric acid. Metastannic acid. Water. Nitric peroxide. 5. Nitric peroxide is also formed by the action of nitric dioxychloride on argentic nitrite: NOAgo + = + AgCI. Argentic Nitric dioxychloride. Nitric Argentic nitrite. (Nitroxylic chloride.) peroxide. chloride. Properties.—Nitric peroxide is a volatile liquid which solidifies at —9° C. (15.8° F.), forming a white fibrous crystalline mass. Nitric peroxide displays remarkable changes of color, dependent upon the temperature. Just above its fusing point it is a colorless liquid. At 0° 0. it assumes a yellow tint, which deepens through orange to brown as the temperature rises to 22° C. (71.6° F.), when the nitric peroxide enters into ebullition, yielding a reddish-brown vapor. This vapor also assumes a darker color as its temperature is raised, becoming at last almost black. The vapor of nitric peroxide possesses a characteristic absorption spectrum. These changes of color correspond to definite changes of molecular condition, as may be seen from a study of the vapor-density of nitric peroxide at different temperatures. At a temperature very little above its boiling point it possesses a vapor-density below that required for the formula 'N1V204, but nearer to this value than to that required for 'NivOa. As the temperature rises the vapor-density diminishes, till at 140° 0. it corresponds exactly with the latter formula. There is, there- 228 INORGANIC CHEMISTRY. fore, even at the boiling point of nitric peroxide, a partial dissociation of the larger molecules, 'Nlv204, into the smaller, 'N1T02; but the greater number of the former still remain intact. The decrease in vapor- density corresponds with an increase in the relative number of disso- ciated molecules. It is probable that this dissociation begins even in the liquid state, as denoted by the change of color (see Note, p. 179). Liquid nitric peroxide is a powerfully corrosive substance, and its vapor is very irritating when inhaled even in small quantity. Reactions.—l. With metallic hydrates and oxides it yields a mix- ture of nitrite and nitrate in equivalent proportions: 'N%O4 + 20KH = N02Ko + NOKo + 0H2. Nitric Potassic Potassic Potassic Water, peroxide. hydrate. nitrate. nitrite. It thus behaves like a compound anhydride—a view of its chemical character which is supported by its formation from nitric dioxychloride and argentic nitrite (see above). 2. A small quantity of water acts like a metallic hydrate, producing a mixture of nitrous and nitric acids: 'NivA + OH2 = N02Ho + NOHo. Nitric peroxide. Water. Nitric acid. Nitrous acid. But an excess of water decomposes it into nitric oxide and nitric acid: 3/NiT204 + 20H2 = 4N02Ho + 2/N"Q. Nitric peroxide. Water. Nitric acid. Nitric oxide. Composition.—The composition of nitric peroxide may be ascertained by passing the vapor of a known weight of the gas over red-hot me- tallic copper. The oxygen of the peroxide combines with the copper, and may be determined by ascertaining the increase in weight of the latter. The nitrogen is liberated, always mixed however with a small quantity of nitric oxide, and may be collected and measured. The proportion of nitric oxide must also be determined. From these data the composition of the peroxide may be calculated. COMPOUNDS CONTAINING NITROGEN, CHLORINE, AND OXYGEN. NITROUS OXYCHLORIDE, Nitrosylic Chloride, Ghloronitrous Gas. Molecular weight 65.5, Molecular volume | | |. 1 litre weighs 32.75 criths. Boils at 0° C. NOCI. Preparation.—1. By the direct union of chlorine and nitric oxide: 2/N//0 + Cl2 = 2NOCI. Nitric Chlorine. Nitrons oxide. oxychloride. (Nitrosylic chloride.) NITRIC DIOXYCHLORIDE. 229 2. It is also evolved along with chlorine from a mixture of nitric and hydrochloric acids (see Aqua-regia, p. 218); N02Ho + 3HCI = NOCI + 20 H2 + Cl2 Nitric Hydrochloric Nitrous Water, acid. ' acid. oxychloride. Properties.—Nitrous oxychloride is an orange-colored gas, which, in a freezing- mixture, condenses to a red fuming liquid possessing an odor of aqua-regia. Reactions.—1. Nitrous oxychloride is decomposed by water into nitrous and hydro- chloric acids: NOCI + OH2 = NOHo + HCI. Nitrous Water. Nitrous Hydrochloric oxychloride. acid. acid. In like manner it yields, with metallic oxides and hydrates, a mixture of nitrite and chloride: NOCI + 20KH = NOKo + KCI + OH2. Nitrous Potassic Potassic Potassic Water, oxychloride. hydrate. nitrite. chloride. Nitrous oxychloride belongs to the class of chlorides of the acid radicals, a view regarding its constitution which is expressed by the name nitrosylic chloride. These chlorides are derived from the corresponding acids by the substitution of chlorine for hydroxyl. Water decomposes them into the corresponding acid and hydrochloric acid, as in the foregoing reaction. 2. Nitrous oxychloride attacks mercury. The chlorine combines with the metal to form mercurous chloride, whilst nitric oxide is liberated : 2NOCI + Hg2 = /Hg/2C12 + 2/N//0. Nitrous Mercurous Nitric oxychloride. chloride. oxide. It is without action on gold or platinum. The corresponding bromine compound, NOBr, has also been prepared. NITRIC DIOXYCHLORIDE, Nitroxylie Chloride, Chloropernitric Gas. N02CI. Molecular weight 81.5, Molecular volume I I I. 1 litre weighs 40.75 criths. Boils at 5° C. (41° F.). Preparation.—l. By passing nitric peroxide and chlorine together through a heated glass tube: 'N‘»A + Cl2 = 2N02CI. Nitric peroxide. Nitric dioxychloride. (Nitroxylic chloride.) 2. By the action of chlorine on argentic nitrate as already described (see Nitric Anhydride, p. 219). 3. By the action of sulphuric dioxychlorhydrate (sulphurylic chlorhydrate) on nitric acid: S02CIHo + N02Ho = S02Ho2 + N0SCI. Sulphuric Nitric Sulphuric Nitric dioxychlorhydrate. acid. acid. dioxychloride. 4. It is most readily obtained by heating plumbic nitrate with phosphoric oxytri- chloride: f N02 3 \ Pbo// + 2POCI3 = P202Pbo//s + 6N02CI. Ino2 Plumbic Phosphoric Triplumbic Nitric nitrate. oxy trichloride. diphosphate. dioxychloride. The action of the chlorine compounds of phosphorus on acids and their salts is a general method for the preparation of the chlorides of the acid radicals. 230 INORGANIC CHEMISTRY. Properties.—Nitric oxychloride is a heavy yellow oil boiling at 5° C. (41° F.). Reaction.—Water decomposes it into nitric and hydrochloric acids : N02CI + OH2 = N02Ho + HCL Nitric Water. Nitric Hydrochloric dioxychloride. acid. acid. Bases effect a similar decomposition, yielding a mixture of nitrate and chloride. COMPOUNDS OF NITROGEN WITH HYDROGEN AND HYDROXYL. AMMONIA. nh3. Molecular weight 17. Molecular volume 1 litre weighs 8.5 criths. Fuses at —7 5° C. (—lo3° F.). Boils at— 38.5° C.(—37.3° F.). History.—The aqueous solution of ammonia was known to the alche- mists. The gas was first obtained by Priestley, who also observed its decomposition by the electric spark. Berthollet first ascertained its com- position. Occurrence.—Ammonia occurs in small quantity in the air as carbo- nate, and in rain-water, especially in that which falls during thunder- storms, as nitrite and nitrate. Most fertile soils contain ammonia. As chloride and sulphate it is found in the neighborhood of active volca- noes. Along with boric acid, it occurs, as salts of ammonia, in the lagoons of Tuscany (p. 191), having probably been formed by the ac- tion of subterranean steam upon boric nitride : BN"' + 30 H2 = BHo3 + NH3. Boric nitride. Water. Boric acid. Ammonia. It also occurs, in the form of its salts, in animal fluids, particularly in putrid urine, and in the juices of plants. Formation.—Ammonia is formed: 1. By the decay of animal and vegetable matters containing nitrogen. It is from this source that the atmospheric ammonia is derived. 2. By the destructive distillation of these nitrogenous matters. The ammonia of commerce is thus obtained. Formerly, horn, hoofs, and bones were distilled for this purpose, and hence the name spirits of hartshorn was given to ammonia; but its chief source at the present day is the ammoniacal liquor of gas works, in which it occurs as a by- product from the distillation of coal. Volcanic ammonia is also a product of the destructive distillation of nitrogenous vegetable matter, being formed only where the lava has flowed over fertile soil. 3. By the action of nascent hydrogen (from zinc and caustic alkali) on nitric and nitrous acids. 4. Ammonia is also formed synthetically from its elements when the silent electric discharge is passed through a mixture of nitrogen and hy- drogen (Donkin). Preparation.—Ammonia may be prepared from any of its salts by AMMONIA. 231 heating these with slaked lime. The chloride is usually employed for this purpose: 2NH4CI + CaHo2 = CaCl2 + 2NH3 + 20H2. Ammonic Calcic Calcic Ammonia, Water. chWride. hydrate. chloride. One part of ammonic chloride is mixed with 2 parts of slaked lime in powder, and the whole is heated in a flask. If gaseous ammonia is re- quired, the gas evolved may be dried by passing over quicklime (calcic chloride absorbs gaseous ammonia), and may be collected either over mercury or by upward displacement. When an aqueous solution is re- quired the gas is passed direct into water, which is contained in a series of Woulff’s bottles fitte'd with safety-tubes. The delivery tubes must pass to the bottom of the liquid, otherwise only the upper layer would be saturated, as the aqueous solution of ammonia is lighter than water. Properties.—Ammonia is a colorless gas, with a very pungent odor. Its density is 0.589 (air = 1). It turns red litmus blue, and yellow turmeric paper brown. It neutralizes acids, uniting directly with them to form salts (see Reactions). Ammonia may be liquefied by cold or pressure. Faraday first ob- tained it in the liquid state by heating argentic ammonio-chloride in one limb of a bent sealed tube, whilst the other was immersed in a freezing mixture. The argentic ammonio-chloride is prepared by pass- ing ammonia over dry argentic chloride, which in this way absorbs 320 Fig. 38. times its volume of the gas. The double compound parts with all its ammonia when heated to 112° C. (233.6° F,). By conducting the heating in a bent sealed tube as above described, the ammonia is lique- fied by the joint action of its own pressure, and of the cold of the freez- ing mixture. Calcic ammonio-chloride may be substituted for the argentic compound in the above experiment. Ammonia may also be liquefied by the action of cold alone at a temperature of —4o° to —so° C. (—4o° to —s7° F.), by passing the gas through a tube immersed in a mixture of ice and crystallized calcic chloride. Liquid ammonia is a mobile, colorless, highly refracting liquid, boil- ing at —38.5° C. (—37.3° F). At —lo° C. (14° F.) it has a sp. gr. of 0.65. When subjected to a temperature below —7s° C. (—103° F.) it solidifies to a white crystalline translucent mass. 232 INORGANIC CHEMISTRY. The cold produced by the rapid evaporation of liquid ammonia has been utilized in Carry’s apparatus for the artificial production of ice. Two strong wrought-iron vessels, A and B (Fig. 38), are connected by a tube of the same material. A contains an aqueous solution of ammonia saturated at 0° C. When ice is to prepared by means of this apparatus, heat is applied to A, whilst B is immersed in cold water. Gaseous ammonia is evolved from A, and condenses under its own pressure be- tween the double walls of the receiver B. When a sufficient quantity of the gas has been driven off, A is cooled by means of water, whilst the water to be frozen is introduced into a metal cylinder, C, into the cavity of the receiver B, the space between receiver and cylinder being filled with alcohol, which does not freeze, and serves as a conducting medium. As the liquid in A cools, it rapidly reabsorbs ammonia, which boils off from B as fast as the pressure is removed, producing a great depression of temperature by means of the heat which becomes latent, and freezing the water contained in the metal cylinder. Ammonia is exceedingly soluble in water. Water at 0° C. absorbs more than 1100 times its volume of the gas, evolving great heat in the process. When the ammonia is pure, the absorption is instantaneous, the water rushing into the space occupied by the gas as into a vacuum. The affinity of the two substances for each other is nevertheless slight, as the solubility of ammonia in water decreases rapidily at higher tem- peratures, and the gas is completely expelled from the liquid by boiling. When exposed to the air, the aqueous solution also parts with nearly all its gas by diffusion. When ammonia is removed in the gaseous state from its solution, the heat which was liberated during the process of solution is again absorbed: thus by sending a rapid current of air from a foot blower through concentrated aqueous ammonia, the gas is expelled, and the temperature sinks belcrw —4o° C. (—4o° F.). Specific Gravity Table of Aqueous Ammonia at 14° C. d. P- d. P- 0.8844 36 0.9347 17 0.8864 35 0.9380 16 0-8885 34 0.9414 15 0.8907 33 0.9449 14 0.8929 32 0.9484 13 0.8953 31 0.9520 12 0.8976 30 0.9555 11 0.9001 29 0.9593 10 0.9026 28 0.9631 9 0.9052 27 0.9670 8 0.9078 26 0.9709 7 0.9106 25 0.9749 6 0.9133 24 0.9790 5 0.9162 23 0.9831 4 0.9191 22 0.9873 3 0.9221 21 0.9915 2 0.9251 20 0.9959 1 0.9283 19 0.9975 0.6 0.9314 18 0.9991 0.2 AMMONIA. 233 The foregoing table (Carius) gives the specific gravity of the aqueous solutions of ammonia of various strengths at 14° C. (57.2° F.). The column d contains the specific gravities, the column p the corresponding percentages of ammonia. Ammonia does not support combustion and does not burn in air unless the latter be heated. When mixed with oxygen, however, it is readily inflammable, burning with a pale yellow flame. At a bright red heat ammonia is decomposed into its elements. This decomposition, which is best effected by electric sparks, affords a means of ascertaining the composition of the gas. Reactions.—l. Ammonia is decomposed by chlorine (see p. 212). Bromine and iodine have a similar action. Under certain conditions, when chlorine and iodine are employed in excess, the explosive com- pounds, nitrous chloride and nitrous iodide (iq.v.) are formed. 2. When ammonia is passed over charcoal heated to redness in a tube, amnionic cyanide is formed and hydrogen is evolved: 2NH3 '+ C = NvH4(CN) + H2. Ammonia. Ammonic cyanide. 3. The metals of the alkalies, when heated in gaseous ammonia, replace the hydrogen atom for atom : NH3 + m = NNaH2 + H. Ammonia. Sodic amide, 4. Ammonia unites directly with acids, forming the ammonium salts, in which the atomicity of nitrogen is v : + HCI = NVH4CI. Hydrochloric Ammonic acid. chloride.* N'"HS + Nv02Ho = NvQ2(NvH40). Nitric acid. Ammonic nitrate.f 2N///H3 + S02Ho2 = S02(]SrvH40)2. Sulphuric acid. Ammonic sulphate.^ When a glass rod moistened with hydrochloric acid is brought close to a liquid evolving ammonia, white fumes of ammonic chloride are H I * H—N—II /\ H Cl H O I li f H—N—O—N. /\ II H H O H O H I II 1 % H—N—O—S—O—N—H. /\ I! /\ HHO H H 234 INORGANIC CHEMISTRY. formed. If the ammonia is in combination, the substance must be warmed with a solution of caustic alkali before applying this test. Amnionic chloride forms with platinic chloride a yellow crystalline double salt of the formula PtOJ4,2NH4CI, almost insoluble in water, and insoluble in alcohol or ether. This salt is employed in the quanti- tative determination of ammonia. Composition.—The composition of ammonia may be ascertained in the following manner. A measured volume of gaseous ammonia is intro- duced into an eudiometer tube over mercury. The tube is furnished with platinum wires fused into the glass for the purpose of passing the electric spark, which is furnished by an induction coil. The spark is allowed to pass through the gas as long as any increase of volume is observed. The resulting mixture of gases is then measured ; an excess of oxygen is added, and the whole is exploded by means of the spark. Two-thirds of the contraction which follows the explosion represents the volume of hydrogen contained in the mixture. The following example will illustrate the use of this method : The mixture of gases resulting from the decom- position of 100 cubic centimetres of am- monia is found to measure 200 c.c. Add 100 c.c. of oxvgen, 100 c.c. Total, 300 c.c. After explosion there remain 75 c.c. Contraction, .... 225 c.c. The hydrogen contained in the 200 c.c. is therefore f X 225 = 150 c.c., and the nitrogen is 200— 150 =6O c.c.; the two gases are therefore present in the proportion of 3 volumes of hydrogen to 1 volume of nitrogen. Further, as the mixed gases occupied twice the volume of the ammonia, it is evident that these 4 volumes in combining have undergone the normal condensation to 2 volumes. Expressing the vol- umes in litres: 1 litre of nitrogen weighs 14 criths. 3 litres of hydrogen weigh 3 “ The proportion by weight in which these elements are combined is therefore, 14 parts by weight of nitrogen to 3of hydrogen. Dividing each of these numbers by the atomic weight of the corresponding ele- ment, the atomic proportion 1: 3, represented by the formula NH3, is arrived at. AMMONIUM—HYDEOXYLAMINE. 235 AMMONIUM. /nh4 I nh4- This monad radical has never been obtained in the free state, but its compounds are perfectly analogous, in crystalline form and other prop- erties, to those of potassium. These facts have led some chemists to consider the group NH4 as a metal, to which they have given the name ammonium, a hypothesis which is considered to receive support from the production of an unstable amalgam of this radical. All the com- pounds of mercury with metals are found to possess metallic lustre; and this is the case with the amalgam of ammonium. It may be pre- pared by two different processes. 1. If a solution of amnionic chloride be electrolyzed, the negative electrotrode being mercury and the positive a platinum plate, the mer- cury is observed to swell up owing to the formation of a spongy metal- lic mass. The solution ought to contain an excess of ammonia, other- wise the explosive compound, nitrous chloride, may be formed at the positive electrode. 2. On pouring into a slightly warmed solution of amnionic chloride an amalgam of potassium or sodium, the amalgam is found to swell enormously, owing to its conversion into ammonium amalgam, whilst potassic or sodic chloride is simultaneously formed : HgJSam + mNH4CI = Hgu(NvH4)m 4- raNaCl. Sodic amalgam. Ammonic chloride. Ammonium amalgam. Sodic chloride. Ammonium amalgam rapidly decomposes into mercury, ammonia, and hydrogen, the ammonia and hydrogen being liberated in the pro- portion of 2NH3 to H2: 2Hgn(NvH4)rn = 2nHg -J~ 2mNHs -f- mH2. Ammonium amalgam. Mercury. Ammonia. Ammonium plays the part of a compound monad radical, and its salts are isomorphous with those of potassium; they are all volatile, unless the acid from which they are derived is fixed. They will be more fully described along with the metals of the alkalies. HYDEOXYLAMINE. NH2Ho. This remarkable compound, which was discovered by Lossen, may be regarded as ammonia in which one atom of hydrogen has been dis- placed by hydroxyl. Preparation.—l. Hydroxylamine is formed by the direct union of nitric oxide with nascent hydrogen : 236 INORGANIC CHEMISTRY. 2'N"O + 3H2 = 2NH2Ho. Nitric oxide. Hydroxylamine. Nitric oxide is passed into a mixture in which hydrogen is being gen- erated—thus into a flask containing tin and dilute hydrochloric acid. 2. Nitric and nitrous acids also yield hydroxylamine when added to the above reducing mixture : N02Ho + 3H2 == NII2Ho -f 20H2. Nitric acid. Hydroxylamine. Water. In these reactions the hydroxylamine remains in solution combined with the hydrochloric acid. Properties.—Free hydroxylamine is known only in its aqueous solu- tion, which is colorless, devoid of odor, and powerfully alkaline. On distilling the solution, part of the base passes over with the steam, but the greater part is decomposed with formation of ammonia. The solu- tion possesses reducing properties and precipitates silver and mercury in the metallic state from the solutions of their salts. Hydroxylamine is a mon-acid base. Its salts, which crystallize well, are formed, like those of all amine bases, by the direct union of base and acid without elimination of water. COMPOUNDS OF NITROGEN WITH CHLORINE, BROMINE, AND lODINE. NITROUS CHLORIDE. NC13? Preparation.—Nitrous chloride is formed when chlorine is passed into a solution of amnionic chloride warmed to about 30° C.: NvH4CI + 3C12 = N'"C1S + 4HCI. The same reaction takes places when a solution of ammonic chloride is electrolyzed, the chlorine which is evolved at the positive electrode acting on the ammonium salt. Properties.—Nitrous chloride is a yellow oil, of specific gravity 1.6, possessing a disagreeable, pungent odor. Its vapor irritates the eyes. Nitrous chloride is the most dangerously explosive substance known. A slight rise of temperature, or the mere contact with certain bodies— such as fats, phosphorus, or arsenic—is sufficient to cause it to decom- pose instantaneously with explosive violence into its elements. Very frequently explosion occurs without apparent cause. Ammonia decomposes it with formation of ammonic chloride and lib- eration of nitrogen. Its formation is therefore prevented by the pres- ence of an excess of ammonia (see Nitrogen, p. 213). THE ATMOSPHEEE. 237 The formula of this compound has not been ascertained with cer- tainty: it may contain hydrogen, and it is possible that the compounds intermediate between ammonia and nitrous chloride may exist: nh3, nh2ci, nhci2, nci3. NITROUS BROMIDE NBr3? This compound is obtained as a dark-red, very explosive oil by adding an aqueous solution of sodic or potassic bromide to nitrous chlo- ride : NC13 + 3KBr = NBr3 + 3KCI. Nitrous Potassic Nitrous Potassic chloride. bromide. bromide. chloride. NITROUS lODIDE. ni3. When aqueous or alcoholic ammonia is poured on finely powdered iodine, a black substance is formed which is highly explosive, and, when dry, detonates on the slightest touch. The product varies in composi- tion, according as aqueous or alcoholic ammonia is employed. A nitrous hydrodiniodide is formed at the same time: 4NH3 + 3I2 = NI3 + 3NHJ. Ammonia. Nitrous iodide. Ammonic iodide. 3NHS + 2I2 = NHI2 + 2NHJ. Ammonia. Nitrous Ammonic hydrodiniodide. iodide. THE ATMOSPHERE. The atmosphere of the earth consists of a mixture of gaseous, liquid, and solid matters. The chief gaseous constituents are nitrogen, oxygen, a small quantity of carbonic anhydride, and a varying proportion of aqueous vapor. Water also occurs in the liquid state in minute parti- cles in the form of mist. The solid matters consist of ice particles, volcanic and other dust, sporules and metallic salts—notably sodic chloride—in a finely divided state. The atmosphere is generally considered to extend to a height of about 45 miles above the earth’s surface, this estimate being based upon ob- servations of the length of time during which the twilight is visible in the zenith. Meteorites, however, ignite at an elevation of about 200 238 INORGANIC CHEMISTRY. miles, proving the presence of a medium which, though of too great tenuity to reflect light, still possesses density, and otfers resistance to the passage of bodies through it. It is probable that even this height does not denote the upper limit of the atmosphere. Owing to the effect of gravitation and the elasticity of the atmos- phere, the lower strata have a much greater density than the higher strata. If the density, instead of thus gradually decreasing with the elevation, were uniform throughout, and identical with that which pre- vails at the earth’s surface, the entire height of the atmosphere would be only about 5 miles. This diminution of density is such that at a height of about 3 miles the barometric pressure is only half as great as at the earth’s surface, and consequently one-half of the atmosphere lies below this height. According to the very accurate determinations of Regnault the weight of 1 litre of pure dry air at 0° C., and under a pressure of 760 milli- metres of mercury (the average barometric pressure at the level of the sea—a pressure commonly referred to as that of 1 atmosphere) in the latitude of Paris, is 1.2932 grams. Air is thus 773 times lighter than water, 10,500 times lighter than mercury, and 14.45 times heavier than hydrogen. A column of the height of the atmosphere and of 1 inch square weighs 15 lbs. Thus 27,000,000 tons rest upon every square mile of the earth’s surface. The luminous rays of the sun pass through the atmosphere without appreciably heating it, except in so far as they are intercepted and ab- sorbed by suspended solid or liquid matter; but the rise of temperature from the latter cause is not great. The dark heat-rays, however, are partly absorbed, and this absorption is due to aqueous vapor. These dark rays represent, however, but a fraction of the total radiant en- ergy of the sun, of which the greater part therefore reaches the earth unimpaired. Here both the visible and the invisible rays are converted by absorption into heat; and radiation from the earth’s surface in the form of dark heat is for the most part intercepted by aqueous vapor. In this way, the earth which has been heated by the sun imparts its heat to the air immediately resting upon it, and the aqueous vapor acts as a trap for the solar rays, allowing them to enter freely in the form of luminous heat, but preventing their escape when they are once con- verted into dark heat. Thus a too rapid cooling of the earth’s surface during the absence of the sun, and the consequent great inequalities of temperature, are prevented. The air, thus heated by contact with the earth, expands, and, becoming lighter, rises, and shares its heat with the strata above, whilst air from some colder quarter flows in to supply its place. The air is thus in constant motion, and differences in com- position of the atmosphere in various places, which might, arise from local causes, are prevented. To this heating and cooling, and to the varying quantities of aqueous vapor present in hot and in cold air, the variations of the barometric pressure are due. Equalization of tempe- rature is also effected by the condensation of aqueous vapor during a fall of temperature, the latent heat of vaporization being recovered in this process. The highest atmospheric temperature (temperature in shade) that has 239 THE ATMOSPHERE. been observed is about 49° C. (120.2° F.); the lowest —49° C. (—56.2° F.). The expansion of air by heat is 0.003665 of its volume measured at 0° C. for every 1° above 0° C. As regards the chemical composition of the atmosphere, the propor- tion of oxygen to nitrogen is nearly constant; the proportions of the other constituents are subject to considerable variation. The following table contains determinations of the relative quantities of oxygen and nitrogen present in dry air freed from carbonic anhydride. As is usual in the analysis of gaseous mixtures, the results are expressed in parts by volume. Composition of Atmospheric Air from various Localities. In 100 parts by volume. Oxygen. Parts by volume. Nitrogen. Parts by volume. St. Bartholomew’s Hospital, . . -j f 20.885 L 20.999 79.115 79.001 Paris, j r 20.913 L 20.999 79.087 79.001 Lyons, ■{ 1 20.918 L 20.966 79-.082 79.034 T onion, ■< f 20.912 l 20.982 79.088 79.018 Berlin, < ( 20.908 [ 20.998 79.092 79.002 Madrid, -1 f 20.916 [ 20.982 79.084 79.018 Geneva, \ f 20.909 ( 20.993 79.091 79.007 Montanvert, ....... 20.963 79.037 Summit of Pichincha, 16,000 ft. . j f 20.949 L 20.988 79.051 79.012 North American Prairie, 20.910 79.090 South America, 20.960 79.040 Liverpool to Vera Cruz, . . . ■< f 20.918 ( 20.965 79.082 79.035 18,000 ft, above London, . . . 20.885 79,115 Manchester, ■< f 20.876 ( 20.888 79.124 79.112 Algiers (June 5, 1851), . . . . - J 20.420 ( 20.395 79.580 79.605 Bay of Bengal (Feb. 1, 1849), . ■< f 20.460 ( 20.450 79.540 79.550 Ganges (March 8, 1849), . . . - f 20.387 \ 20.390 79.613 79.610 These analytical results, except in the case of the three localities last mentioned, display a remarkable uniformity. The cause of the varia- tion in the case of the sample from Algiers is unexplained ; but as re- 240 INORGANIC CHEMISTRY. gards the sample from the Bay of Bengal and the Ganges, it is to be noted that these were collected during an outbreak of cholera when the water contained large quantities of putrefying organic matter.* The presence of a very small quantity of the oxygen as ozone has al- ready been referred to (p. 166). The average proportion of carbonic anhydride present in air is about 0.03 per cent.; but the amount may vary considerably owing to local causes. Thus the effect of animal life is to increase the proportion of car- bonic anhydride; that of vegetable life to diminish it (see p. 202). In putrefaction and in combustion, large quantities of this gas are given off. In London, combustion and respiration daily send into the air at least 200,000,000 cubic feet of carbonic anhydride. Each ton of coal consumed furnishes about 3 tons of carbonic anhydride, and abstracts 2.75 tons of oxygen from the air. The variations due to the above causes are very noticeable: thus in crowded and ill-ventilated rooms, the air may contain as much as 0.3 per cent, of carbonic anhydride; air from the centre of London contains 0.11 per cent. Near the surface of the ocean, both oxy- gen and carbonic anhydride are slightly in excess during the day, and slightly deficient during the night. This is due to the fact that these gases are more soluble in water than nitrogen: in the night time the cold water dissolves them in larger quantity, and this dissolved excess is again expelled when the water is heated by the sun’s rays during the day. At great altitudes the proportion of carbonic anhydride appears to increase: thus the air at the Grands Mulcts was found to contain 0.1 per cent. The proportion of aqueous vapor present in the air varies greatly. The maximum quantity of aqueous vapor which a given volume of air can take up is constant for a given temperature, and independent of the pressure. When air has taken up this maximum quantity it is said to be saturated with moisture. The amount necessary for saturation at a given temperature can be calculated from the tension of the vapor of water for that temperature. In this way it is found that 1 cubic metre of air can take up the following weights of aqueous vapor: At 0° C. ( 32° F.),. . . . At 10° C. ( 50° F.),. . . . At 20° C. ( 68° F.),. . . . At 30° C. ( 86° F.), . . . . At 40° C. (104° F.), . . . . The air is very seldom saturated with moisture. When the tempera- ture of air containing aqueous vapor falls, as soon as the point is passed at which the quantity of aqueous vapor present corresponds to satura- tion, a separation of the excess of this vapor in the form of mist, rain, snow, or hail begins. This point is known as the dew-point, and by * The oxygen in the foregoing samples was determined by exploding the air with hydrogen and noting the contraction. If, as is quite conceivable, the air in the above abnormal cases contained traces of marsh-gas derived from the decomposition of organic matter, a smaller contraction would be observed, and the percentage of oxygen would be found too low. THE ATMOSPHERE. 241 determining it accurately, the quantity of aqueous vapor present in in- completely saturated air may be ascertained. The usual proportion by volume of aqueous vapor in air, varies from 0 to 5 per cent. The question of the proportion of aqueous vapor present in air is of great importance in meteorology; but in the chemical examination of air the aqueous vapor is taken into account only in so far as by its vol- ume it diminishes the absolute quantity of the other constituents present in a given bulk. It is usual, in the analyses of gases, to eliminate the aqueous vapor from the result by calculation. Other constituents of the air, which are, however, present only in minute quantity, are salts of ammonia, namely, the carbonate, nitrate, and nitrite. Ammonia is given off in the putrefaction of animal and vegetable matter. Oxides of nitrogen are formed whenever a flash of lightning passes through air : rain-water, especially if collected after a thunderstorm, contains nitrates and nitrites. The presence of these nitrogenous compounds in the air is of great importance to plant life, as it is from this source alone that plants which have not been supplied with a nitrogenous manure obtain the nitrogen necessary for their growth. Plants cannot assimilate free nitrogen. Another product of putrefaction which is constantly being given off into the air is marsh-gas. It is doubtful, however, whether the pres- ence of this compound in air has been proved, except in the neighbor- hood of putrefying matter. Although the various gases which together make up the atmosphere possess very different specific gravities, they display no tendency to sepa- rate from each other. On the contrary, by the laws of diffusion, any number of gases which are brought into contact have a tendency to be- come thoroughly mixed, even although there are no actual currents in the gases, and even although the lighter gases may be uppermost at the commencement of the process. The influence of currents of air in pre- serving uniformity of composition has already been referred to. As regards the suspended matter in the atmosphere, this may, as already stated, be both solid and liquid. These particles, even when present in small quantity, are rendered visible to the eye by their property of reflecting light; thus when a ray of light passes through a dark room, the path of the ray appears luminous. By filtration through cotton wool, or by subsidence, the particles are removed, and the path of a ray of light through air thus purified ceases to be visible. These particles are never absent from air under ordinary conditions. When solid particles are present in quantity sufficient to obstruct visibly the passage of light, they constitute a dust-haze. Piazzi Smyth ob- served a strong dust-haze on the summit of the Peak of Teneritfe at an altitude of 12,000 feet. Minute liquid particles constitute ordinary mist or fog. When the surface of the sea is violently agitated by the wind, particles of sea-water are thrown into the air in the form of spray : these are carried far inland by the wind, yielding by evapora- tion solid particles of sea salt, a substance which is scarcely ever absent from air. The yellow flashes which a Bunsen flame emits from time to time, while burning in air, are due to sodium compounds, as may be proved by spectroscopic examination. In the neighborhood of the sea 242 INORGANIC CHEMISTRY. the quantity of sodic chloride present in air is of course greater than further inland. At Land’s End for example, the rain water contains as much as 0.033 per cent, of this salt. At great altitudes in Switzer- land the air almost always contains minute particles of snow, which may be seen by putting the eye in shadow and looking into sunshine. Among the organic solid particles present in air, are to be reckoned the germs of putrefactive and other fermentations. This is shown by the fact that air which has been effectually freed from all suspended matter by filtration does not induce putrefaction in milk, flesh, urine, and other readily alterable animal substances, however long these may be left in contact with it. If a Bunsen flame be placed under the path of a ray of light in a dark room, the heated air rising from the flame appears like a black smoke, owing to the absence of suspended matter in the products of combus- tion. The same phenomenon may be shown, though in a less striking manner, by substituting for the flame a flask filled either with oil heated to I2o°—l3o° C. (248°-266° F.), or with ice-cold water, and concentrating the ascending or descending current of air upon the path of the ray by means of a conical paper funnel. This phenomenon has not yet received any satisfactory explanation. It has been shown by Lodge that the electrification of air also rapidly removes the suspended particles contained in it. That the oxygen and nitrogen, which form the chief constituents of the atmosphere, are present in a state of mere mechanical mixture and not, as was formerly supposed, in chemical combination, is proved by a variety of considerations. Thus the proportion by volume of the two gases to each other is highly complex, 21 volumes of oxygen to 79 vol- umes of nitrogen being the simplest proportion that can be assumed; whereas in compounds of only two elements much simpler relations prevail. No contraction occurs when oxygen and nitrogen form air, and there is no case known in which two gases unite chemically in un- equal proportions by volume without contraction. When oxygen and nitrogen are mixed in the above proportions, no heat is evolved, nor is there any other sign of chemical combination; nevertheless, the mix- ture displays all the properties of air. When air is dissolved in water, the proportion of its constituents is totally altered, owing to the greater solubility of oxygen ; thus, dissolved air contains, in 100 volumes, 32.5 volumes of oxygen and 67.5 volumes of nitrogen. Again, air which has been forced through a thin caoutchouc membrane contains 41.6 volumes of oxygen to 58.4 volumes of nitrogen, owing to the property which oxygen possesses of passing more readily through caoutchouc. If air were a chemical compound, the proportion of its constituents could not be thus altered by solution or by osmosis. 243 SULPHUR. CHAPTER XXVII HEXAD ELEMENTS. Section I. SULPHUR, S2. Atomic weight 32. Molecular weight = 64. Molecular volume I 11 at 1000° C., but only one-third of this at its boiling-point. 1 litre of sulphur vapor weighs 32 criths. Rhombic variety fuses at 114.5° C. (238.1° F.). Boils at 445° C. (833° F.). Atomicity ", iT, and Ti. Evidence of atomicity: Sulphuretted hydrogen, Triethylsulphinic iodide, S"H2. SiTEt3I. Sulphuric oxydichloride (Sulphurylic chloride), Svi02Cl2. Sulphuric iodide, S"I«- Sodic dinitrosulphate, SviO(NO)2Nao2. History.—This element has been known from the earliest historical times. Occurrence.—Sulphur occurs both native and in combination. Na- tive sulphur is found chiefly in the neighborhood of volcanoes: thus in Sicily, whence the greater part of the native sulphur of commerce is obtained. In combination it occurs either with metals alone as sul- phides, or with metals and oxygen as sulphates. Of the former the most important as sources of sulphur are ferric disulphide or iron py- rites (FeS"2) and copper pyrites (Fe2Cu2S4). The sulphides of zinc, lead, mercury, and antimony are important ores of these metals. The most commonly occurring sulphates are calcic sulphate—which is found in two forms, as gypsum and as anhydrite (So2Cao")— baric sulphate, or heavy spar (SOgBao"), and magnesic sulphate, which also occurs in two forms, as kieserite (SOHo2Mgo/r),and as Epsom salts (SOHo2Mgo",6OH2). The sulphates of calcium, magnesium, and so- dium occur in natural waters. Of gaseous compounds, both sulphurous anhydride (SO2) and sulphuretted hydrogen (SH2) are of frequent oc- currence in volcanic exhalation, the latter being also found in many mineral waters. Sulphur is a constituent of many complex organic compounds in the animal and vegetable kingdoms. Formation of Volcanic Sidphur.—This is probably due to the mutual decomposition of the two volcanic gases, sulphurous anhydride and sulphuretted hydrogen, in presence of water. In this reaction penta- thionic acid and water are formed, whilst sulphur is liberated: f S02Ho 5S02 + 5SH2 = J (S"3)" + 40H2 4 SS. (SO2Ho Sulphurous Sulphuretted Pentathionic Water, anhydride. hydrogen. acid. 244 INORGANIC CHEMISTRY. Manufacture.—1. Native sulphur is usually mixed with large quan- tities of earthy matters, from which it is separated by fusion. In Sicily, the heat for this purpose is obtained by the combustion of a portion of the sulphur itself. The sulphur ore is built up into a large heap over a pit sunk into the ground. The heap is ignited from beneath, and as the heat slowly penetrates through the mass, the sulphur melts and flows into the pit, which is so arranged that the liquid product can be drawn off during the process. By this method more than half the sul- phur burns away as sulphurous anhydride. 2. Sulphur is also obtained by distilling iron pyrites: 3FeS2 = iv(Fe3)TiiiS4 + S2. Ferric Triferric disulphide. tetrasulphide. The reaction is analogous to that which takes place in the preparation of oxygen from manganic peroxide (p. 161). The distillation is per- formed in fire-clay cylinders. It is, however, in every way more eco- nomical to burn the pyrites in kilns, a method which has been generally adopted. The kiln is lighted from below; part of the sulphur which, in the process of distillation in cylinders, remains in combination with the iron, burns, forming sulphurous anhydride; the remainder distils off and is condensed. The exhausted pyrites is from time to time with- drawn from the lower part of the kiln, and a fresh charge is introduced at the top, thus rendering the process more continuous. By this method one-half of the total sulphur is obtained from the pyrites: 3FeS2 + 502 = iv(Fe3)viiiO, + 3S02 + 3S. Ferric Triferric Sulphurous disulphide. tetroxide. anhydride. By passing the products through heated charcoal a larger yield can be obtained. Sulphur is obtained in a similar manner from copper pyrites in the process of roasting the ore in the first stage of copper-smelting. 3. The alkali-waste obtained in the manufacture of sodic carbonate (q.v.) may be made to yield considerable quantities of sulphur. This waste, which remains after the extraction of the sodic carbonate from the black-ash by lixiviation, consists essentially of insoluble calcic oxy- sulphide, a combination of calcic sulphide with calcic oxide in varying proportions. Without removing the waste from the lixiviating vats, a current of air is blown through it, by which means the calcic sulphide contained in the oxysulphide is oxidized with considerable rise of tem- perature, yielding a mixture of soluble polysulphides of calcium and calcic thiosulphate: {a.) 20aS + O + 0H2 = Oa(S")" + CaHo2. Calcic sulphide. Water. Calcic disulphide. Calcic hydrate. (b.) Oa(S"j)" + 30 = S02J®Ca)''. Calcic disulphide. Calcic thiosulphate. SULPHUR. 245 Calcic sulphide is also liberated from its combination with the lime and becomes soluble. The oxidation is allowed to proceed till one-half of the sulphur has been converted into thiosulphate, and the remainder into calcic sulphide or polysulphide, after which the whole is lixiviated, and the solution treated with hydrochloric acid. The sulphur is lib- erated, as represented by the following equation : 2CaS -f S02|gCaj" + 6HCI = 4S + 3CaCJ2 -f 30H2. Calcic Calcic Hydrochloric Calcic Water, sulphide. thiosulphate. acid. chloride. Calcic polysulphides undergo an analogous decomposition with calcic thiosulphate when treated with hydrochloric acid, but the quantity of sulphur liberated is proportionately larger. The sulphur thus obtained is melted under superheated water. 4. Sulphur is obtained in the purification of coal-gas. The crude gas contains sulphuretted hydrogen. In order to remove this impurity, the gas is passed through ferric hydrate, which absorbs the sulphuretted hydrogen with formation of ferrous sulphide and separation of sulphur : Fe2Ho6 + 3SH2 = 2FeS + S + 60H2. Ferric Sulphuretted Ferrous Water, hydrate. hydrogen. sulphide. When the mixture has lost its absorptive power, it is exposed to the air in a moist state, ferric hydrate being thus regenerated and sulphur set free: 2FeS -f 30H2 + 30 = Fe2Ho6 + Ferrous sulphide. Water. Ferric hydrate. In this condition the mixture is again employed in the removal of sul- phuretted hydrogen. These alternate processes of absorption and oxi- dation are repeated till the mixture contains half its weight of sulphur, when the latter is separated by distillation. Refining.—Crude sulphur is generally contaminated with earthy im- purities, from which it is separated by distillation. The operation is conducted as shown in Fig. 39. The crude sulphur is first introduced into the iron pot A, where it is melted. The greater part of the im- purities sink to the bottom, and the melted sulphur is run off into the retort B, whence it is distilled into the large brick-work chamber C. When the distillation is conducted rapidly, so as to keep the tempera- ture of the chamber above the melting-point of sulphur, the latter con- denses in the liquid state and collects on the floor of the chamber, whence it may be drawn off by the tap D, to be run into slightly conical box-wood moulds. The sulphur thus obtained is known as roll sulphur. When the distillation proceeds slowly, and the temperature of the chamber is consequently lower, the sulphur is deposited as a fine crystal- 246 INORGANIC CHEMISTRY line dust on the walls and floor of the chamber. This form is termed flowers of sulphur. Fig. 39. Properties.—Sulphur is capable ot existing in several allotropic modi fications, of which the following are the most important: Specific Behavior with carbonic Condition. gravity. disulphide. a. Rhombic (octahedral), . . 2.05 Soluble. fi. Monoclinic (prismatic), . . 1.98 Transformed into «. y. Plastic, . 1.95 Insoluble. 8. Powder, . 1.95 Insoluble. The rhombic form is that in which sulphur occurs in nature. This form displays great variety of crystalline combinations: the most fre- quently occurring combination, in which the rhombic octahedron is dominant, is shown in Fig. 40. Rhombic sulphur is insoluble in water, somewhat soluble in alcohol, ether, and hydrocarbons, readily soluble in carbonic disulphide and disulphur dichloride. From these solvents it is again deposited in the rhombic form. Rhombic sulphur fuses at 114,5° C. (238.1° F.). The behavior of melted sulphur is anomalous. Just above its fusing- point it forms a clear, yellow, mobile liquid; but on raising the tem- perature the color deepens, changing to a reddish-brown, whilst the liquid becomes viscid. At about 230° C. (446° F.) it is almost black, SULPHUR. 247 and is so thick that the vessel in which it is contained may be inverted without spilling the contents. Heated above this temperature it again becomes liquid, still however preserving its dark color, till at 447° C. (836° F.) it boils, giving off* a reddish-brown vapor. One litre of this vapor at 524° C. (975° F.) weighs 96 criths, whereas above 860° C. (1580° F.) the weight of 1 litre of sulphur vapor is 32 criths, or only Fig. 40. one-third of the first-mentioned value. From this it follows that above 860° C. the molecule of sulphur contains two atoms, but that just above its boiling-point, the molecule is hexatomic. If sulphur, heated to its boiling-point, be allowed to cool gradually, the above changes are observed in the reverse order. The rhombic variety of sulphur may also be obtained by melting sul- phur in large masses, and, by slow cooling with exclusion of air, allow- ing it to remain in a state of superfusion, or suspended solidification. Fig. 41. At a temperature of about 90° C. (194° F.), the superfused sulphur deposits rhombic crystals. If the melted sulphur be allowed to cool more rapidly, the second or monoclinio variety is obtained. This last experiment is best performed by fusing about a kilogram of sulphur in a Hessian crucible, and allowing it to cool till a crust has been formed over the surface. Two holes are then broken in this crust, and the crucible is inclined so as to allow the sulphur which still remains liquid to run out. The interior of the crucible (Fig. 41) is found to be lined with long thin transparent prisms, belonging to the monoclinic system. These fuse at 120° C. (248° F.). The system in which sulphur crystallizes is determined by the con- 248 INORGANIC CHEMISTRY. ditions of temperature under which the crystallization occurs, and the crystals of each system are unstable at the temperature of formation of those of the other system. Thus, when a transparent crystal of rhom- bic sulphur, which has been deposited at ordinary temperatures, is ex- posed for some time to a temperature just below its fusing-point, it loses its transparency and, on examination, is found to have been converted into an aggregation of minute monoclinic crystals. On the other hand, the transparent crystals of monoclinic sulphur, which are formed at a higher temperature, become opaque after remaining for some time at the ordinary temperature, having changed into aggregations of small rhora-* bic crystals. This latter change may also be effected by scratching the monoclinic crystals: in this case the transformation takes place rapidly, and is found to be accompanied by a liberation of heat. The rhombic modification is that into which all other forms of sulphur (except the 5 variety) spontaneously change at ordinary temperatures. If melted sulphur at a temperature just above its fusiug-point be poured into cold water, it solidifies to a yellow, brittle mass. But if the temperature of the melted sulphur be raised above the point of maximum viscosity, and the dark-colored mobile liquid thus obtained be poured in a thin stream into water so as to effect its cooling as rap- idly as possible, a totally different phenomenon is observed. Under these conditions, the sulphur forms plastic, amber-colored, transparent threads, which may be drawn out or kneaded between the fingers. This is the variety known as plastic sulphur. After standing for some time at ordinary temperatures it becomes brittle and opaque. At a temper- ature of 100° it is suddenly converted into rhombic sulphur, the change being accompanied by evolution of heat. If the brittle sulphur resulting from the spontaneous change of the plastic variety be treated with carbonic disulphide, part of it is dissolved, whilst part remains behind as a brown amorphous powder. A light yellow amorphous powder, insoluble in carbonic disulphide, is also ob- tained by treating flowers of sulphur with this solvent as long as any- thing is dissolved. The same insoluble variety separates out when a solution of sulphur in carbon disulphide is exposed to sunlight concen- trated by means of a lens. At a temperature of 100° C., these amor- phous varieties pass into the ordinary rhombic modification. All the varieties of sulphur are insoluble in water. The so-called milk of sulphur is nothing more than sulphur in a finely divided state, obtained by decomposing calcic pentasulphide, or any other polysulphide, with hydrochloric acid: ™(SS)"Ca + 2HCI = CaCl2 + SH2 + 4S. Calcic Hydrochloric Calcic Sulphuretted pentasulphide. acid. chloride. hydrogen. It is soluble in carbonic disulphide, and is probably the rhombic variety. Reactions—1. When heated in air or oxygen to its temperature of ignition, sulphur burns with a blue flame, forming sulphurous anhy- dride : 249 SULPHURETTED HYDROGEN. S -f- 02 S02. Sulphurous anhydride. A slow, phosphorescent combustion occurs when sulphur is heated to about 180° C. (356° F.) in air. No flame is visible in daylight; but in the dark a grayish-white flame, quite distinct from the ordinary blue flame of burning sulphur, appears to hover over the heated surface. The product of combustion is in this case also sulphurous anhydride. In presence of air and moisture, finely divided sulphur is spontane- ously oxidized at ordinary temperatures to sulphurous and sulphuric acids. 2. Sulphur also unites directly with chlorine, bromine, iodine, phos- phorus, hydrogen, and various other non-metals. 3. It combines directly with many metals when heated with them, forming sulphides: k2 + s = sk2. Potassic sulphide. Fe -f S = FeS. Ferrous sulphide. When united exclusively with positive elements or radicals, sulphur is almost invariably a dyad; it is then analogous to oxygen, as will be seen from the following formulae: Oxygen compounds, . . ok2, OKH, 0O2, COKo,. Sulphur “ . sk2, SKH, CS"2, OSKs2.“ Uses.—Sulphur is employed in the arts in the manufacture of gun- powder and for tipping common lucifer matches. In the form of sul- phurous anhydride it is a useful bleaching agent. Its most important application, however, is in the manufacture of sulphuric acid. COMPOUNDS OF SULPHUR WITH HYDROGEN. Sulphuretted hydrogen, .... . . sh2. Hydrosulphyl, . . 'S'2H2 or Hs2. Hyposulphurous hydrosulphate, . . SHs2. SULPHURETTED HYDROGEN, Hydrosulphur ic Acid, Sulphy- dric Acid. SH2. H—S—H. Molecular weight = 34. Molecular volume I I I. 1 litre weighs 17 criths. Solid at —85.5° C. (—121.9° F.). Liquefied under a pressure of 17 atmospheres at 10° C. (50° F.). History.—This compound was first investigated by Scheele. Occurrence.—Sulphuretted hydrogen is evolved along with other gases from volcanoes and furaaroles. It occurs also in hepatic mineral 250 INORGANIC CHEMISTRY. waters, such as those of Harrogate, and in waters which contain sul- phates along with organic matters. Formation and Preparation.—1. Sulphuretted hydrogen is formed in small quantity by the direct union of its elements when hydrogen, together with the vapor of sulphur, is passed through a red-hot tube, or even when hydrogen is passed into boiling sulphur: h2 + s = sh2. Sulphuretted hydrogen. 2. The most convenient method of preparing the gas for laboratory purposes consists in acting on ferrous sulphide with dilute sulphuric acid: FeS" + S02Ho2 = SH2 + S02Feo" Ferrous Sulphuric Sulphuretted Ferrous sulphide. acid. hydrogen. sulphate. The ferrous sulphide, broken into coarse fragments, is introduced into a flask similar to that used in the preparation of hydrogen, and the acid, diluted with about 6 times its bulk of water, is poured in through a funnel. The gas is washed by passing it through water. Hydrochloric acid may be substituted for sulphuric acid in the above reaction : FeS" + 2HCI = SH2 + FeCl2. Ferrous Hydrochloric Sulphuretted Ferrous sulphide. acid. hydrogen. chloride. The use of sulphuric acid is, however, much more convenient in practice. 3. Sulphuretted hydrogen prepared from ferrous sulphide generally contains free hydrogen, generated by the action of the acid upon metallic iron, which is often present in the sulphide as an impurity. Pure sul- phuretted hydrogen may be obtained by decomposing precipitated anti- monious sulphide, or native antimonious sulphide (gray antimony ore), with hydrochloric acid aided by a gentle heat: Sb2S"3 + 6HCI = 3SH2 + 2SbCl3. Antimonious Hydrochloric Sulphuretted Antimonious sulphide. acid. hydrogen. chloride. If the native compound be employed, it ought to be first treated with dilute hydrochloric acid, in order to remove any carbonates that may be present. 4. Sulphuretted hydrogen is formed in small quantity along with sulphurous anhydride when steam is passed over boiling sulphur, or even when sulphur is boiled with water: 3S + 20H2 = 2SH2 + S02. Water. Sulphuretted Sulphurous hydrogen. anhydride. The sulphuretted hydrogen and sulphurous anhydride mutually decom- pose each other in the distillate with separation of sulphur, only a por- tion of the former gas remaining (see p. 243). SULPHURETTED HYDROGEN. 251 5. Sulphuretted hydrogen is formed when sulphur is heated along with paraffin, aniline, and various other organic bodies. The reac- tions which take place in these cases are very complicated and cannot be followed by means of equations. 6. It is evolved during the putrefaction of organic bodies containing sulphur, and also when these bodies are subjected to destructive distil- lation. It thus finds its way into illuminating gas, from which it has to be removed in the process of purification. Properties.—Sulphuretted hydrogen is a colorless gas possessing the disgusting odor of putrid eggs. The very offensive odor of the gas pre- pared from ferrous sulphide is, however, in part due to the presence of volatile sulpho-carbon compounds derived from the iron. It is slightly heavier than air. It is combustible, burning in air or oxygen with a bluish flame, and forming sulphurous anhydride and water: SH2 + 30 = S02 + 0H2. Sulphuretted Sulphurous Water, hydrogen. anhydride. When the supply of oxygen is insufficient for complete combustion, water only is formed, and sulphur is deposited. Water absorbs about three times its volume of sulphuretted hydrogen, yielding a colorless solution possessing the taste and odor of the gas. The aqueous solution is a useful laboratory reagent. It parts with the whole of its gas on boiling. Exposed to the air, the gas in solution is quickly oxidized with separation of sulphur, water being formed at the same time. Sulphuretted hydrogen has a powerfully poisonous action when in- haled, especially in the case of small animals. The intensity of the action in various animals appears to be connected with the rapidity of circulation of the blood. An atmosphere containing °f the gas suffices to kill a bird, whilst is necessary to kill a dog, and to kill a horse. Cold-blooded animals are totally unaffected by this pro- portion of sulphuretted hydrogen. Reactions.—l. Sulphuretted hydrogen is immediately decomposed by chlorine with separation of sulphur : SH2 + Cl2 = 2HCI + S. Sulphuretted Hydrochloric hydrogen. acid. A similar reaction takes place with bromine. In the case of iodine, the formation of hydriodic acid and the liberation of sulphur take place only in the presence of water. The reason of this is that the reaction SH2 + I 2 = 2HI + s, Sulphuretted Hydriodic hydrogen. acid. 18 attended with an absorption of heat, and consequently, according to the laws of thermochemistry (p. lift), cannot take place without the 252 INORGANIC CHEMISTRY. aid of some extraneous energy. When water is present, the heat evolved by the absorption of the hydriodic acid by water, furnishes this energy ; the thermal sign of the equation becomes positive and the reaction possible. 2. Sulphuretted hydrogen is decomposed by many compounds rich in oxygen, such as ferric hydrate : 'Fe'"2Ho6 + 3SH2 = 2FeS" + S + fiOH2. Ferric Sulphuretted Ferrous Water, hydrate. hydrogen. sulphide. This reaction is employed on a large scale in the purification of coal- gas (see p. 245). In like manner it reduces concentrated sulphuric acid, which cannot therefore be employed in drying the gas: SH2 + S02Ho2 = S02 + S + 20H,. Sulphuretted Sulphuric Sulphurous Water, hydrogen. acid. anhydride. Fuming nitric acid, when dropped into ajar of sulphuretted hydrogen, oxidizes it with explosive violence. 3. The sulphydrates and sulphides of the metals are produced by the action of sulphuretted hydrogen on the hydrates and oxides; thus: OKH + SH2 = SKH + 0H2. Potassic Sulphuretted Potassic Water, hydrate. hydrogen. sulphydrate. BaHo2 + 2SHa = BaHs2 + 20H2< Baric Sulphuretted Baric Water, hydrate. hydrogen. sulphydrate. OAg2 + SH2 = SAg2 + 0H2. Argentic Sulphuretted Argentic Water, oxide. hydrogen. sulphide. CuO + SH2 = CuS + 0H2. Cupric Sulphuretted Cupric Water, oxide. hydrogen. sulphide. Upon this property, and upon the varying behavior of the different metallic sulphides towards weak acids, is based the use of sulphuretted hydrogen as a reagent in analysis. Some of these sulphides are insol- uble in weak acids: sulphuretted hydrogen, therefore, precipitates them from an acid solution of the salts of their metals: S02Cuo" + SH2 = CuS -f S02Ho2. Cupric Sulphuretted Cupric Sulphuric sulphate. hydrogen. sulphide. acid. Others are soluble in weak acids, but insoluble in alkaline solutions. The precipitation of these sulphides is most conveniently effected by the SULPHURETTED HYDROGEN. 253 addition of an alkaline sulphide (ammonic sulphide is most commonly employed for this purpose) to the neutral or alkaline solution of the salt, when double decomposition takes place, thus: ZnCl2 + S(NH4)2 = ZnS + 2NH4CI. Zincic Amnionic Zincic Ammonie chloride. sulphide. sulphide. chloride. A third class of metals yields sulphides which are soluble in water, and are therefore not precipitated either in acid or in alkaline solutions. It is thus possible to divide the metals into three groups, according to the behavior of their sulphides, and this division forms one of the foundations of inorganic qualitative analysis. 4, Most metals when heated in sulphuretted hydrogen combine with the sulphur to form sulphides, whilst hydrogen is liberated: Sn + SH2 = Sn"S -f H2. Sulphuretted Stannous hydrogen. sulphide. Silver becomes tarnished when exposed at (ordinary temperatures to the action of sulphuretted hydrogen in presence of air, owing to the formation of a superficial coating of argentic sulphide {q. v.), but the ac- tion is very slow unless moisture be present. Composition.—The composition of sulphuretted hydrogen is best ascertained by heating in it some metal which combines with the sul- phur liberating the hydrogen. Tin is usually employed for this purpose (see above). (Potassium or sodium cannot be used, as in these cases the metal displaces only one-half of the hydrogen, combining with a semi-molecule of hydrosulphyl to form a sulphydrate.) The operation is performed in a bent tube over mercury as described in the analysis of hydrochloric acid (p. 159). After the action is complete and the tube has been allowed to cool, it will be found that the hydrogen occupies exactly the same volume as the sulphuretted hydrogen employed. Sul- phuretted hydrogen thus contains its own volume of hydrogen. There- fore:: Weight of 1 litre of sulphuretted hydrogen, Deduct weight of 1 litre of hydrogen,, . . 17 criths. 1 crith. There remain ...... 16 criths, which is the weight of half a litre of normal sulphur vapor. Calcu- lating to whole volumes, 2 volumes of hydrogen combine with 1 volume of sulphur vapor to form 2 volumes of sulphuretted hjalrogen. By weight, the proportion of hydrogen to sulphur is as 1 : 16 or as 2 : 32, and the formula of the compound is therefore SH2. 254 INORGANIC CHEMISTRY. HYDROSULPHYL, Hydrio Persulphide. 'S'2n2 or Hs2. H—S—S—H. Probable molecular weight = 66. Sp. gr. 1.769. Preparation.—When a solution of calcic disulphide is poured into an excess of cold concentrated hydrochloric acid, hydrosulphyl separates out as a heavy yellowish oil: 'S'2Ca" + 2HCI = 'S'2H2 + CaCl2. Calcic Hydrochloric Hydrosnlphyl. Calcic disulphide. acid. chloride. The calcic disulphide is prepared by boiling milk of lime with an excess of sulphur and filtering. The solution must be poured into the acid, and not the reverse, as hydrosulphyl is much more stable in con- tact with acids than in contact with alkalies. The calcic disulphide prepared as above, is always mixed with higher polysulphides, but these also yield hydrosulphyl, mixed however with sulphur. Properties.—Hydrosulphyl is a heavy yellowish liquid possessing a fetid odor. It closely resembles hydroxyl in its properties, bleaching organic coloring matters and reducing argentic oxide. It is very un- stable, and is gradually decomposed into sulphuretted hydrogen and free sulphur. Owing to this fact and to the property which hydro- sulphyl possesses of dissolving sulphur, it has been found almost im- possible to obtain it in a state of purity, and its composition is more a matter of conjecture, based upon its analogy with hydroxyl, than a strict analytical result. HYPOSULPHUROUS HYDRO SULPHATE. f Hs SHs2, or 4 S". (Hs Probable molecular weight = 98. Preparation.—When a cold saturated solution of strychnine in alcohol is mixed with an alcoholic solution of yellow ammonic sulphide, a compound is formed crystallizing in orange needles of the formula B2|H22N202,H2S3. By the action of concentrated sul- phuric acid upon this compound, and subsequent dilution with water, hyposulphurous hydrosulphate is liberated as a yellow oily body. It closely resembles in its proper- ties hydrosulphyl, and, like that substance, undergoes spontaneous decomposition into sulphuretted hydrogen and sulphur. COMPOUNDS OF SULPHUR WITH THE HALOGENS. Disulphur dichloride, . . . 'S'2C12. Hyposulphurous chloride, . . . . . . . SCI, Sulphurous chloride, . . . SC14. Disulphur dibromide, , . . . 'S'2Br2. Disulphur diniodide, . . . 'S'aIr Sulphuric iodide, . . . SI6. DISULPHUR DICHLORIDE—SULPHUROUS CHLORIDE. 255 DISULPHUR DICHLORIDE. 'S'2C12. Molecular weight = 135. Molecular volume I I I. 1 litre of disulphur dichloride vapor weighs 67.5 criths. Specific gravity of liguid 168. Boils at 139° C. (282.2° F.). Preparation.—A current of thoroughly dried chlorine is passed over the surface of heated sulphur contained in a retort. The disulphur dichloride distils over as fast as it is formed and collects in the cooled receiver. The process must be interrupted before all the sulphur is converted into the chloride, and the product must be purified by rectifi- cation. S2 + ci2 = 'S'2C12. Disulphur dichloride. Properties.—Disulphur dichloride is an amber-colored, fuming liquid, possessing a disagreeable pungent odor. Its vapor irritates the eyes. It dissolves sulphur freely, a property which is utilized in the manufacture of vulcanized india-rubber. Reaction.—ln contact with water it is gradually decomposed with formation of hydrochloric acid and sulphurous anhydride, whilst sul- phur is deposited : 2/S'2Cl2 + 20H2 = 4HCI + S02 + 3S. Disulphur Water. Hydrochloric Sulphurous dichloride. acid. anhydride. HYPOSULPHUROUS CHLORIDE. SC12. This compound is prepared by saturating disulphur dichloride with chlorine at o°. On removing the excess of chlorine by a stream of dry carbonic anhydride, the hyposul- phurous chloride remains behind as a dark-red liquid. It is very unstable, sponta- neously decomposing at ordinary temperatures into disulphur dichloride and chlorine. On attempting to distil it, this decomposition takes place rapidly. With water it is decomposed like disulphur dichloride. SULPHUROUS CHLORIDE. SC14. Sulphurous chloride is obtained as a yellowish-brown, very mobile liquid by satu- rating disulphur dichloride with chlorine at a temperature of from —2o° to —22° C. ("—4° to —B° F.). It is even less stable than the foregoing compound, and can exist only at temperatures below —2o° C. (—4° F.). When removed from the freezing Mixture it rapidly evolves chlorine, and is converted into hyposulphurous chloride. *v ater decomposes it with violence, forming sulphurous anhvdride and hydrochloric acid; SC14 + 20 H5 = S0s + 4IT Cl. Sulphurous Water. Sulphurous Hydrochloric chloride. anhydride. acid. 256 INORGANIC CHEMISTRY. DISULPHUR DIBROMIDE. /S/2Br2. This compound is formed by the direct union of its elements. It forms a heavy red liquid which distils with partial decomposition between 210° and 220° C. DISULPHUR D INIODIDE, /s/2i2. Disulphur diniodide is obtained as a dark-gray crystalline mass by heating sulphur and iodine together under water. SULPHURIC lODIDE. sr6. This substance is obtained in crystals when a solution of iodine and sulphur in carbonic disulphide is allowed to evaporate. It is interesting as a compound of hexadic sulphur in which all the six bonds are satisfied by monads. COMPOUND OF SULPHUR WITH CARBON. CARBONIC DISULPHIDE, Bisulphide of Carbon. cs2. Molecular weight = 76. Molecular volume I I I. 1 litre of carbonic disulphide vapor weighs 38 criths. Bp. qr. of liquid 1.293. Fuses at —loo° C. (—lBo° F.). Boils at 46.6° C. (115.9° F.). History.—Carbonic disulphide was discovered by Lampadius in 1796. Preparation.—l. This compound is formed by the direct combina- tion of its elements at a high temperature. A tubulated earthenware retort, filled with pieces of charcoal and furnished with a vertical porce- lain tube luted to the tubulure and passing to the bottom of the retort, is heated to redness. Fragments of sulphur are introduced one at a time through the porcelain tube, the latter being closed at the top after each addition. The sulphur volatilizes and its vapor combines with the carbon forming carbonic disulphide, which distils over and is con- densed as a liquid and collected under water: c + s2 = os2. Carbonic disulphide. Sulphuretted hydrogen is formed at the same time owing to the combi- nation of the sulphur with the hydrogen which is invariably present in charcoal. The crude product is redistilled in order to free it from dis- solved sulphur. Thus prepared it possesses a peculiar, fetid odor, due to the presence of other volatile sulphur compounds. These may be CARBONIC DISULPHIDE. 257 removed by shaking the liquid with mercury or corrosive sublimate, subjecting it afterwards to a further distillation. 2. It is also formed when charcoal is heated with iron- or copper- pyrites. This was the method employed by Lampadius. The reaction is due to the sulphur which is given off by the pyrites on heating, and is essentially the same as the foregoing: C + 2FeS2 = CS"2 + 2FeS". Iron-pyrites Carbonic Ferrous (Ferric disulphide), disulphide. sulphide. It is to the occurrence of iron-pyrites in coal that the presence of car- bonic disulphide vapor in coal-gas is due. This impurity, on account of the difficulties attending its removal, has long been the source of annoyance both to the gas manufacturer and the consumer. Properties.—Carbonic disulphide is a colorless, powerfully refracting, mobile liquid. When pure, it possesses a sweetish, ethereal odor. It solidifies at 116° C. ( 177° F.) and fuses at— 110° C. ( 166° F,). It dissolves sulphur, phosphorus, iodine, caoutchouc, oils, and fats. Sulphur and phosphorus may be obtained in crystals by the spontaneous evaporation of their solutions in carbonic disulphide. It is extensively employed in manufacturing processes as a solvent. Carbonic disulphide is exceedingly inflammable. Its vapor inflames in the air at 149° C. (300° F.), and may be ignited by bringing a test tube of paraffin heated to this temperature in contact with it. It burns with a blue flame, yielding carbonic anhydride and sulphurous anhy- dride : CS"2 + 302 = C02 + 2S02. Carbonic Carbonic Sulphurous disulphide. anhydride. anhydride. A mixture of the vapor with air or oxygen explodes with great violence on the approach of a flame. Mixed with nitric oxide and inflamed, the vapor burns, emitting a brilliant blue light, very rich in rays of high refrangibility. Carbonic disulphide is highly poisonous. Its vapor, inhaled in large quantities, proves speedily fatal, and even in minute quantity is very dangerous when habitually inhaled (as, for instance, in factories in which it is employed), owing to a specific action on the nervous system. Reactions.—1. Heated potassium burns in the vapor of carbonic disulphide with formation of potassic sulphide and liberation of carbon : OS" + 2K2 = 2SK2 + C. Carbonic Potassic disulphide. sulphide. 2, When brought into contact with a solution of an alkaline hydrate, carbonic disulphide is decomposed, a carbonate and a sulphocarbonate being formed: s 60KH + 3CS"2 = 20S//Ks2 + COKo2 + 30H2. Potassic Carbonic Potassic Potassic Water, hydrate. disulphide. sulphocarbonate. carbonate. 258 INORGANIC CHEMISTRY. 3. In contact with solutions of alkaline sulphides, carbonic disul- phide also forms alkaline sulphocarbonates : SK2 + CS"2 = CS"Ks2. Potassic Carbonic Potassic sulphide. disulphide, sulphocarbonate. 4. When the vapor of carbonic disulphide is passed over heated calcic hydrate it is decomposed, carbonic anhydride and sulphuretted hydrogen being evolved: CS2 + 2CaHo2 = 2CaO + C02 + 2SH2. This reaction has been successfully employed in removing carbonic disulphide from illuminating gas. Carbonic disulphide is, as has already been pointed out, the sulphur compound corresponding to carbonic anhydride. A carbonic mono- sulphide, corresponding to carbonic oxide, has not been prepared. SULPHOCARBONIC ACID. CS//Hs2. Preparation.—This compound is obtained as a reddish-brown oily liquid by the action of hydrochloric acid on ammonic sulphocarbonate : CS"(NH4S)2 + 2HCI = CS//Hs2 + 2NFT4CI. Ammonic Hydrochloric Sulpho- Amnionic sulphocarbonate. acid. carbonic acid. chloride. COMPOUND OF SULPHUR WITH CARBON AND OXYGEN CARBONIC OXYSULPHIDE. COS". Molecular ineight = 60. Molecular volume I I I. 1 litre of carbonic oxysulphide weighs 30 criths. Gaseous. History.—This gas, which in composition lies intermediate between carbonic anhydride and carbonic disulphide, was discovered by C. von Than. Occurrence.—It appears to exist in solution in the waters of certain mineral springs. Preparation.—1. Carbonic oxysulphide is formed when a mixture of carbonic oxide and sulphur vapor is passed through a heated tube: CO + S = COS". Carbonic Carbonic oxide. oxysulphide. COMPOUNDS OP SULPHUR WITH OXYGEN AND HYDROXYL. 259 2. It is most readily obtained by the action of moderately strong sul- phuric acid upon potassic sulphocyanide: ONKs + 2S02Ho2 + 0H2 = COS" Potassic Sulphuric Water. Carbonic sulphocyanide. acid. oxysulphide. + S02HoKo + S02Ho(NvH40). Hydric potassic Hydric ammonic sulphate. sulphate. By regulating the temperature a steady evolution of the gas is obtained. Propeities.—Carbonic oxysulphide is a colorless gas with a peculiar odor. It is readily inflammable, and forms with oxygen a mixture which explodes on the approach of a flame. It is soluble in its own volume of water, to which it imparts its characteristic odor. Reactions.—1. A platinum wire heated to whiteness by means of the voltaic current decomposes the gas into sulphur and carbonic oxide, the latter occupying the same volume as the carbonic oxysulphide employed. 2. With caustic alkalies it yields a mixture of carbonate and sulphide: COS" + 4KHo = COKo2 + SK2 + 20H2. Carbonic Potassie Potassic Potassic Water, oxysulphide. hydrate. carbonate. sulphide. COMPOUNDS OF SULPHUR WITH OXYGEN AND HYDROXYL. In these compounds the sulphur is either a dyad, a tetrad, or a hexad. Sulphurous anhydride, S02. o=S=o o •I Sulphurous acid, . . SOHo2. H—O—S—O—H. o O . . 11 Sulphuric anhydride, . S03. o=B=o O Sulphuric acid {Hy- 1 SO Ho2. H—O—l—O—H. dnc sulphate), . . j 22 O o o Pyrosulphuric acid f S02Ho j| ]| {Dihydric disul- -< O . H—O—S—O—S—O—H. phate), hS.,OHo |j || o o 260 INORGANIC CHEMISTRY. o o fSO-O) I 1 Persul phuric anhydride, <0 > o=S—O—B=o. (so2-oj II II o o o Thiosulphuric acid IsO.HoHs. H—O—S—S—H. (Jdyposutphurous acta), J 1 Dithionous acid (Hydro- j SOHo tt c q r\ tt sulphurous acid), . . \ SOHo' O O o o ( = so-b>°" + {l^ho- Baric Sulphuric Baric Dithionic dithionate. acid. sulphate. acid. The solution of dithionic acid may be evaporated in vacuo over sulphuric acid till it attains a specific gravity of 1.347, but beyond this point it decomposes into sulphuric acid and sulphurous anhydride. The dilute acid undergoes the same change on boiling. 2. Dithionic acid is also formed when a dilute solution of iodine in potassic iodide is added to a dilute solution of hydric sodic sulphite : 2SOHoNao + I 2 = {lq'ho + 2NaL Hydric sodie Dithionic Sodic sulphite. acid. iodide. About 20 per cent, of the sulphite is thus transformed. The remainder is converted into sulphate. Dithionates.—The dithionates mostly crystallize well. They may be obtained either by neutralizing a solution of the acid with a base, or by exactly precipitating a solu- tion of baric dithionate with a soluble sulphate. TRITHIONIC ACID, Sulpho dithionic Acid, Sulphuretted Hyposulphuric Acid. ( S02Ho \8" . ( S02Ho Preparation.—1. By digesting flowers of sulphur at a temperature of between 50° and 60° Ci with a conceptrated solution of hydric potassic sulphite, potassic trithionate and potassic thiosulphate are formed: f S02Ko 6SOKOHO + 2S = 2\ 8" + S02KoKs + 30H2. i S02Ko Hydric potassic Potassic Potassic Water, sulphite. trithionate. thiosulphate. 2. Potassic trithionate may also be obtained by saturating a very concentrated solu- tion of potassic thiosulphate with sulphurous anhydride; r so2Ko 2S02KoKs + 3S02 = +B. i S02Ko Potassic Sulphurous Potassic thiosulphate. anhydride. trithionate. 3. The same salt is formed when a solution of potassic argentic thiosulphate is boiled: rso2Ko 2S02KoAgs = -j S'7 + SAg2. I S02Ko Potassic argentic Potassic Argentic thiosulphate. trithionate. sulphide. 4. By adding iodine to a solution of sodic sulphite and thiosulphate, sodic trithionate and sodic iodide are formed: 280 INORGANIC CHEMISTRY. f S02Nao SONao2 + S02NaoNas + I 2 = JS" + 2NaI. i S02Nao Sodic Sodie Sodie Sodic sulphite. thiosulphate. trithionate. iodide. An aqueous solution of trithionic acid may be obtained by decomposing the potas- sium salt with hydrofluosilicic acid : f S02Ko f S02Ho \S" + SiF4,2HF = ]S" + SiF„2KF. ( S02Ko I S02Ho Potassic Hydrofluo- Trithionic Potassic trithionate. silicic acid. acid. silicofluoride. The liquid is filtered from the insoluble potassic silicofluoride. The free acid is very unstable, and spontaneously decomposes into sulphuric acid, sulphurous anhydride, and free sulphur: r so2Ho <8" = S02Ho2 + so2 + s. ( S02Ho Trithionic Sulphuric Sulphurous acid. acid. anhydride. Sodium amalgam converts a trithionate into a mixture of sulphite and thiosulphate, thus reversing the process of its formation from these salts: f S02Nao \S" + Na2 = SONao, + S02NaoNas. f SC)2Nao Sodic Sodlc Sodic trithionate. sulphite. thiosulphate. TETRATHIONIC ACID, Disulphodithionic Acid, Bisulphuretted Hyposulphuric Acid. f S02Ho I S" 1 S'' L S02Ho Preparation.—When iodine is added to a solution of a thiosulphate, an iodide and a tetrathionate of the base are formed: f S02Nao 2S02NaoNas +l2= J |" + 2NaI. t S02Nao Sodic Sodic Sodic thiosulphate. tetrathionate. iodide. This action of iodine, in coupling together two atoms of sulphur in two molecules of substances containing the group Hs (or its equivalent, Ks, Nas, etc.) is character- istic of this element, and meets with many applications in organic chemistry. If baric thiosulphate be employed, baric tetrathionate will be formed, and by de- composing this salt with dilute sulphuric acid an aqueous solution of tetrathionic acid may be obtained. The dilute solution may be boiled without decomposition; but, when concentrated, the acid breaks up into sulphurous acid, sulphuric acid, and free sulphur. Sodium amalgam reconverts the tetrathionates into thiosulphates: f S02Nao I S" } S" + Na2 = 2S02NaoNas. [ S02Nao Sodic tetrathionate. Sodic thiosulphate. SULPHUROUS OXYDICHLORIDE. 281 PENTATHIONIC ACID, Trisulphodithionic Acid, Trisulphuretted Hyposulphuric Acid. S02ITo 8" - s" . S" S02Ho Preparation.—1. This acid may be obtained by passing sulphuretted hydrogen into a solution of sulphurous anhydride: ' S02Ho S" 5SH2 + 5S02 = ]S" + 40H2 + SS. | 8" S02Ho Sulphuretted Sulphurous Pentathionic Water, hydrogen. anhydride. acid. 2. It is also formed by the action of disulphur dichloride on baric thiosulphate: fso2_ o IS" | 2S02aßa + = -I 8" Bao" + BaCl2 + S. I 8" | L so,-1 Baric Disulphur Baric Baric thiosulphate. dichloride. pentathionate. chloride. The aqueous solution of the acid may be concentrated till it attains a specific gravity of 1.6, but beyond this point it decomposes, evolving sulphurous anhydride. The pentathionates are unstable, and have been but imperfectly examined. COMPOUNDS OF SULPHUR WITH OXYGEN AND CHLORINE (OXYCHLORIDES, ACID CHLORIDES). These compounds may be regarded as derived from the corresponding oxy-acids of sulphur by the substitution of chlorine for hydroxyl (see acid chlorides of the nitro- gen acids, p. 229). Acid chloride. Corresponding acid. Sulphurous oxychloride (Thionylic chloride), SOCl2 Sulphurous acid, . . . S011o2 Sulphuric oxydichloride (Sulphurylie 1 chloride), Sulphuric oxychlorhydrate (Sulphur- so2ci2 1 Sulphuric acid, .... S02Ho2 ylic chlorhydrate), | S02C1HoJ r S02C1 ( S02Ho Pyrosulphurylic chloride, . , . . Pyrosulphuric acid, . . -| 0 1 :so2ci S02Ho SULPHUROUS OXYDICHLORIDE, Thionylic Chloride. SOC12. Molecular weight = 119. Molecular volume [ | |. 1 litre of sulphurous oxydichloride vapor weighs 59.5 criths. Specific gravity of liquid 1.675. Boils at 78° C. (172.4° F.). Preparation.—l. When dry sulphurous anhydride is passed over phosphoric chlo- ride, sulphurous oxydichloride and phosphoric oxytrichloride are formed: S02 + PC15 = SOC12 + POC13. Sulphurous Phosphoric Sulphurous Phosphoric anhydride. chloride, oxydichloride. oxytrichloride. 282 INORGANIC CHEMISTRY. 2. It may also be obtained by heating together calcic sulphite and phosphoric oxy- trichloride in sealed tubes to 150° C. (302° F.) 3SOCao" + 2POC13 = 3SOC12 + | p^Cao'^. Calcic Phosphoric Sulphurous Calcic sulphite. oxytrichloride. oxydichloride. phosphate. Properties.—Sulphurous oxydichloride is a colorless liquid, possessing a pungent odor. Reaction.—Water gradually decomposes sulphurous oxydichloride into sulphurous and hydrochloric acids: SOCl2 + 20H2 = SOHo2 + 2HCI. Sulphurous Water. Sulphurous Hydrochloric oxychloride. acid. acid. SULPHURIC OXYDICHLORIDE, Sulphurylic Chloride. S02C12. Molecular weight 135. Molecular volume | | |. 1 litre of sulphuric oxydichloride vapor weighs 67.5 criths. Specific gravity of liquid 1.66. Boils at 70° C. (158° F.). Preparation.—1. Sulphuric oxydichloride is formed by the direct union of sulphur- ous anhydride and chlorine, either in sunlight or when the two gases are passed into glacial acetic acid or through camphor which immediately liquefies, and the saturated solution, after standing for some time, subjected to distillation: S02 + Cl2 = S02C12. Sulphurous Sulphuric anhydride. oxydichloride. 2. It may also be prepared by heating sulphuric oxychlorhydrate (see below) in sealed tubes for 12 hours to a temperature of from 170° to 180° C. (338°-356° F.). 2S02CIHo = S02C12 -f S0211o2. Sulphuric Sulphuric Sulphuric oxychlorhydrate. oxydichloride. acid. Properties.—Sulphuric oxydichloride is a colorless fuming liquid with a suffocating odor. Reactions.—1. A small quantity of water decomposes it into sulphuric oxychlorhy- drate and hydrochloric acid : S02C12 + OH2 = S02CIHo + HCI. Sulphuric Water. Sulphuric Hydrochloric oxydichloride. oxychlorhydrate. acid. 2. An excess of water converts it into sulphuric and hydrochloric acids: S02C12 + 20H2 = S02Ho2 + 2HCI. Sulphuric Water. Sulphuric Hydrochloric oxydi chloride. acid. acid. SULPHURIC OXYCHLORHYDRATE, Sulphurylic Chlorhydrate. SO2C]Ho. Molecular weight = 116.5. Molecular volume f I I- 1 litre of sulphuric oxychlorhydrate vapor weighs 58.25 criths. Specific gravity of liquid 1.776 at 18° C. (64.4° F.). Boils at 158° C, (316.4° F.). Preparation.—1. Sulphuric anhydride and hydrochloric acid unite directly to form sulphuric oxychlorhydrate: 283 SELENIUM. S03 + HCI = S02CIHo. Sulphuric Hydrochloric Sulphuric anhydride. acid. oxychlorhydrate. 2. It may be obtained by distilling a mixture of sulphuric acid and phosphoric chloride. 3502H02 + PC15 = 3S02CIHo + PO2HO + 2HCI. Sulphuric Phosphoric Sulphuric Metaphosphoric Hydrochloric acid. chloride. oxychlorhydrate. acid. acid. Properties.—Sulphuric oxychlorhydrate is a colorless, strongly fuming liquid. When distilled it undergoes partial dissociation into sulphuric anhydride and hydrochloric acid. Reaction.—Water decomposes it with violence, forming sulphuric and hydrochloric acids: SChCIHo + OH2 = S02Ho2 + HCI. Sulphuric Water. Sulphuric Hydrochloric oxychlorhydrate. acid. acid. PYROSULPHURYLIC CHLORIDE. rso2ci \o . ( S02CI Molecular weight = 215. Molecular volume | | |. 1 litre of pyrosulphurylic chloride vapor weighs 107-.5 criths. Specific gravity of liquid 1.819 at 18° C. (64.4° F.). Boils at 146° C. (294.8° F.). Preparation.—l. This compound is formed when sulphuric anhydride is heated with phosphoric chloride; f SO2CI 2S03 + PC15 = ] O + POCl3. ( S02CI Sulphuric Phosphoric Pyrosulphurylic Phosphoric anhydride. chloride. chloride. oxytrichloride. 2. It is also produced by the action of disulphur dichloride on sulphuric anhydride; r so2ci 'S'.CI, + 5S03 = \ O + 5S02. [ S02CI Disulphur Sulphuric Pyrosulphurylic Sulphurous dichloride. anhydride. chloride. anhydride. Properties.—Pyrosulphurylic chloride is a heavy, colorless, fuming liquid. Reaction.—In contact with water it is slowly decomposed into sulphuric and hydro- chloric acids: r SO2CI 1 O + 30H2 . = 2S02H02 + 2HCI. ( S02CI Pyrosulphurylic Water. Sulphuric Hodrochloric chloride. acid. acid. SELENIUM, Se2. Atomic weight 79. Molecular weight 158. Molecular volume □Q 1 litre of selenium vapor weighs 79 criths. Sp. gr., amorphous, 4.28; crystallized, 4.8. Fuses at 217° C. (422.6° F.). Boils about 700° C. (1292° F.). Atomicity ", IT, and Tl. Evidence of atomicity: Seleniuretted hydrogen, . . Se"H2. Selenious chloride, .... . . SeiTCl4. Selenic acid, . . Sevi02Ho2. History.—Selenium (from the moon) was discovered in 1817 by Berzelius in a deposit from a sulphuric acid chamber. The name 284 INORGANIC CHEMISTRY. was given on account of the analogy of this element with tellurium (tellus, the earth). Occurrence.—Selenium is generally found in very small quantities along with sulphur, both native and combined. Less frequently it occurs alone in combination with metals in a few rare minerals, as the selenides of lead, copper, silver, and mercury. When iron- or copper-pyrites containing selenium is employed in the manufacture of sulphuric acid, the selenium forms a red deposit in the chambers. Preparation.—The red deposit from the sulphuric acid chambers is digested with a warm solution of potassic cyanide until the red color disappears. Soluble potassic selenocyanide is formed : KCy + Se = SeKCy. Potassic Potassic cyanide. selenocyanide. cyanide. On adding an excess of hydrochloric acid to the filtered solution, sele- nium is precipitated as a red amorphous powder, the liberated seleno- cyanic acid being instantly decomposed in presence of strong acids into hydrocyanic acid, which remains in solution, and selenium. Properties.—Selenium, like sulphur, exists in various modifications. When precipitated from solutions by means of acids, it forms an amor- phous brick-red powder, which, when heated along with the liquid, turns black and cakes together below 100° C. When melted and rapidly cooled, selenium solidifies to a black, shining, amorphous mass, with a conchoidal fracture. This variety is soluble in carbonic disulphide, and possesses a specific gravity of 4.28. The solution deposits mono- clinic crystals, isomorphous with those of monoclinic sulphur. The fusing point of soluble selenium cannot be determined, as this substance softens gradually on heating. When amorphous selenium is heated for some time to a temperature of 97° C. (206,6° F.), it is converted into the crystalline modification. This change is attended with evolution of great heat, the temperature of the mass rising above 200° C. Crystalline selenium is of a dark gray color, with a metallic lustre and granular fracture. Its specific gravity is 4.5. The same variety is obtained when melted selenium is allowed to cool very slowly. It is insoluble in carbonic disulphide. This modification conducts the electric current. Its electrical resistance is greatly diminished by exposing the substance to light, but is again restored on shading it from the light—a property which is turned to account in the construction of the photophone. When a solution of an alkaline selenide is exposed to the air, minute black crystals of selenium separate out, possessing a specific gravity of 4.8, They are insoluble in carbonic disulphide. The vapor-density of selenium, like that of sulphur, decreases as the temperature rises. Above 1400° C. (2552° F.) it possesses the normal vapor-density corresponding with the molecular weight Se2 158. The following determinations of the vapor-density (air = 1) illustrate this decrease: SELENIURETTED HYDROGEN, 285 Temperature. Vapor-density. 860° C. (1580° F.) 7.67 1040° “ (1804° “) 6.37 1420° “ (2588° “) 5.68 Selenium dissolves in fuming sulphuric acid, with a green color. Reaction.—When heated in the air selenium burns, forming selenious anhydride, Se02, at the same time giving off an odor of decayed horse- radish. Nitric acid oxidizes selenium to selenious acid, SeOHo2, whereas sul- phur under the same conditions yields sulphuric acid. COMPOUNDS OF SELENIUM WITH HYDROGEN AND CHLORINE. SELENIURETTED HYDROGEN, Hydroselenic Add. SeH2. Molecular weight =Bl. Molecular volume I I I. 1 litre weighs 40.5 criths. Preparation.—This compound is formed by the action of dilute hy- drochloric acid upon ferrous selenide : FeSe" + 2HCI = SeH2 + FeCl2. Ferrous Hydrochloric Seleniuretted Ferrous selenide. acid. hydrogen. chloride. Properties.—Seleniuretted hydrogen is a colorless gas, possessing an odor resembling that of sulphuretted hydrogen, but much more power- ful. Inhalation of a single bubble of seleniuretted hydrogen through the nose destroys for some time the sense of smell. Like sulphuretted hydrogen it produces precipitates in solutions of most of the heavy metals. It is decomposed by heat into its elements. The degree of this dissociation varies in a remarkable manner, being less at a higher than at a lower temperature. Thus the dissociation begins at 150° C. (302° F.), increases gradually up to 270° C. (518° F.), then decreases gradually as the temperature rises, till at 520° C. (968° F.) it almost entirely ceases. At still higher temperatures it again increases. When ignited, seleniuretted hydrogen burns in air with a blue flame, yielding selenious anhydride and water: SeH2 + 30 = Se02 + 0H2. Seleniuretted Selenious Water, hydrogen. anhydride. There are two chlorides of selenium, and SeCl4. 286 INORGANIC CHEMISTRY. COMPOUNDS OF SELENIUM WITH OXYGEN AND HYDROXYL. Selenious anhydride, Se02. Selenious acid, SeOHo2. Selenic acid, Se02Ho2. SELENIOUS ANHYDRIDE Se02. Preparation.—Selenious anhydride is formed by the direct com- bination of its elements, when selenium is burned in a stream of oxygen. It may also be obtained by heating selenious acid: Se()Ho2 = Se02 + OIL, Selenious Selenious Water, acid. anhydride. Properties.—Selenious anhydride crystallizes in prisms, and when heated sublimes without fusing. Reaction.—Dissolved in water it forms selenious acid by a reaction the reverse of the foregoing. SELENIOUS ACID. SeOHo2. Preparation.—l. As above, by dissolving selenious anhydride in iter. 2. It is formed when selenium is oxidized with nitric acid : water. Se + 02 + OH2 = SeOHo2. Water. Selenious acid. Properties.—Selenious acid is a white, very soluble substance, crystallizing in prisms. It forms normal, acid, and superacid salts : Normal potassic selenite, .... SeOKo2. Hydric potassic selenite, .. .. SeOHoKo. Superacid potassic selenite, . . . SeOHoKo,SeOHo2. Reaction.—Reducing agents, such as sulphurous acid, stannous chlo- ride, etc., precipitate red amorphous selenium from its solutions : SeOIIo2 + 2SOHo2 = Se + 2S02Ho2 + OH2. Selenious acid. Sulphurous acid. Sulphuric acid. Water. SELENIC ACID—TELLURIUM. 287 SELENIC ACID. Se02Ho2. Preparation.—l. The most convenient method of obtaining this acid consists in suspending argentic selenite in water, and adding bromine until a perceptible reddish coloration is visible: SeOAgo2 -f- Br2 -}- 0H2 = Se02Ho2 + 2Agßr. Argentic selenite. Water. Selenic acid. Argentic bromide. On evaporating the filtered liquid a concentrated solution of selenic acid remains. 2. Potassic seleniate is prepared by fusing selenium or metallic selen- ides with nitre. The potassic salt thus formed is then converted into a plumbic salt, and, by decomposing the latter with sulphuretted hydro- gen, selenic acid is obtained. Properties.—Selenic acid is not known in a state of purity. The most concentrated aqueous solution contains 97.4 per cent, of the acid. Further evaporation causes it to decompose into selenious anhydride, oxygen, and water. The solution has a specific gravity of 2.627, and closely resembles in its properties concentrated sulphuric acid. It is remarkable as being the only single acid which dissolves gold. In this process it undergoes reduction to selenious acid. Reaction.—When heated with hydrochloric acid, selenic acid is re- duced to selenious acid, chlorine being liberated : Se02Ho2 + 2HCI = SeOHo2 + 0H2 + Cl2. Selenic Hydrochloric Selenious Water, acid. acid. acid. Selenic anhydride has not been prepared TELLURIUM, Te2. Atomic weight = 125. Molecular weight = 250. Molecular volume CD. 1 litre of tellurium vapor weighs 125 criths. Sp. gr. 6.2. Fuses at 490°-500° C. (914°-932° F.). Atomicity ",iT, and yi. Evidence of atomicity: Telluretted hydrogen, , . Tellurous chloride, . . . . TeivCl4. Telluric acid, . . . . TeTiG2Ho2, History.—Tellurium (from tellus, the earth) was first recognized as a distinct substance by Muller von Reichenstein, in 1782. Occurrence.—It is found in very small quantities both in the native state and as the tellurides of metals. Preparation.—Bisrauthic telluride, Bi2Te"3, a substance occurring in 288 INORGANIC CHEMISTRY. nature as the mineral tetradymite, is fused with a mixture of sodic car- bonate and finely-powdered charcoal. The fused mass yields on lixivi- ation with water, a solution of sodic telluride, which on exposure to the air, deposits tellurium as a gray powder. The pulverulent tellurium may be fused into a coherent mass under sodic chloride. Properties.—Tellurium is a silver-white crystalline substance with a metallic lustre. At a high temperature it may be distilled. It dis- solves in fuming sulphuric acid with a deep red color. Reaction.—When heated in air it burns with a blue flame, forming tellurous anhydride, Te02. COMPOUNDS OF TELLURIUM WITH HYDROGEN\ CHLORINE, AND OXYGEN. TELLURETTED HYDROGEN. TelT2. Molecular weight = 127. Molecular volume I I I. 1 litre weighs 63.5 criths. Preparation.—Telluretted hydrogen is obtained by the action of dilute hydrochloric acid on ferrous or zincic telluride: ZnTe" + 2HCI = TeH2 + ZnCl2. Zincic Hydrochloric Telluretted Zincic telluride. acid. hydrogen. chloride Properties.—Telluretted hydrogen is a colorless gas of a fetid odor, resembling that of sulphuretted hydrogen. It exhibits the same anom- alies of dissociation as seleniuretted hydrogen. It may be ignited in air, and burns with a blue dame, forming tellurous anhydride and water: TeH2 + 30 = Te02 + OIL. Telluretted Tellurous Water, hydrogen. anhydride. There are two chlorides of tellurium, 'Te'2Cl2and TeCl4. Tellurous Anhydride, Te02 —This compound is prepared like seleni- ous anhydride (p 286), which it closely resembles in properties. Tellurous Acid.—TeOHo2.—This acid is obtained as a white floccu- lent precipitate, when a solution of tellurium in dilute nitric acid is poured into water. It is decomposed at a temperature of 40° C. (104° F.) into anhydride and water. It dissolves more readily in hydrochloric acid than in water. Sulphurous acid precipitates tellurium from the solution (see Selenious Acid, p. 286). Tellurous acid is a dibasic acid, forming acid and normal salts. Thus: Hydric potassic tellurite, TeOHoKo. Normal potassic tellurite, TeOKo2. TELLURIC ACID. 289 Tetratellurites, produced by the combination of the normal tellurites with tellurous anhydride, are also known : o o o o Di potassic tetratel- lurite, K—O—Te—O—Te—o—Te—O—Te—O—K. Telluric Anhydride, TeOs.—Telluric anhydride is prepared by care- fully heating telluric acid. It forms an orange-yellow mass. When strongly heated it is decomposed into tellurous anhydride and oxygen. It is insoluble in water. Boiling concentrated hydrochloric acid dis- solves it slowly, converting it, with evolution of chlorine, into tellurous anhydride: Te03 + 2HCI = Te02 + 0H2 + Cl2. Telluric Hydrochloric Tellurous Water, anhydride. acid. anhydride. Telluric Acid, Te02Ho2.—ln order to prepare this compound tellu- rium is fused with a mixture of caustic potash and potassic chlorate. The tellurium is oxidized at the expense of the oxygen of the potassic chlorate to telluric anhydride, which combines with the alkaline base to form potassic tellurate: Te + {OKO = Teo» + KCI- - Telluric Potassic chlorate. anhydride. chloride. TeOa + 2KHo = Te02Ko2 + 0H2. Telluric Potassic Potassic Water, anhydride. hydrate. tellurate. The fused mass is dissolved in water, and a solution of baric chloride is added, when insoluble baric tellurate is precipitated: TeQ2Ko2 + BaCl2 = Te02Bao" + 2KCI. Potassic Baric Baric Potassic tellurate. chloride. tellurate. chloride. The baric tellurate is suspended in water, and decomposed with the exact quantity of sulphuric acid. In this way insoluble baric sulphate and free telluric acid are formed. On evaporating the filtered solution, large colorless monoclinic crystals of hexabasic telluric acid, TeHo6, are deposited. On heating to 160° these crystals part with two molecules of water, yielding dibasic telluric acid, Te02Ho2, as a white amor- phous mass. Telluric acid forms a series of somewhat complex salts. Among the potassium salts, for example, tellurates, ditellurates, and tetratellu rates are known. 290 INORGANIC CHEMISTRY. Tetrahydric dipotassic tellurate, .. . TeHo4Ko2,30H2. ( TeHo4Ko“ Octohydric dipotassic ditellurate, . . < O I TeHo4Ko f TeHo4Ko O Te°2 Octohydric dipotassic tetratellurate,. ■{ O Te02 O TeHo4Ko CHAPTER XXYIII. MONAD ELEMENTS. Section 11. (Continued from Chapter XXII.) BROMINE, Br2. Atomic weight = 80. Molecular weight = 160. Molecular volume I I I. 1 litre of bromine vapor weighs 80 criths. Sp. gr. 3.187. Fuses at— 24.5° C, (—12.1° F.). ' Boils at 63° C. (145.4° F.). Atomicity'. Evidence of atomicity: Hydrobromic acid, . . . . HBr. Potassic bromide, . . . . . . . KBr. Argentic bromide, .... History.—Bromine (from ftpa/ws, a stench) was discovered in 1826, by Balard, in the mother-liquors obtained in the crystallization of common salt from sea-water. Occurrence.—Bromine does not occur in the free state in nature. It is found in combination with metals as bromides, sodic bromide being the most common. This salt occurs in small quantity in sea-water, particularly in the water of the Dead Sea, and in greater abundance in many salt springs and deposits of rock salt. The salt mines of Stass- furt furnish 20,000 kilos, of bromine yearly. Preparation.—l. The mother-liquors of saline waters containing bromides are treated with chlorine as long as the color of the liquid continues to become darker. In this way bromine is liberated, and may be distilled otf and collected in a cooled receiver: 2Naßr + Cl2 = 2NaCI + Br2. Sodic bromide. Sodic chloride. An excess of chlorine must be avoided, as this would occasion the for- mation of a chloride of bromine. BROMINE. On a large scale the mother liquors are mixed with an excess of sul- phuric or hydrochloric acid, and a quantity of manganic peroxide exactly sufficient to liberate the bromine present (see Equation, Prepa- ration 2) is added. As long as an excess of the peroxide is avoided, there is no danger of obtaining a product contaminated with chlorine, since any chlorine which might be liberated would at once set free its equivalent of bromine. 2. Bromine may also be obtained from pure bromides, in a reaction similar to that employed in the preparation of chlorine, by heating them with sulphuric acid and manganic peroxide: 2Naßr -f Mn02 + 2S02Ho2 = Br2 + Sodic Manganic Sulphuric bromide. peroxide. acid. S02Nao2 -f~ S02Mno" + 20H2 Sodic sulphate. Manganous Water, sulphate. Properties.—Bromine is a heavy reddish-brown liquid, transparent only in thin layers. Its vapor possesses a considerable tension at ordi- nary temperatures. If a few drops be poured into a flask, the latter will be speedily filled with the reddish-brown vapor. At a tempera- ture of —24.5° C. (—12.1° F.) bromine solidifies to a crystalline mass with a slight metallic lustre. Bromine has a powerful and un- pleasant odor, resembling that of chlorine. Its vapor attacks the eyes and the organs of respiration. It is an irritant poison. When brought in contact with the skin, it produces dangerous wounds. Throughout a considerable range of temperature above its boiling point, bromine has a vapor-density corresponding with the molecular formula Br2. At higher temperatures the vapor-density diminishes, owing to a partial dissociation of the molecules of the vapor into single atoms. This dissociation, which occurs more readily than in the case of chlorine, but less readily than in the case of iodine, is not complete at 1600° C. (2912° F.), the highest temperature that has been em- ployed in such determinations. Bromine is soluble in about thirty times its weight of water at ordi- nary temperatures, the solubility decreasing as the temperature rises. The solution is of a reddish color, and, when exposed to a temperature of 0° C. deposits crystals of a hydrate, Br2,100H2, melting at 15° C. (59° F.). Bromine is more soluble in ether and carbonic disulphide than in water, and when an aqueous solution is agitated with either of these solvents, the bromine is extracted from the water and passes into the new solvent, which separates from the water as a dark-colored layer, on allowing the liquid to stand. Moist bromine bleaches vegetable colors, but less powerfully than chlorine. Bromine combines directly with many of the metals to form bro- mides. Antimony and tin inflame spontaneously in the vapor, and burn with great brilliancy. Potassium and bromine, when brought together at ordinary temperatures, unite, frequently with explosion; 292 INORGANIC CHEMISTRY. but sodium must be heated to 200° C. in contact with bromine vapor, before any action occurs. HYDROBROMIC ACID. HBr. Molecular weight = 81. Molecular volume I I I. 1 litre of hydrobromic acid weighs 40.5 criths. Fuses at —73° C. (—99.4° F.). Boils at —696 C. (—92.2° F.). Preparation.—l. When a mixture of hydrogen and bromine vapor is passed through a red-hot tube, or when a mixture of hydrogen and bromine vapor is burned in air, hydrobromic acid is formed by the direct combination of its elements : H2 + Br2 = 2HBr. Hydrobromic acid. 2. It may be obtained by heating potassic bromide with phosphoric acid: 3KBr + POHo3 = POKO3 + 3HBr. Potassic Phosphoric Potassic Hydrobromic bromide. acid. phosphate. acid. Sulphuric acid cannot be substituted for phosphoric acid in this re- action, as a portion of the hydrobromic acid is then decomposed, with liberation of bromine: S02Ho2 + 2HBr = Br2 + 20H2 + S02. Sulphuric Hydrobromic Water, Sulphurous acid. acid. anhydride. 3. It is formed by the action of water upon phosphorous tribromide or phosphoric pentabromide; P'"Br3 + 30H2 = PoH3 + 3HBr. Phosphorous Water. Phosphorous Hydrobromic tribromide. acid. acid. PBr5 + 40H2 = POHo3 + SHBr. Phosphoric Water. Phosphoric Hydrobromic pentabromide. acid. acid. These reactions may be most conveniently applied by gradually dropping the requisite quantity of bromine into water containing amor- phous phosphorus. The bromides of phosphorus are decomposed at the moment of their formation : P + Br5 -f 40 H2 = POHO3 + SHBr. Water. Phosphoric Hydrobromic acid. acid. COMPOUNDS OF BROMINE WITH OXYGEN AND HYDROXYL. 293 This is the method most usually employed in the laboratory for the preparation of hydrobromic acid. 4. It may also be obtained in aqueous solution by passing sulphur- etted hydrogen through water containing bromine: SH2 + Br2 = 2HBr + S Sulphuretted Hydrobromic hydrogen. acid. Properties.—Hydrobromic acid is a colorless gas, with a pungent odor. It fumes strongly in contact with moist air. By means of pressure and cold it may be liquefied, and when cooled to —73° C. (—99.4° F.) solidifies to a colorless crystalline mass. Water absorbs more than its own weight of the gas, yielding a powerfully acid liquid. When a solution, saturated at a low temperature, is subjected to distil- lation, the liquid in the retort gradually becomes weaker, until it con- tains 48 per cent, of hydrobromic acid, when it distils unchanged between 125° and 126° C. (257°-259° F.), and possesses a specific gravity of 1.49 at 14° C. (57° F.). When an acid containing less than 48 per cent, is distilled, the liquid in the retort gradually becomes more concentrated till the above percentage is attained. This aqueous solu- tion does not correspond with any definite hydrate, and its composition may be altered by altering the pressure under which the distillation takes place. Reactions.—l. Chlorine decomposes the acid with liberation of bromine : 2HBr + Cl2 = 2HCI -f Br2. Hydrobromic acid. Hydrochloric acid. 2. By the action of atmospheric oxygen a small quantity of bromine is liberated from hydrobromic acid in aqueous solution, but the decom- position is soon arrested; 4HBr -j- 02 = 20H2 -f 2Br2. Hydrobromic acid. Water. 3. In contact with metallic oxides and hydrates bromides are formed. Argentic bromide, Agßr, and mercurous bromide, /Hg’/2Br2, are insol- uble in water; plumbic bromide, Pbßr2, is sparingly soluble; all the other bromides dissolve readily. COMPOUNDS OF BROMINE WITH OXYGEN AND HYDROXYL. Hypobromous acid, . . . . OBrH. f OBr Bromic acid, <0 . (OH The graphic formulae of these compounds are analogous to those of the corresponding chlorine compounds, given on page 177. 294 INORGANIC CHEMISTRY. HYPOBROMOUS ACID. OBrH. Preparation.—An aqueous solution of this very unstable compound may be obtained by agitating mercuric oxide with bromine-water: f Hgßr 2HgO + OH2 + 2Br2 = 20BrH + { O (Hgßr Mercuric Water. Hypobromous Mercuric oxide. acid. oxybromide. The corresponding anhydride, OBr2, has not been prepared. BROMIC ACID. /OBr \OHo- Preparation.—Bromic acid is best prepared by decomposing a solu- tion of baric bromate with the requisite quantity of sulphuric acid : f OBr | O rAP ■{ Bao" + S02Ho2 =2 -J + S02Bao". I OBr Baric Sulphuric Bromic Baric bromate. acid. acid. sulphate. The aqueous solution may be concentrated in vacuo till it contains 1 molecule of acid to 7of water. Beyond this point it is decomposed into water, bromine, and oxygen. The same decomposition takes place when the dilute solution is boiled : 4|°®ro = 2Br2 + 20H2 + 502. Bromic acid. Water. Bromic acid closely resembles chloric acid in its properties. Preparation of Bromates.—1. When bromine is added to a solution of a metallic hydrate, a mixture of bromate and bromide is formed : 6KHo + 3Br2 = SKBr + | + 30H2. Potassic Potassic Potassic Water, hydrate. bromide. bromate. The potassic bromate is much less soluble than the bromide, and may be separated from it by crystallization. lODINE. 295 2. Potassic bromate is also formed when bromine is added to a solu- tion of potassic hydrate or carbonate, and chlorine is passed into the liquid : 6KHO +Br + SCI = SKCI + | + 30H2. Potassic Potassic Potassic Water, hydrate. chloride. bromate. In this way the whole of the bromine is converted into bromate. Character of the Bromates.—Some of the bromates, when heated, lose oxygen, and are transformed into bromides: 2{oKo = 2KBr + 30- Potassic Potassic brotnate. bromide. Others evolve bromine and a portion of their oxygen, leaving metallic oxides : f OBr "h" - + + OBr Magnesic Magnesic bromate. oxide. lODINE, 12.I2. Atomic weight = 127. Molecular weight 254. Molecular volume I I I. 1 litre of iodine vapor weighs 127 criths. Bp. gr. 4.95. Fuses at 114° C. (237° F.), Boils above 200° C. (392° F.). Atomicity'A Evidence of atomicity: Hydriodic acid, HI. Potassic iodide, KI. Argentic iodide, Agl History.—lodine was discovered in 1812 by Conrtois in the mother- liquors of soda prepared from the ashes of sea-weed. The first thorough investigation of its properties is due to Gay-Lussac. Occurrence.—lodine is always found in combination with metals, generally associated with chlorine. In this form it occurs in small quantities in some mineral springs and in sea-water, from which last it is absorbed in larger quantity by various kinds of sea-weed. From these the iodine of commerce is obtained. It has also been detected in some marine animals, such as sponges and oysters. The iodides of silver and lead occur as rare minerals. * See, however, Periodates. 296 INORGANIC CHEMISTRY. Manufacture.—Sea-weed is burned in pits, the temperature being kept as low as possible in order to prevent loss from volatilization of the salts of iodine. The ash thus obtained is known as help. The soluble salts, consisting of alkaline carbonates, sulphates, chlorides, bromides, and iodides, are extracted from the ash with water. The so- lution is evaporated, and the carbonates, sulphates, and chlorides are removed by crystallization. To the mother-liquor, containing the Fig. 43. bromides and iodides, sulphuric acid is added, which causes a separa- tion of sulphur, owing to the presence of sulphides and sulphites. The sulphur and crystals of sulphate are removed, and the liquid is trans- ferred to a large iron retort A (Fig. 43), lined with lead. Heat is ap- plied and manganic peroxide is added in small portions at a time. lodine is thus liberated according to the equation : 2NaI + Mn02 + 2S02Ho2 = S02Nao2 Sodic Manganic Sulphuric Sodic iodide. oxide. acid. sulphate. + S02Mno" -f I 2 + 20H2. Manganous sulphate. Water. and, distilling over, is condensed in a series of stoneware receivers, hhh, fitting one into the other as in the figure. When the iodine ceases to distil over, the receiver is changed, and more manganic peroxide is added. This liberates the bromine, which, on account of its superior affinity for hydrogen and bases, is given off later than the iodine (see equation, p. 297). The bromine is distilled off and collected. Sometimes the dried sea-weed is carbonized in retorts and the result- ing charcoal lixiviated with water. In this way the loss of iodine by volatilization is avoided ; but, on the other hand, it is found impossi- ble to extract the whole of the iodine salts from the charcoal. Properties.—lodine forms bluish-black tabular rhombic crystals, with a metallic lustre. It possesses a peculiar and irritating odor, dis- tantly resembling that of chlorine. When heated, it gives off a vapor lODINE. 297 of a magnificent violet color (hence the name of this element, from i()£iSyj<;, violet-colored). At higher temperatures and when free from admixture of air, this vapor assumes a deep blue tint. The vapor pos- sesses a characteristic absorption-spectrum. The vapor-density of iodine at temperatures up to 700° C. (1292° F.) corresponds with the molecular formula 12.I2. Above this tempera- ture the vapor-density diminishes as the temperature rises, till at 1400° C. (25.52° F.) it is somewhat less than two-thirds of the vapor-density below 700° C. This diminution is due to a partial dissociation of the molecules of iodine into free atoms. If the iodine vapor be mixed with four-fifths of its volume of air, in order to reduce the pressure of the iodine vapor and thus increase the dissociation, the vapor-density of the iodine at 1400° C. is only half as great as at 700° C.; that is to say, the vapor-density corresponds with the molecular formula I, and the iodine vapor at that temperature is mon-atomic. At temperatures above 1400° C. no further diminution occurs under these circumstances. lodine is very slightly soluble in water, but dissolves readily in pres- ence of hydriodic acid or of soluble iodides. Alcohol dissolves it more freely, whilst in ether, chloroform, and carbonic disulphide, it is very readily soluble. The aqueous, ethereal, and alcoholic solutions are brown ; those in chloroform and carbonic disulphide are violet. The smallest trace of free iodine imparts to starch paste a splendid blue color, which disappears on heating, but returns, although with diminished intensity, on subsequent cooling. jßeactions.—1. lodine is expelled by chlorine and bromine from all its compounds with electro-positive elements: 2KI + Cl2 = 2KCI + 12.I2. Potassic iodide. Potassic chloride. 2KI + Br2 = 2KBr + 12.I2. Potassic iodide. Potassic bromide. 2, With a solution of calcic hydrate, iodine yields a liquid which bleaches in alkaline solution, and therefore probably contains calcic iodohypiodite : CaHo2 + I 2 = Ca(OI)I + OII2. Calcic hydrate. Calcic Water. iodohypiodite. The bleaching power diminishes gradually on standing, and more rap- idly on boiling or by exposure to light. When the bleaching property has disappeared, the solution contains only a mixture of calcic iodate and calcic iodide. 3. lodine unites directly with metals and non-metals, the process of combination being frequently accompanied with evolution of heat and light. Phosphorus ignites when brought into contact with solid iodine, and powdered antimony, when thrown into iodine vapor, bursts into flame. 298 INORGANIC CHEMISTRY. HYDRIODIC ACID. Molecular weight = 128. Molecular volume I I I. 1 litre of hy dr iodic acid weighs 64 criths. Fuses at —ss° C. (—67° F.). Preparation.—l, Hydriodic acid is formed by the direct union of its elements when a mixture of iodine vapor and hydrogen is passed through a red-hot tube or over spongy platinum gently heated: H2 + I 2 = 2HT. 2. It is formed when an iodide is heated with phosphoric acid: SKI + POHO3 = SHI -f POKO3. Potassic Phosphoric Hydriodic Potassic iodide. acid. acid. phosphate. Sulphuric acid cannot be substituted for phosphoric acid in this reac- tion, as it liberates iodine from hydriodic acid : 2HI + S02Ho2 = I 2 + S02 + 20H2. Hydriodic Sulphuric Sulphurous Water, acid. acid. anhydride. An aqueous solution of hyriodic acid may however be prepared by de- composing a solution of baric iodide with the exact quantity of dilute sulphuric acid, the sulphuric acid being in this case immediately with- drawn from the reaction in the form of insoluble baric sulphate. 3. It is also formed by decomposing phosphorous triiodide by water : PI3 + 30 H2 = PHo3 + SHI. Phosphorous Water. Phosphorous Hydriodic triiodide. acid. acid. 4. It may be prepared by heating together water, potassic iodide, iodine, and amorphous phosphorus : 4KI + P2 + 5I2 + 80H2 = 14HI + 2POHoKo2. Potassic Water. Plydriodic Hydric dipotassic iodide. acid. phosphate. An aqueous solution of hydriodic acid prepared by Method 5 (see below) may be advantageously substituted for the solution of potassic iodide in the above reaction. The amorphous phosphorus is placed in a retort with the neck sloped slightly upwards, and a solution of 2 parts of iodine in 1 part of aqueous hydriodic acid (b. p. 127°) is allowed to drop gradually through the tubulure frora-a stoppered funnel. Gaseous hydriodic acid is evolved in a steady stream. When the action begins to slacken, a gentle heat may be applied. HYDRIODIC ACID. 299 5. A solution of hydriodic acid may be readily obtained by passing sulphuretted hydrogen through water in which iodine is suspended: 2SH2 + 2I2 = 4HI + S2. Sulphuretted Hydriodic hydrogen. acid. As the reaction proceeds the unattacked iodine dissolves in the aqueous hydriodic acid and facilitates the decomposition. Properties.—Hydriodic acid is a colorless gas, fuming in contact with moist air, and possessing a pungent odor. At a temperature of 0° C. and under a pressure of 4 atmospheres, it condenses to a colorless liquid which solidifies at —ss° C. (—67° F.). It is readily decomposed by heat. A hot glass rod plunged into a vessel filled with the gas, causes the immediate separation of violet vapors of iodine. It is readily absorbed by water, forming a strongly acid liquid. A solution saturated at 0° C. has a sp. gr. of 2. Aqueous hydriodic acid behaves on distillation like hydrochloric and hydrobromic acids (q.v.). The strongest acid obtainable by distillation has a sp. gr. of 1.67, con- tains 57.7 per cent, of hydriodic acid, and boils at 127° C. (260.6° F.). When a weaker or a stronger acid is distilled, the composition of the distillate gradually becomes stronger or weaker until an acid of the above strength and boiling-point distils over unchanged. This acid does not correspond with any definite hydrate and, as in the case of hydrochloric and hydrobromic acids, the composition of the distillate may be made to vary by varying the pressure under which distillation takes place. The aqueous solution when pure is colorless, but in contact with the oxygen of the air, it rapidly becomes brown from separation of iodine, which dissolves in the aqueous acid : 4HI + 02 = 20H2 + 2T2. Hydriodic acid. Water. Oxidizing agents have a similar action. Owing to this property of readily parting with its hydrogen, hydriodic acid is frequently em- ployed as a reducing agent, particularly at higher temperatures and in the case of organic substances. jßeactions.—l. Chlorine and bromine decompose hydriodic acid, liberating iodine: 2HI + Cl2 = 2HCI + 12.I2. Hydriodic acid. Hydrochloric acid. 2. Mercury rapidly decomposes it, with liberation of hydrogen: 2HI + 2Hg = 'Hg'jl, + IL. Hydriodic acid. Mercurous iodide. 3. With metallic oxides, hydrates, and some salts, it forms iodides. 300 INORGANIC CHEMISTRY. Even argentic chloride is transformed by hydriodic acid into argentic iodide: AgCl + HI = Agl + HCI. Argentic Hydriodic Argentic Hydrochloric chloride. acid. iodide. acid. lodides.—The iodides closely resemble the chlorides and bromides. Argentic iodide, Agl, mercurous iodide, mercuric iodide, Hgl2, and cuprous iodide, are insoluble in water; plumbic iodide, Pbl2, dissolves very slightly; the other iodides are readily soluble. COMPOUNDS OF lODINE WITH CHLORINE. Hypiodous chloride, ICI. lodous chloride, IC13. These compounds are formed by the direct union of their elements. HYPIODOUS CHLORIDE. ICI. Molecular weight = 162.5. Fuses at 24.7° C. (76.5° F.). Boils at 101° C. (213.8° F.). Preparation.—This compound is obtained by passing dry chlorine over iodine, interrupting the operation as soon as the whole of the iodine has liquefied. The reddish-brown liquid thus obtained solidifies on standing. Properties.—Hypiodous chloride forms large prismatic crystals of a hyacinth-red color. Reaction.—Water decomposes it with formation of iodic acid, hydro- chloric acid, and free iodine: SICI + 30H3 - |®*jo + SHCI + 21, Hypiodous Water. lodic acid. Hydrochloric chloride. acid. lODOUS CHLORIDE. Molecular weight = 233.5. IC13. Preparation.—lodous chloride is formed by the action of an excess of chlorine upon iodine or upon the foregoing compound. Properties.—lt forms long yellow crystals which sublime at ordinary temperatures. It fuses at 20°-25° C. (68°-77° F.), with dissociation into hypiodous chloride and free chlorine. lODIC ANHYDRIDE. 301 Reaction.—With water it is decomposed, yielding the same products as hypiodous chloride (see preceding compound): 510, + 90H3 = 31 °)Io + 15HC1 + C lodous Water. lodic acid. Hydrochloric chloride. acid. COMPOUNDS OF lODINE WITH HYDROXYL. OXYGEN AND fS1 t O . o I 01 |o°L {S1 (OHo lodic anhydride, lodic acid, Periodic acid, The graphic formulae of these compounds are the corresponding chlorine compounds given on p analogous to those of .. 177. lODIC ANHYDRIDE f 01 o I 2Os or O . O 101 Preparation.—This compound is formed when iodic acid is heated to 170° C.: r oi r oi 2 { Olio = + 0H°' 01 lodic acid. lodic Water, anhydride. Properties.—lodic anhydride is a white crystalline powder possessing a sp. gr. of 4.48. Reactions.—1. When heated to 300° C. (572° F.) it is decomposed into iodine and oxygen. 2. Gaseous hydrochloric acid decomposes it with elevation of tem- perature, iodous trichloride and water being formed, and chlorine libe- rated : 302 INORGANIC CHEMISTRY. I 205 + 10HC1 = 2IC13 + 50H2 + 2CI2. lodic Hydrochloric lodous Water, anhydride. acid. trichloride. 3. It dissolves in water, forming iodic acid. lODIC ACID. / 01 \ OHo* Preparation.—l. lodic acid may be obtained by decomposing a solu- tion of baric iodate with the equivalent quantity of sulphuric acid : rOl I Of oi < Bao" + S02Ho2 = 2|2ho + S02Bao" I 01 Baric iodate, Sulphuric acid. lodic acid. Baric sulphate. The aqueous solution of iodic acid may be evaporated at 100° C. without decomposition. 2. It is best prepared by oxidizing iodine with strong boiling nitric acid: 6N02Ho + I.= 21 oho + 20H* + 2NA + 'N-A. Nitric acid. lodic acid. Water. Nitrous Nitric anhydride. peroxide. 3. It is also formed when chlorine is passed into water in which finely powdered iodine is suspended: I 2 60H2 -f- 5C12 2 Qjq0 + 10HC1. Water. lodic acid. Hydrochloric acid. Properties.—lodic acid forms colorless rhombic crystals of sp. gr. 4.629. It is very soluble in water. At a temperature of 170° C. (338° F.) it gives off water, and is converted into anhydride. Reactions.—1. In contact with hydriodic acid it forms water and iodine: (OHo 30H2 + 3I2. lodic acid. Hydriodic Water. acid. 2. It is reduced by many other deoxidizing agents; lODIC ACID. 303 2{oHo + 6SH* = I 2 + 5S + 60H2. lodic acid. Sulphuretted ■ Water, hydrogen. 21 qjj0 + 5S02 -f- 40H2 I 2 + 5S02Ho2, lodic acid. Sulphurous Water, Sulphuric acid, anhydride. Preparation of lodates.—lodates may be obtained by the following methods: 1. By treating solutions of metallic hydrates with iodine, and sep- arating the iodate by crystallization : 6KHo + 31* = SKI + |°J-o + 30H*. Potassic Potassic Potassic Water, hydrate. iodide. iodate. 2. By dissolving iodine in potassic hydrate and treating the mixture with chlorine: 12KHo + I 2 + 5CI2 = 10KC1 + 2{qk0 + 60H2. Potassic Potassic Potassic Water, hydrate. chloride. iodate. In this way the whole of the iodine is converted into iodate. 3. By heating together potassic chlorate and iodine: h + {OKo = ICI + {OKo' Potassic Hypiodous Potassic chlorate. chloride. iodate. Character of lodates.—The iodates are nearly all insoluble in water; those of the alkalies are the most soluble. lodic acid, though a monobasic acid, forms hyper-acid salts. Thus in the case of potassium, the following salts are known Normal potassic iodate, < OKo' Acid potassic iodate, | | q^o Di-acid potassic iodate, | OKo’ { OHo All the iodates are decomposed by heat. Some break up into iodides aud oxygen, others into metallic oxides, iodine, and oxygen: 304 INORGANIC CHEMISTRY, {oL = KI + 3°- Potassic Potassic iodate. iodide. fOI I o i Mgo" = MgO + I 2 + 50. O 101 Magnesic iodate. Magnesic oxide. PERIODIC ACID (OHo Preparation.—1. Periodic acid is obtained by decomposing plumbic periodate with sulphuric acid : f 01 o o roi - Pbo" + S02Ho2 = 2< O + SOaPbo". O (OHo O 101 Plumbic Sulphuric Periodic Plumbic periodate. acid. acid. sulphate. 2. Argentic periodate is decomposed on boiling with water into an insoluble basic salt of the formula I0SHAg2,0H2 and free periodic acid: roi roi 2< O + 20H2 = I0SHAg2;OH2 -f < O . (OAgo (OHo Argentic Water. Basic argentic Periodic periodate. periodate. acid. The periodic acid remains in solution and may be obtained on evap- roi oration in crystals of the formula < O ,20H2. ( OHo 3. It is also formed when iodine is added to an aqueous solution of perchloric acid: roci roi + I 2 = + Cl2. ( OHo ( OHo Perchloric acid. Periodic acid. PERIODIC ACID. 305 fOl Properties.—Normal periodic acid, < O , has not been obtained. (OHo The crystals which are formed when an aqueous solution of the acid is evaporated, contain two molecules of water of crystallization, which they retain at 100° C. They fuse between 130° and 136° C. (266°- 277° F.), and are slowly converted into iodic anhydride, oxygen, and water. At 200° C. (392° F.) this change takes place rapidly. Preparation of Periodates.—l. Sodic periodate may be prepared by passing chlorine through mixed solutions of sodic hydrate and sodic iodate: | + 2NaHo + Cl2 =IO + 0H2 + 2NaCI. 1 ojNao I ONao Sodic Sodic Sodic Water. Sodic icxiate. hydrate. periodate. chloride. A basic baric periodate may be obtained by heating baric iodate: f OI f 01 O o o s[pb°-2PBo Pentargentic periodate IOAgo5 ivO//3fAgo,20Ag2 Trihydric diargentic periodate lOHosAgo2 2,vO//3IAg0,OAg2,3OH2 Pentabaric periodate | // 10Bao 5 ;Vvg7/jBao^,4BaO f i02Ago2 Tetrargentic anhydroperiodate 1° I I02Ago2 2ivO//3IAgo,OAg2 f IO Zno//2 ::°;>{Zno"f3ZnO Tetrazincic anhydroperiodate 0 (lOZno", The periodates are, as a rule, only sparingly soluble in water. FLUORINE, F2? Atomic weight =l9. Molecular weight —3B (?). Atomicity Evi- dence of atomicity: Hydrofluoric acid, . HF. Occurrence.—Fluorine occurs in nature in combination with metals as fluorides. The most common of these is calcic fluoride or fluorspar, CaF2, known also as the Derbyshire spar. Cryolite,a mineral occurring in Greenland, is a double fluoride of sodium and aluminium, possessing the formula Fluorine also occurs in small quantity in various other minerals, such as apatite, topaz, etc. In the animal kingdom it has been found in minute traces in the enamel of the teeth and, in the bones. Attempts to isolate Fluorine.—Little is known of fluorine in the free state. So great is the affinity of this element that as soon as it is ex- pelled from one combination it enters into another. Its isolation has from time to time been announced, but a repetition of the experiments by other investigators has, till lately, failed to confirm the supposed results. Argentic fluoride is decomposed at a red heat by chlorine, bromine, or iodine, with formation of a chloride, bromide, or iodide of silver; but the liberated fluorine instantly combines with the material of which the vessels employed in the experiment are composed. Ves- sels of glass, silver, gold, platinum, and graphite have been tried, but 307 HYDROFLUORIC ACID. without success. In like manner in the electrolysis of fused fluorides, the fluorine combines with the material of the positive electrode. The attempt to employ vessels of fluorspar in the above decomposi- tions has proved unsuccessful. Latterly, however, Brauner, in heating ceric tetrafluoride, has found that it is converted into diceric hexafluoride, whilst a gas is evolved which smells like chlorine, and liberates iodine from potassic iodide. The reaction probably occurs according to the equation : 2CeF4 = 'Ce'"2F6 + F2. Ceric Diceric tetrafluoride. hexafluoride. COMPOUND OF FLUORINE WITH HYDROGEN HYDROFLUORIC ACID. HF. Molecular weight = 20.* Molecular volume I I I. 1 litre weighs 10 criths. Boils at 19.4° C. (66.9° F.). Bp.gr. of liquid 0.9875 at 13° C. (55° F.). Preparation.—l. Hydrofluoric acid may be obtained by heating calcic fluoride or cryolite with concentrated sulphuric acid in a leaden or platinum retort (Fig. 44), which is connected with a U-tube of the same metal: CaF2 + S02Ho2 = 2HF + S02Cao". Calcic Sulphuric Hydrofluoric Calcic fluoride. acid. acid. . sulphate. A very concentrated acid distils over and condenses in the U-tube, which is cooled by a freezing mixture. If an aqueous solution is required, the acid may be passed at once into water. 2. In order to obtain a perfectly anhydrous acid, the double fluoride of hydrogen and potassium (HF,KF), which must be previously fused in order to free it from the last traces of moisture, is heated in a plati- num retort. The condenser and receiver must also be of platinum. The anhydrous hydrofluoric acid distils over, whilst potassic fluoride remains behind in the retort. The condensation is effected by means of * Kletzinsky has found that hydrofluoric acid at a temperature just above its boiling- point possesses a vapor-density corresponding with the molecular weight 40, and there- forepvith the molecular formula H2F2. Mallet, experimenting at a temperature of 25° C. (77° .F.), arrived at the same result. The vapor-density at these temperatures is twice as great as at 100° C., at which temperature it corresponds as above with the formula HF. The existence of such a molecule as H.2F2 could best be accounted for on the supposition that fluorine is a triad in this compound, thus: H—F=F—H. This view finds further support in the existence of a hydric potassic fluoride, which would thus be formulated: H—F=F—K. The greater jnolecular weight of hydrofluoric acid at lower temperatures accounts also for the relatively high boiling-point of this acid as compared with the other hydracids. 308 INORGANIC CHEMISTRY. a freezing mixture, and great care is required in performing the opera- tion, owing to the dangerous properties of the anhydrous acid. Fig. 44. Properties.—Anhydrous hydrofluoric acid is a colorless, mobile liquid which fumes strongly in contact with the air. It may be cooled to —B4° C. (—29.2° F.) without solidifying. Water absorbs the gaseous acid readily, forming a solution which, when saturated, possesses a sp. gr. of 1.25. This solution gives off* a portion of its acid on distillation until the sp. gr. has decreased to 1.15, when it distils unchanged at T2o° C. (248° F.), and contains from 36 to 38 per cent, of anhydrous acid. This acid of constant boiling-point does not correspond with any definite hydrate. The concentrated acid when brought in contact with the skin produces dangerous wounds which are very difficult to heal. The vapor of the anhydrous acid when inhaled has been known to prove fatal. Reactions.—l. Aqueous hydrofluoric acid dissolves many of the metals with evolution of hydrogen and formation of fluorides; Fe + 2HF - FeF2 + H2. Hydrofluoric Ferrous acid. fluoride. 2. It acts upon silicic anhydride and silicates, forming silicic fluoride and water: Si02 + 4HF = SiF, + 20H2. Silicic Hydrofluoric Silicic Water, anhydride. acid. fluoride. Thus hydrofluoric acid dissolves glass. This characteristic property is employed as a test for hydrofluoric acid and fluorides. All metallic fluorides, when treated with sulphuric acid, evolve hydrofluoric acid. The substance to be tested is placed in a small platinum or leaden dish with concentrated sulphuric acid, and the dish is covered with a piece of glass coated with wax, on which characters have been traced, so as to remove the wax from the parts written upon. The vessel is very gently warmed, and the glass is allowed to remain over it for about a quarter of an hour. On removing the wax, the presence of hydro- fluoric acid will be indicated by the etching of the exposed parts of the glass. This method is frequently employed in etching scales on glass, the fumes from a mixture of powdered fluorspar and sulphuric SILICON. 309 acid being employed for this purpose. Etchings produced by means of the aqueous solution of the acid are transparent and cannot be seen at a distance; when the gaseous acid is employed, the etched surface is dull, for which reason the use of the gaseous acid is preferred. It is evident from the above that neither glass nor porcelain vessels can be employed in the preparation or storing of hydrofluoric acid. The aqueous solution is generally kept in vessels of caoutchouc or guttapercha. Pure and perfectly dry hydrofluoric acid is without action upon glass (Gore); but the slightest trace of moisture induces the action just described. Fluorides.—The fluorides are formed by dissolving metals in hydro- fluoric acid or by acting with this acid on oxides, hydrates, or carbon- ates. The fluorides of the alkalies and of silver are soluble; those of the alkaline earths are insoluble. Nearly all the fluorides form molecu- lar compounds with hydrofluoric acid, such as the double fluoride of hydrogen and potassium already mentioned. CHAPTEP XXIX. TETRAD ELEMENTS. Section I. {Continued from Chapter XXV.). SILICON, SUioium, Si Atomic weight 28.2. Sp. gr. (crystallized) = 2.49. Atomicity iv, also a pseudo-triad. Evidence of atomicity: Silicic chloride, . . . . SiivCl4. Silicic fluoride, . . . . . SiivF4. Disilicic hexachloride, . . / SiClg •\sicl3. History,—Silicon was first isolated by Berzelius in 1810. Occurrence.—Silicon is, with the exception of oxygen, the most abun- dant and widely distributed of the elements. It does not occur in the free state. In combination with oxygen it forms the mineral quartz or silica, which is the anhydride of silicic acid: whilst the compounds of silica with bases constitute the chief constituents of the rocks which compose the earth’s crust, and consequently of the soils, which have all been produced by the disintegration of the rocks. From the soils the silicon is absorbed by plants, in the ashes of which it may always be detected. Preparation.—1. Silicon is liberated when silicic anhydride is re- duced by heating it with sodium: 310 INORGANIC CHEMISTRY, Si02 + 2Na2 = Si + 20Na2. Silicic Sodic anhydride. oxide. This method is not, however, adapted for the preparation of pure silicon. The reaction may be shown by heating sodium in a test-tube of Bohe- mian glass, when the glass speedily blackens owing to the reduction of the silica. 2. Pure silicon may be readily obtained by heating potassic silico- fluoride with potassium: SiK2F6 + 2K2 = Si + 6KF. Potassic Potassic silicofluoride. fluoride. Sodium may be substituted for potassium in this reaction. The fused mass is allowed to cool, and the potassic fluoride is then dissolved in water, when the silicon remains behind as a brown amorphous powder. 3. Silicon is deposited in the same amorphous condition when sodium is heated in a current of the vapor of silicic chloride : SiCl4 + 2Na2 = Si + 4NaCI. Silicic Sodic chloride. chloride. 4. In order to obtain silicon in the crystallized condition, advantage is taken of the property which this element possesses of dissolving at a high temperature in certain metals, such as zinc or aluminium, and crystallizing from these metallic solutions on cooling. A mixture of 15 parts of dry potassic silicofluoride, with 4 parts of sodium in thin slices, is thrown into a red-hot Hessian crucible; 36 parts of granulated zinc are quickly added, and the mass is covered with a layer of fused sodic chloride. The lid is then replaced and the whole is heated for some time to a temperature below the boiling point of zinc. On dissolving the cooled regulus of zinc in acids, the crystallized silicon remains be- hind. 5. Crystallized silicon may also be obtained by heating together in a crucible 1 part of aluminium with 5 parts of glass free from lead, and 10 parts of powdered cryolite (/Al///2F6,6NaF). The silica of the glass is reduced at the expense of the aluminium. Properties.—Amorphous silicon is a brown powder, devoid of lustre. It inflames when heated in the air, but cannot be entirely burnt, even in oxygen, as the silica which is formed coats the particles and prevents further oxidation. It is insoluble in water, and is not attacked by acids, except hydrofluoric acid, which dissolves it readily. When heated with exclusion of air it becomes denser, and is no longer combustible. Crystallized silicon forms dark lustrous octahedra, which possess a sp. gr. of 2.49 and are hard enough to scratch glass. It may be heated to whiteness in oxygen without burning. At a very high temperature it fuses. It conducts electricity imperfectly. Acids are without action upon it, with the exception of a mixture of nitric and hydrofluoric acids, which dissolves it slowly. SILICIC HYDEIDE. 311 Reactions.—1. When amorphous silicon is heated in oxygen, silicic anhydride is formed. 2. Both varieties of silicon may be burned in a stream of chlorine, silicic chloride being produced. Owing to the volatile nature of the silicic chloride, the whole of the silicon may be thus converted. 3. When amorphous silicon is treated with hydrofluoric acid, or the crystallized variety with a mixture of nitric and hydrofluoric acids, hy- drofluosilicic acid is formed ; Si + 6HF = SiH2F6 + 2H2. Hydrofluoric Hydrofluo- acid. silicic acid. 4. Amorphous silicon when boiled with caustic alkalies, yields an. alkaline silicate, with evolution of hydrogen : Si + 4KHo = SiKo, + 2H2. Potassic Potassic hydrate. silicate. Crystallized silicon must be fused with the alkali in order that this re- action may take place. COMPOUND OF SILICON WITH HYDROGEN. SILICIC HYDRIDE, Siliciuretted Hydrogen. SiH,. Molecular weight = 32,2. Molecular volume I I I. Preparation.—l. When dilute sulphuric acid is decomposed by a feeble electric current passing from electrodes of aluminium containing silicon, silicic anhydride is evolved at the negative electrode. 2. It may also be obtained by decomposing magnesic silicide with hydrochloric acid: SiMg" + 4HCI = 2MgCl2 + SiH4. Magnesic Hydrochloric Magnesic Silicic silicide. acid. chloride. hydride. The magnesic silicide is prepared by heating together in a closed cruci- ble 40 parts of anhydrous magnesic chloride, 35 parts of dried sodic silicofluoride, 10 parts of fused sodic chloride, and 20 parts of sodium in thin slices. The fused mass is broken into fragments and intro- duced into a flask fitted with safety and delivery tubes. The flask and the delivery tube are filled with water from which the air has been ex- pelled by boiling, and hydrochloric acid is then poured through the funnel of the safety tube into the flask. Silicic hydride is evolved and is collected over previously boiled water in the pneumatic trough. 312 INORGANIC CHEMISTRY. 3. Silicic hydride prepared by either of the foregoing processes is always contaminated with hydrogen ; but if ethylic silico-orthoformate, a substance obtained by the action of silicon-chloroform (q.v.) on abso- lute alcohol, be placed in contact with sodium, it breaks up into ethylic orthosilicate and pure silicic hydride, the sodium remaining unaffected: 4SiH(C2HSO)3 = SiH4 + 3Si(C2HgO)4. Ethylic silico- Silicic Ethylic orthoformate. hydride. orthosilicate. Properties.—Silicic hydride is a colorless gas. When prepared from magnesic silicide it inflames spontaneously in contact with air, and burns with a brilliant white flame evolving dense clouds of silicic an- hydride. The pure gas is not spontaneously inflammable, but it acquires this property when it is gently warmed, or when the pressure is reduced, or when it is diluted with hydrogen. Reactions.—l. Burned in the air or oxygen, silicic hydride yields silicic anhydride and water: SiH4 + 202 == Si02 + 20H3. Silicic Silicic Water, hydride. anhydride. 2. With chlorine it explodes spontaneously, forming silicic chloride and hydrochloric acid: SiH4 + 4C12 = SiCl4 + 4HCI. Silicic Silicic Hydrochloric hydride. chloride. acid. 3. When heated, it is decomposed into amorphous silicon and free hydrogen, the latter occupying twice the volume of the original gas. 4. It is decomposed at ordinary temperatures by a solution of potas- sic hydrate, yielding four times its volume of hydrogen : SiPI4 + 2KHo + 0H2 = SiOKo2 + 4H2. Silicic Potassic Water Potassic hydride. hydrate. silicate. 5. It precipitates some of the heavy metals in the form of silicides from the solutions of their salts : 2S02Cuo" + SiH4 = SiCu"2 + 2S02Ho2. Cupric Silicic Cupric Sulphuric sulphate. hydride. silicide. acid. SILICIC CHLORIDE—DISILICIC HEXACHLOEIDE. 313 COMPOUNDS OF SILICON WITH THE HALOGENS. SILICIC CHLORIDE. SiCl4. Molecular weight 170.2. Molecular volume L_U. 1 litre of the vapor weighs 85.1 criths. Sp. gr. of liquid 1.52. Boils at 59° C. (138.2° F.). Preparation.—l. Silicic chloride is formed by the direct combination of its elements when silicon is burnt in chlorine. 2. It is most conveniently prepared by heating a mixture of finely divided carbon and silicic anhydride in a stream of dry chlorine: Si02 + 2C + 2C12 = SiCl4 + 2CO. Silicic Silicic Carbonic anhydride. chloride. oxide. Properties.—Silicon tetrachloride is a colorless mobile liquid, fuming strongly in contact with air. Reaction.—Water decomposes it instantaneously with formation of silicic and hydrochloric acids: SiCl4 + 40H2 = SiHo4 + 4HCI. Silicic Water. Silicic Hydrochloric chloride. acid. acid. DISILICIC HEXACHLORIDE. f SiClg I SiClg • Molecular weight = 269.4. Molecular volume | | |. 1 litre of the vapor weighs 134.7 criths. Sp. gr. of liquid 1.58. Fuses at —l° C. (30.2° F.). Boils at 147° C. (296.6° F.). Preparation.—1. This compound is formed in small quantity when the vapor of silicic chloride is passed over silicon heated above 1000° C.: BSIOI. + Si _ Silicic Disilicic chloride. hexachloride. 2. It is more easily prepared by gently heating disilicic hexiodide (q.v.) with mer curie chloride: {Ill: + 3Hsa = {“s + SH*I, Disilicic Mercuric Disilicic Mercuric hexiodide. chloride. hexachloride iodide. Properties.—Disilicic hexachloride is a mobile, colorless liquid, which at a tempera- ture of —l° C. solidifies to a crystalline mass. It possesses the peculiarity of being stable only below 350° C. and above 1000° C., whilst at intermediate temperatures it dissociates into silicic chloride and silicon. A similar abnormal behavior has already been noted in the case of seleniuretted and telluretted hydrogen. Reaction.—With water it is decomposed into silicon-oxalic acid and hydrochloric acid: 314 INORGANIC CHEMISTRY. {IIS: + '•OH, = {IS + «HCL Disilicic Water. Silicon-oxalic Hydrochloric hexachloride. acid. acid. SILICON CHLOROFORM, Silicic Hydrotrichloride. SiIICl3. Molecular weight = 135.7. Molecular volume | | |. 1 litre of the vapor weighs 67.85 criths. Sp. gr. of liquid 1,6. Boils at 36° C. (96.8° F.) Preparation.—Silicon chloroform is formed when silicon is heated to dull redness in a current of hydrochloric acid gas: Si + 3HCI = SiHClg + H2. Hydrochloric Silicon acid. chloroform. Properties.—Silicon chloroform is a colorless liquid. It is very inflammable, and burns with a green-edged flame. A mixture of its vapor with air explodes in contact with a heated body. Reactions.—1. It is decomposed by chlorine at ordinary temperatures: SiHClg + Cl, = SiCl4 + HCI. Silicon Silicic Hydrochloric chloroform. chloride. ' acid. 2. By contact with water it is transformed into silicoformic anhydride, or disilicic hydrotrioxide : f Si HO 2SiHCl3 + 30H2 = \ O + 6HCI. [ Si HO Silicon Water. Silicoformic Hydrochloric chloroform. anhydride. acid. Silicon bromoform, SiHBr3, and silicon iodoform, SiHl3, have also been prepared. SILICIC BROMIDE. Molecular weight = 348.2. Molecular volume || |. Fuses at —l3° C. (6.6° F.). Boils at 153° C. (307.4° F.). Sp. gr. of liquid 2.813 at 0° C. Sißr4. Preparation.—This substance is obtained by a method analogous to that employed in the preparation of the chloride, bromine-vapor being substituted for chlorine. Properties.—It is a fuming, colorless liquid. Reaction.—Water decomposes it with formation of silicic and hydrobromic acids; Sißr4 + 40 H2 = SiHo4 + 4HBr. Silicic Water. Silicic Hydrobromic bromide. acid. acid. Disilicic hexabromide, | is also known. SILICIC lODIDE sir4. Molecular weight = 536.2. Molecular volume | | |. Fuses at 120.5° C. (248.9° F.). Boils in carbonic anhydride at 290° C. (554° F.j. Preparation.—This compound is formed by the direct union of its elements when a mixture of iodine vapor and carbonic anhydride is passed over red-hot silicon. The SILICIC FLUORIDE. 315 carbonic anhydride serves to carry off the vapor of the silicic iodide as fast as it is formed, and thus to prevent its decomposition. Properties.—Silicic iodide crystallizes in colorless octahedra. It may be distilled in a current of carbonic anhydride. It is soluble in carbonic disulphide. Reactions.—1. Water decomposes it into silicic and hydriodic acids. 2. Absolute alcohol decomposes it, with production of silicic anhydride, ethylic iodide, and hydriodic acid: Sil4 + 2EtHo = Si02 + 2EtI + 2HT. Silicic Alcohol. Silicic Ethylic Hydriodic iodide. anhydride. iodide. acid. r gy Disilicie hexiodide, -< gy3, has been obtained by heating silicic iodide with finely divided silver: 2SiT4 + Ag2 = + 2AgT. Silicic Disilicic Argentic iodide. hexiodide. iodide. It forms hexagonal crystals, fusing with decomposition at 250° C. SILICIC FLUORIDE. SiF4. Molecular weight = 104.2. Molecular volume I I I. 1 litre weighs 52.1 criths. Fuses at —l4o° C. (—22o° F.). Boils at —lo7° C. (—160.6° F.). Preparation.—Silicic fluoride is prepared by heating together, in a flask furnished with a delivery tube, quartz sand, fluorspar, and an excess of concentrated sulphuric acid : Si02 + 2CaF2 + 2S02Ho2 = SiF4 + 2SOHo2Cao". Silicic Calcic Sulphuric Silicic Dihydric calcic anhydride. fluoride. acid. fluoride. sulphate. The gas may be collected in perfectly dry glass vessels over mercury. Properties.—Silicic fluoride is a colorless gas with a very pungent odor. It fumes strongly in contact with air. Under a pressure of 30 atmospheres, or at a temperature of—lo7° C. (—160.6° F.), it con- denses to a colorless liquid, which at a still lower temperature solidities. It is not altered by exposure to the heat of powerful electric sparks. Reaction.—Water decomposes it with formation of silicic and hydro- fluosilicic acids: 3SiF4 + 40H2 = SiHo4 + 2SiH2F6. Silicic Water. Silicic Hydrofluosilicic fluoride. acid. acid. When the gas is passed into water, the silicic acid separates out as a gelatinous mass, whilst the hydrofluosilicic acid remains in solution. To prevent the delivery tube from being stopped up, it must dip under mercury at the bottom of the vessel in which the water is contained. The liquid is afterwards filtered from the silicic acid and evaporated at 316 INORGANIC CHEMISTRY. a low temperature. The aqueous solution of hydrofluosilicic acid thus obtained forms a fuming acid liquid, which on further evaporation de- composes into silicic fluoride and hydrofluoric acid. With metallic oxides, hydrates, and some salts, hydrofluosilicic acid produces silicofluorides: SiH2F6 + 2KHo = SiK2F6 + 20H2. Hydrofluo- Potassic Potassic Water, silicic acid. hydrate. silicofluoride. In contact with an excess of base the silicofluorides are decomposed, yielding silicates and fluorides : SiK2F6 + BKHO = SiKo4 + 6KF + 40H2. Potassic Potassic Potassic Potassic Water, silicofluoride. hydrate. silicate. fluoride. The silicofluorides of barium and potassium are insoluble in water. Disilicic hexafluoride has been prepared by passing silicic fluoride over melted silicon ; 3SiF4 + Si = 2 {li Fg- Silicic Disilicic fluoride. hexafluoride. It forms a fine white powder. COMPOUNDS OF SILICON WITPI OXYGEN AND HYDRYOXYL. Silicic anhydride, Si02. Silicic acid, SiHo4 and SiOHo2. Other Modifications of Silicic Acid. 5i.,03H02 SLO10Ho4 Si305Ho2 SiG08Ho8 Si407Ho2 Si8G15Ho2 Si404Ho8. SILICIC ANHYDRIDE, Silica. Si02. Molecular weight = 60.2. Sp. gr. (amorphous) 2.2, (tridymite), 2.3, (quartz} 2.69. Occurrence.—Some of the forms in which silicic anhydride is found in nature have already been alluded to (p. 309). It -occurs in the crys- tallized condition as quartz and tridymite, and in an amorphous form as opal. SILICIC ANHYDRIDE. 317 Preparation.—lt may be obtained by heating silicic acid to 100° C. Water is given off and amorphous silicic anhydride remains. Properties.—As quartz or rock crystal, silicic anhydride occurs in the form of hexagonal prisms terminated by a hexagonal pyramid (Fig. 45). The crystals are sometimes colorless, sometimes colored by the presence of various oxides. Amethyst quartz, rose quartz, smoky quartz, are Fig. 45. varieties of this description named according to their color. Occa- sionally quartz occurs in large crystalline masses as quartzose rock. It has a sp. gr. of 2.69, and is hard enough to scratch glass. Tridymite is a second crystallized variety of silicic anhydride found in various trachytic rocks. Like rock crystal, it crystallizes in forms belonging to the hexagonal system; but the relations of the axes vary in the two minerals, so that the forms of the one cannot be referred to those of the other. The sp. gr. of tridymite is 2.3. Amorphous silicic anhydride, when artificially prepared, forms a white, very fine powder. As opal, amorphous silica occurs in transparent or translucent masses with a conchoidal fracture. The sp.gr. of the arti- ficial variety is 2.2 ; that of the natural 2.3. Silicic anhydride in all its forms is insoluble in water at ordinary temperatures. It dissolves slightly, however, if heated with water un- der pressure to low redness, and, on cooling, crystallizes from the solu- tion in the form of quartz. In like manner, when a solution of an al- kaline silicate is heated in a sealed glass tube, a portion of the silica from the glass is dissolved, forming an acid silicate. Gn cooling, the excess of silica separates out. If the separation takes place above 180° 0.(356° F.) the silica is obtained as quartz; below this temperature tridymite is formed; at ordinary temperatures it is deposited in the hy- drated condition as amorphous silicic acid. Acids, with the exception of hydrofluoric, are without action upon silicic anhydride. With aqueous hydrofluoric acid hydrofluo-silicic acid is formed : Si02 + 6HF = SiH2F6 + 20H2. Silicic Hydrofluoric Hydrofluosilicic Water, anhydride. acid. acid. All the modifications of silicic anhydride, when fused with an excess of a caustic alkali or an alkaline carbonate, combine with the base to form a soluble silicate : Si02 -f- 2CO]\ao2 = SiNao4 -f- 2C02. Silicic Sodic Sodic Carbonic anhydride. carbonate. silicate. anhydride. 318 INORGANIC CHEMISTRY. The amorphous variety, if it has not been ignited too strongly, dissolves in boiling solutions of caustic alkalies. SILICIC ACID. Tetrabasic, . . SiHo4, Dibasic, . . . SiOHo2. Preparation.—l. Silicic acid may be obtained by decomposing a solution of sodic or potassic silicate with hydrochloric acid: SiNao, + 4HCI = SiHo4 + 4NaCI. Sodic Hydrochloric Silicic Sodic silicate. acid. acid. chloride. If the solution of the alkaline silicate is concentrated, the silicic acid separates out as a white gelatinous precipitate; but if a dilute solution of the silicate be poured into an excess of hydrochloric acid, the silicic acid remains dissolved. The clear solution obtained by the latter method maybe freed from the sodic chloride and excess of hydrochloric acid by dialysis (see Introductions p. 130). The silicic acid, being a colloid, is unable to pass through the membrane of the dialyzer, whilst the other substances in solution diffuse freely through into the surround- ing liquid. The solution of silicic acid may be concentrated by boiling in a flask until it contains 22 per cent, of the tetrabasic acid, but beyond this point it solidifies to a jelly. When evaporated in a dish the solu- tion is apt to gelatinize round the edges, and then the whole mass solid- ifies. The concentrated solution also gelatinizes spontaneously when allowed to stand for a few days, and the same effect is produced instan- taneously by passing carbonic anhydride into the solution, or by the addition of a trace of an alkaline carbonate. 2. Gelatinous silicic acid may be obtained by passing a stream of car- bonic anhydride through a solution of an alkaline silicate: SiNao4 + 40H2 + 4C02 = SiHo4 + 4COHoNao. Sodic Water. Carbonic Silicic Hydric sodic silicate. anhydride. acid. carbonate. A reaction similar to this is the cause of the disintegration of granitic rocks. The carbonic anhydride which is held in solution in all natural waters acts upon the alkaline silicates contained in the rocks. 3. Gelatinous silicic acid is also formed when silicic fluoride is passed into water (p. 315). Properties.—Silicic acid, like most other weak poly basic acids of even basicity, has a great tendency to give off water and form an anhydride. It is therefore exceedingly doubtful whether any of the silicic acids have been prepared in a state of purity. By allowing gelatinous silicic acid to dry in the air, a compound having approximately the composi- tion represented by the formula Si,;08Ho8 is obtained, and this, when dried at 100° C., parts with more water, yielding a hydrate of the formula Si6H10Ho4. These substances are, however, very difficult to obtain of SILICIC ACID. 319 fixed composition, and they possess none of the other characteristics of definite chemical compounds. Silicates.—The preparation of alkaline silicates has already been de- scribed (p. 317). Silica and the silicates form a very important class of minerals. The following list contains a few examples; Sand Flint Quartz- Silicic anhydride, , . . . Si02. Opal Chalcedony Peridote (Dimagnesic silicate), SiMgo"2. Phenacite (Diberyllic silicate), Sißeo"2. Willemite (Dizincic silicate), ..... SiZno"2. Zircon [Zir conic silicate), SiZroiT. Enstatite (Monomagnesic silicate),.... SiOMgo". ( SiNao3 Yorke’s Sodic silicate, <0 ( SiNao3 fSi~- Ophite (Noble Serpentine), < O Mgo/r3. ISi 1 f Si Diopside (Calcic magnesia disilicate),. . . \ SiCao"Mgo". Talc (Tetramagnesic pentasilicate), . . . Sir)06Mgo"4. r SiHo2—, Okenite (Tetrahydric calcic disilicate), . .< O Cao". I SiHo2—1 ( SiHoMgo" Serpentine (Dihydric trimagnesic disilicate), < Mgo" ( SiHoMgo" Steatite (Trimagnesic tetrasilicate), . . . Si}05Mgo//3. f SiHoMgo7' Meerschaum (Tetrahydric dimagnesic trisili- j g^jj0 cold), j 0 2 [ SiHoMgo" fSiOHo-1 Pyrophyllite (.Dihydric aluminic tetrasili- \ SiO IJvi * cate), 1 SiO , ? SiOHo—! Anorthite (Aluminic calcic disilicate),. . . Si.>(/AF//206)TiCao'/. (SiO Labradorite (Aluminic calcic trisilicate),. . < SiCao"—AloTl. (SiO——1 * Alovi = (/Al///206)vi. 320 INORGANIC CHEMISTRY. ( SiCao" , Grossularia (Aluminic tricalcic trisilicate), SiCao"—AloTI. ( SiCao" 1 Emerald (Triberyllic aluminic hexasilicafe),. Si6OOAIov,Beo"3. rsio , Cliloropal (Ferric trisilicate), < SiO—Feov 1,30 ( SiO 1 Felspar. Orthose {Dipotassic aluminic hexa- /gj q p^n vi silicate), p 6s COMPOUNDS OF SILICON CONTAINING SULPHUR. SILICIC SULPHIDE. SIS"* Preparation.—1. Silicic sulphide is formed by the direct union of its elements when amorphous silicon is heated in sulphur vapor. 2. It is more conveniently obtained by passing the vapor of carbonic disulphide over a mixture of silicic anhydride and charcoal heated to redness ; Si02 + CS2 + C = SiS" + 2CO. Silicic Carbonic Silicic Carbonic anhydride. disulphide. * sulphide. oxide. Properties.—Silicic sulphide forms white silky needles resembling asbestos in ap- pearance. It may be sublimed without decomposition. In contact with water it forms silicic acid and sulphuretted hydrogen: SiS//2 + 40H2 = SiHo4 + 2SH2. Silicic Water. Silicic Sulphuretted sulphide. acid. hydrogen. SILICIC TRICHLOHSULPHHYDRATE. SiCl3Hs. Molecular weight = 167.7. Molecular volume \ I |. Boils at 96° C. (204.8° F.). Preparation.—This compound is obtained by passing a mixture of silicic chloride vapor and sulphuretted hydrogen through a red-hot porcelain tube: SiCl4 + SII2 = SiClgHs + HCI. Silicic Sulphuretted Silicic trichlor- Hydrochloric chloride. hydrogen. sulphhydrate. acid. Properties.—Silicic trichlorsulphhydrate is a colorless fuming liquid, boiling at 96° C. (204.8° F.). Water decomposes it, forming silicic acid, hydrochloric acid, and sulphuretted hydrogen: SiCl3Hs + 40H2 = SiHo4 + BHCI + SHa. Silicic trichlor- Water. Silicic Hydrochloric Sulphuretted sulphhydrate. acid. acid. hydrogen. * Feovi = (/Fe'//206)Ti. 321 TIN, Sn.* Atomic weight = 118. Sp. gr. 7.28. Fuses at 228° C. (442.4° F.). Atomicity " and IT, and also a pseudo-triad. Evidence of atomicity: Stannous chloride (at 900° C.), Sn"Cl2. Stannic chloride, SnivCl4. History.—Tin has been known from the earliest historical times. The tin -mines of Cornwall were celebrated before the Roman invasion, and from these the Phoenician merchants supplied the metal to the ancient world. Occurrence.—Tin is never found in the free or native state. In combination with oxygen as tin-stone or stannic anhydride, it occurs m veins in the primitive rocks, and sometimes in alluvial deposits (stream tin). Tin-stone is the only ore from which the metal is ex- tracted. The mines of Cornwall, above referred to, and those of Devonshire, furnish the chief supply ; those of Malacca and Banca come next in importance. Extraction.—The tin-stone is first crushed and washed in order to free it from earthy impurities. It is then roasted in a reverberatory furnace, by which means the iron- and copper-pyrites with which it is contaminated are oxidized. The iron is thus converted into ferric oxide, with evolution of sulphurous anhydride, whilst the copper forms cupric sulphate. The roasted mass is again washed, the cupric sul- phate being thus dissolved and the ferric oxide mechanically removed. The finely divided tin-stone thus purified is mixed with charcoal and reduced in a furnace: Sn02 + C 2 = Sn + 2CO. Stannic anhydride. Carbonic oxide. The tin obtained by the above process is generally contaminated with various foreign metals (iron, copper, lead, arsenic, antimony), from which it may be separated by liquation. This process consists in melting the crude tin at the lowest possible temperature on the bed of a reverberatory furnace. The tin, by virtue of its lower fusing-point, melts first, and flows off, leaving a less fusible alloy of tin with the other metals. Properties.—Tin is a white metal with a high metallic lustre. When Warm it emits a peculiar odor. In hardness it is intermediate between lead and zinc. It is malleable and may be beaten into thin leaves (tin-foil). At a temperature of 200° C. it becomes brittle. It fuses at 228° C. (442.4° F.), and when exposed to the air in a molten condition Tliis element, whilst exhibiting all the physical properties of a metal, behaves in uiost of its chemical relations like a non-metal. Its compounds resemble those of car- on, silicon, and titanium, and it can be most conveniently studied in connection with ese dements. For similar reasons antimony, bismuth, and a few other metallic ele- ments have, in the present work, been classed with the non-metals. 322 INORGANIC CHEMISTRY. undergoes superficial oxidation. At a white heat it enters into ebul- lition and burns with a brilliant white light, forming stannic anhy- dride. It is also oxidized when heated to redness in a current of steam. At ordinary temperature it resists the action of air and moisture. If a bar of tin' be bent backwards and forwards a faint crackling sound is heard, and the point ot flexure becomes hot. These effects depend upon the breaking and friction of the crystals within the mass. The crystalline structure of tin may be readily shown by brushing the surface of a piece of the metal (which has been cast but not hammered) with warm dilute aqua-regia, when it becomes covered with fine crystal- line markings, resembling in appearance, watered silk. Tin thus pre- pared was formerly much used for ornamental purposes under the name of moirSe metallique. Crystals of tin maybe readily obtained by fusing a large quantity of the metal, allowing it partially to solidify in the crucible, then breaking a hole in the crust which forms on the surface, and pouring out the molten metal. The interior of the crucible will be found to be lined with crystals of tin. Reactions.—1. Hot concentrated hydrochloric acid dissolves tin with evolution ot hydrogen and formation of stannous chloride : Sn + 2HCI = SnCl2 + H2. Hydrochloric Stannous acid. chloride. 2. Heated with concentrated sulphuric acid it forms stannous sul- phate, sulphurous anhydride being evolved : Sn + 2S02Ho2 = S02Sno" + S02 + 20Ha. Sulphuric Stannous Sulphurous Water, acid. sulphate. anhydride. 3. Nitric acid of sp. gr. 1.3 acts upon it violently, oxidizing it to metastannic acid (SnsOsHolo). Nitric acid of sp. gr. 1.5 does not attack tin. 4. Cold dilute nitric acid dissolves it slowly without evolution of gas, stannous nitrate being formed. At the same time a portion of the nitric acid undergoes reduction to ammonia, which combines with the excess of nitric acid: 4Sn + 9N02Ho = 4^Q2Sno" + NH3 -f 30 H2. Nitric acid. Stannous nitrate. Ammonia. Water. 5. Caustic alkalies dissolve tin when fused with it, a soluble stannate being formed, whilst hydrogen is evolved : Sn + 20KH + 0H2 == SnOKo2 + 2H2. Potassic Water. Potassic hydrate. , stannate. 6. It combines directly with sulphur, phosphorus, chlorine, bromine, and iodine. COMPOUNDS OF TIN. 323 Uses.—Tinning.—Tin is frequently employed in coating other metals to preserve them from rust, a process known as tinning. Ordinary tin-plate is iron which has been thus treated. The surface of the metal to be tinned is thoroughly freed from every trace of oxide, which would otherwise prevent the adhesion of the tin, and the metal is then plunged into a bath of melted tin, covered with a layer of grease to exclude the air. The film of tin which adheres to the surface forms an alloy with the metal, and cannot be separated from it mechanically. The tinning of copper is effected in a similar manner. Alloys.—Numerous alloys of tin are employed in the arts. Plumber’s solder is an alloy of tin and lead, the proportion of tin increasing with the degree of fusibility required.* Fine solder consists of 2 parts of tin and 1 of lead; common solder of equal parts of tin and lead; and coarse solder of'l of tin and 2of lead. Britannia metal consists of equal parts of brass, tin, and antimony, and is employed as a cheap substitute for silver in the manufacture of teapots, etc: Pewter is a similar alloy, in which, however, the lead and tin greatly predominate. The alloys of tin with copper will be treated of under the heading of the latter metal. COMPOUNDS OF TIN. The following are the names and probable constitutional formulae of the principal compounds of this metal: Stannous chloride (at 900°), SnCl2. Stannic chloride, . . . SnCl4. Stannous oxide, .... SnO. Stannic oxide or anhydride, Sn02. Stannous oxydichloride", .• o=Sn=SnSn. * With regard to the fusing points of alloys, or of any mixtures of fusible substances which do not chemically combine, the law holds that the fusing point of the mixture is lower than the main fusing point of the constituents in the proportion in which they are present. 324 INORGANIC CHEMISTRY. f SnHo3 O SnHo2 Metastannic acid (dried at 100“ C.), . . . . .■> ® ° • SnHo2 O LSIIHO3 f SnHo.,K<> O Snllo., I O Dipotassic metastannate, -{ SnO O SnHo2 O SnHo2Ko Stannous sulphide, . . . SnS". Stannic sulphide, . . . SnS"2. S ( / \ Distannic trisulphide, . . gag//®", S=Sn—Sn=S S Stannous sulphostannate, . S=SnSn, or „ II o Stannous sulphate, . . . S02Sno". S2 + 2HCI = SnS" + 2KCI + SH2. Potassic Hydrochloric Stannic Potassic Sulphuretted sulphostannate. acid. sulphide. chloride. hydrogen. Stannic sulphide, SnS"2.—This compound cannot be prepared by merely heating tin and sulphur together. The addition of some volatile substance is necessary in order to lower the temperature during the reaction. An amalgam of 12 parts of tin and 6 parts of mercury is powdered, and heated with 7 parts of sulphur and 6 parts of am- nionic chloride in a glass retort. Ammonic chloride, mercury, and sulphur, along with mercuric sulphide and mercurous chloride, vola- tilize, and the stannic sulphide remains in the flask as a mass of golden- yellow flakes with a metallic lustre. It is not certain whether the ammonic chloride takes part in the reaction or whether it acts merely by its volatilization. Amorphous stannic sulphide is obtained as a brown precipitate by passing sulphuretted hydrogen into an acid solution of a stannic salt. After drying at ordinary temperatures, it still contains water of hydra- tion, with which it parts on heating. Amorphous stannic sulphide dissolves in hot concentrated hydro- chloric acid, and the solution contains stannic chloride. Hot concen- trated nitric acid also decomposes it. It is soluble in alkaline sulphides with formation of sulphostannates : SnS"2 + SK2 = SnS"Ks2; Stannic Potassic Potassic sulphide. sulphide. sulphostannate. COMPOUNDS OF TIN. 329 and in caustic alkalies with formation of a mixture of stannate and sulphostannate: 3SnS"2 + 60KH = SnOKo2 + 2SnS"Ks2 + 30H2. Stannic Potassic Potassic Potassic Water, sulphide. hydrate. stannate. sulphostannate. Crystalline stannic sulphide is insoluble in all single acids, but solu- ble in aqua-regia. Alkalies and alkaline sulphides also dissolve it. Both the varieties of stannic sulphide are decomposed at a bright red heat into free sulphur and stannous sulphide. Crystalline stannic sulphide is employed in the arts under the name of mosaic gold in the production of imitation bronze surfaces. It was known to the alchemists. Sulphostannates.—Only the alkaline sulphostannates are soluble in water. Potassic sulphostannate is uncrystallizable. The sodium salt, SnS//Na«2,70H2, crystallizes in yellow regular octahedra. D [STANNOUS TRISULPHIDE, Or STANNOUS SULPHOSTANNATE, or SnS^Sns."—This compound is prepared by heating to low redness a mixture of 3 parts of stannous sulphide and 1 part of sulphur. It forms a grayish-yellow mass with a metallic lustre. When treated with hot concentrated hydrochloric acid, one half of the tin goes into solution as a stannous salt, the other half remaining behind as stannic sulphide. This reaction would seem to denote that the sub- stance is not, as is frequently assumed, a distinct sulphide of tin, but a stannous sulphostannate. All the sulphides of tin are reduced to the metallic state when heated to redness in a current of hydrogen. General character and reactions of the salts of tin.— The stannous salts, when in solution, absorb oxygen from the air, and are converted into stannic salts. Caustic alkalies precipitate from the solutions white stannous hydrate, which is soluble in an excess of alkali. When an alkaline solution of stannous oxide is boiled, metallic tin separates out and an alkaline stannate remains in solution. Ammonia and the alkaline carbonates produce a precipitate of stannous hydrate, which is, however, not dissolved by an excess of the precipitant. With sulphuretted hydrogen in acid or neutral solutions, the whole of the tin is precipitated as brown stannous sulphide, almost insoluble in colorless ammonic sulphhydrate, readily soluble in yellow arnraonic sulphide. In alkaline solutions of stannous salts the precipitate is either not formed at all or else the precipitation is incomplete. With a solution of auric chloride the stannous salts yield, if added in small quantity, a purple precipitate of aurostannous stannate (Sn2o2Auo2- Sno'blOHjj), known ns, purple of Cassius; an excess of the stannous salt produces a brown precipitate of metallic gold. The stannic salts yield with caustic alkalies a white precipitate of stannic acid soluble in excess of alkali ;■ and the solution gives no pre- cipitate on boiling. With sulphuretted hydrogen a yellow precipitate of stannic sulphide is formed, soluble in alkalies and alkaline sulphides. 330 INORGANIC CHEMISTRY. TITANIUM, Ti. Atomic weight 48. Sp. gr. 5.3. Atomicity " and IT, also a pseudo-triad. Evidence of atomicity: Titanous oxide, . . Ti"G. Titanic tetrachloride, . . . Ti^Cl*. Dititanic hexachloride,. . . j TiCI,. * • | TlCIy History.—Titanium was discovered by Gregor in 1789. Occurrence.—Titanium is one of the rarer elements. It is never found in the free state. As titanic anhydride (Tio2) it occurs in three rare minerals—rutile, anatase, and brookite—and as ferrous titanate (TiOFeo") in titaniferous iron ore. Preparation.—1. Metallic titanium is most readily obtained by heat- ing potassic titanofluoride with potassium in a covered crucible: TiK2F6 + 2K2 = .Ti + 6KF. Potassic Potassic titanofluoride. fluoride. On dissolving the product of the reaction in water the titanium remains as a gray amorphous powder. 2. It may also be obtained in the form of prismatic crystals by heat- ing,sodium in the vapor of titanic chloride: TiCl4 + 2Na2 = Ti + 4NaCI. Titanic chloride. Sodic chloride. Properties.—Amorphous titanium forms a gray powder which, when heated in the air, or when thrown into a flame, burns with brilliant scintillations, forming titanic anhydride. At ordinary temperatures it does not decompose water, but at 100° C. hydrogen is evolved and titanic acid is formed : Ti + 30H2 = TiOHo2 + 2H2. Water. Titanic acid. It dissolves in hydrochloric and dilute sulphuric acids with evolution of hydrogen and formation of titanous salts. The following are the names and probable formulae of the chief com- pounds of titanium: Cl Titanic tetrachloride, TiCl4. Cl—Ti—Cl. I Cl 331 COMPOUNDS OF TITANIUM, Cl Cl Dititanic hexachloride, | TiCJ3, Cl—Ti—Ti—Cl. Cl Cl Titanous oxide, TiO. Ti=o. Titanic oxide or anhydride (Rutile, 1 Anatase, Brookite), . . .'. . j 2’ 0 1 Titanic acid, TiOHo2. ll—O—Ti—O—H. Titanic sulphide, TiS'^. Ti S % Titanic dinitride, IS- N. N Trititanic tetranitride, N—Ti Ti Ti=N. I COMPOUNDS OF TITANIUM WITH CHLORINE. Titanic Chloride, TiCl4. Molecular weight =l9O. Molecular volume I I I. Sp. gr. of liquid 1.76. Boils at 136° C. (276.8° F.). This substance is prepared by heating a mixture of titanic anhydride and finely divided carbon in a current of chlorine: TiOs + 2C + 2C12 = TiCl, + 200. Titanic Titanic Carbonic anhydride. chloride. oxide. It is a colorless strongly fuming liquid, which combines with a small quantity of water to form a crystalline compound, but is decomposed by an excess of water with separation of titanic acid. f TiCl Dititanic hexachloride, ■< ,1j3, is formed when a mixture of the vapor of the tetra- chloride with dry hydrogen is passed through a red-hot tube: 2TiCl4 + Hz = {™33 + 2HCI. Titanic Dititanic chloride. hexachloride. It forms dark violet scales, which cannot be re-sublimed without decomposition. It is deliquescent, and dissolves in water to form a violet solution, which absords oxygen from the air, and becomes colorless. 332 INORGANIC CHEMISTRY. COMPOUNDS OF TITANIUM WITH OXYGEN AND HYDROXYL. Titnnous oxide, TiO, has not been prepared in a state of purity. A hydrate, which has also not been isolated, is formed as a black precipitate when ammonia is added to the solution of a titanous salt prepared by dissolving titanium in a dilute acid. On boiling the liquid with the precipitate, the color of the latter changes to blue and ulti- mately to while, the oxide having been converted into titanic acid at the expense of the oxygen of the water, whilst hydrogen is evolved. Titanic oxide or anhydride, Ti02.—The hydrate of this oxide, tetrabasic titanic acid, Tildo4, is obtained as a white precipitate when ammonia is added to a solution of titanic chloride. This hydrate pos- sesses both basic and acid properties, combining both with acids and with alkalies. When dried in vacuo, it parts with the elements of one molecule of water, and is converted into the acid TiOHo2. At a higher temperature the rest of the water is eliminated, and titanic anhydride is left as a white amorphous powder, which on ignition becomes denser, and of a dark reddish-brown color. Titanic anhydride occurs in nature as rutile, crystallizing in reddish-brown quadratic prisms of sp. gr. 4.3 ; as anatase in quadratic pyramids, irreducible to the forms of rutile, and having a sp. gr. of 3,9; and as brookite in rhombic crystals of 4.1 sp. gr. Titanic anhydride is thus trimorphous. It may be obtained artificially in the same forms by passing a mixture of hydrochloric acid and steam over heated titanofluoride. At very high temperatures rutile is formed; at temperatures between the boiling-points of zinc and cad- mium, crystals of brookite are deposited; whilst below the boiling- point of cadmium anatase is obtained. Titanic anhydride is insoluble in alkalies, and in all acids except hydrofluoric and hot concentrated sulphuric. The titanates have not been thoroughly investigated. All the normal titanates are insoluble in water. Dititanic trioxide, Ti203, is obtained as a black powder by igniting titanic anhydride In a current of hydrogen. When heated strongly in air it is oxidized to titanic anhy- dride. Hydrochloric and nitric acids are without action upon it. Sulphuric acid dis- solves it, yielding a violet solution. COMPOUND OF TITANIUM WITH SULPHUR. Titanic sulphide, TiS//2, is formed when a mixture of the vapor of titanic chloride with dry sulphuretted hydrogen is passed through a red-hot tube : TiCl4 + 2SEI2 = TIS", + 4HCI. Titanic Sulphuretted Titanic Hydrochloric chloride. hydrogen. sulphide. acid. It forms brass-yellow scales resembling mosaic gold. It burns when heated in the air, yielding titanic and sulphurous anhydrides. By exposure to moist air it is slowly de- composed, with evolution of sulphuretted hydrogen. COMPOUNDS OF TITANIUM WITH NITROGEN AND WITH NITROGEN AND CARBON. Titanic dinitride, /N//2Ti, is obtained by heating titanic anhydride in a current of nitrogen : Ti02 + 2NH3 = 'N"sTi + 20 H2 + Hs. It is a dark violet-colored powder with a coppery tinge. ZIRCONIUM. 333 A second nitride, Ti3N4, trititanic tetranitride, is obtained in the form of a copper- colored metallic mass when the double compound of titanic chloride with ammonia (TiCI4,4NH3) is heated in a current of gaseous ammonia: 3TiCl4 + 4NH3 = Ti3N4 + 12HC1. Titanic Ammonia. Trititanic Hydrochloric chloride. tetranitride. acid. This compound was formerly mistaken for metallic titanium. When trititanic tetranitride is strongly heated in a current of hydrogen, a third nitride, Ti5N6, pentatitanic hexanitride, is produced in the form of golden-yellow scales, with a strong metallic lustre. All the nitrides of titanium, when heated with easily reducible oxides, such as those of copper, lead, and mercury, deflagrate brilliantly, the oxides undergoing reduction to the metallic state. Titanic cyanonitride.—TisN3(CN).—This remarkable compound, which was also formerly mistaken for metallic titanium, is frequently found in blast-furnaces which have been used for smelting titaniferous iron. It forms copper-colored metallic cubes, which are hard enough to scratch glass, and possess a sp. gr. of 5.3. The process by which this substance is formed may be imitated on a small scale by heating titanic an- hydride, mixed with charcoal, in a current of nitrogen : 5Ti02 + 11C + 2N2 = Ti5N3(CN) + 10CO. Titanic Titanic Carbonic anhydride. cyanonitride. oxide. It is insoluble in acids. Heated in a current of steam it yields titanic anhydride, ammonia, and hydrocyanic acid. Heated in chlorine, titanic and cyanic chlorides are formed, whilst nitrogen is liberated. General character and reactions of the titanium com- pounds.—The tltanous salts are unknown except in solution. With alkaline carbonates they yield a black precipitate, which becomes blue, and ultimately white. The alkaline titanates are of a yellowish color. They are insoluble in water, but soluble in hydrochloric acid. On boiling the hydrochloric acid solution, white titanic acid is precipitated ; ammonia produces the same effect. With raicrocosmic salt the titanates yield in the reducing flame of the blowpipe a violet glass which becomes colorless in the oxidizing flame. ZIRCONIUM, Zr Atomic weight = 90. Sp. gr. 4.15. Atomicity iv. Evidence of atom- icity : Zirconic chloride, ZrCl4. Zirconic fluoride, ZrF4. Occurrence.—ln combination with silicon and oxygen as zirconic silicate, it forms the rare mineral, zircon, SiZroiv. Preparation.—Zirconium is obtained by heating potassic zircono- fluoride with potassium : ZrK2F6 + 2K2 = Zr + 6KF. Potassic Potassic zirconofluoride. fluoride. On treating the mass with dilute hydrochloric acid the zirconium remains behind as a black amorphous powder. By employing alumin- 334 INORGANIC CHEMISTRY. mm to reduce the potassic zirconofluoride the zirconium may be obtained in crystalline plates. Reaction.—When heated in air. amorphous zirconium readily burns, forming zirconic oxide. The crystallized variety is oxidized only superficially, even at a white heat, but may be burnt with the aid of the oxyhydrogeu blowpipe. Zirconic chloride. ZrCfi (molecular volume f I Ih is prepared like titanic chloride (p. 331). It is a white crystalline mass, which, when treated with water, yields an oxy- chloride of the formula ZrOCl3,80H2. Zirconic bromide, Zrllr4, is also known, and resembles the chloride in its properties and reactions. Zirconic fluoride. ZrF4, is obtained by heating a mixture of zirconic oxide and fluor- spar to whiteness in a current of gaseous hydrochloric acid : Zr02 + 2CaF2 + 4HCI = ZrF4 + 2CaCI2 + 2QH2. Zirconic Calcic Hydrochloric Zirconic Calcic Water, oxide. fluoride. acid. fluoride. chloride. It is a colorless crystalline transparent substance, volatile at a white heat, and soluble in a solution of hydrofluoric acid. With the fluorides of the metal it forms zircono- fluorides, of which the most important is potassic zirconofluoride, ZrK2F6. Zirconic oxide, zirconia, Zr02, is formed by burning zirconium in air, or by heating the hydrate. It is a white infusible powder. When heated in the oxyhydrogen blow- pipe it emits a very intense light. Concentrated sulphuric acid dissolves it with diffi- culty. When fused with alkaline carbonates, it expels carbonic anhydride, and com- bines with the base to form a zirconate. On treating the fused mass with water, the zirconate is decomposed, and zirconic hydrate, ZrHo4, separates out as a voluminous pre- cipitate. The same precipitate is obtained by adding ammonia to the cold solution of a salt of zirconium. It dissolves readily in dilute acids. When ammonia is added to a hot solution of a zirconium salt a hydrate of the formula ZrOHo2 is precipitated. This second hydrate dissolves with difficulty in acids. The method of fusing with an alkaline carbonate is employed in obtaining zirconia from its minerals. Atomic weight 233.4. Sp. gr. 11.23. Atomicity,v. THORIUM, Th, Occurrence.—This substance is of even rarer occurrence than zirconium. It is a con- stituent of the very rare minerals thorite, monazite, and euxenite. Preparation.—It may be obtained as a dark gray powder by heating thoric chloride with potassium or sodium. The following are some of its principal compounds: Thoric chloride, ThCl4. Thoric fluoride, ThF4,40H2. Potassic thorofluoride, ThK2F6,2 d Thoric oxide, thoria, Th02. Thoric silicate (thorite) SiTho,v,2oH2. PHOSPHORUS. 335 CHAPTER XXX. PENTAD ELEMENTS. Section" I. {Continuedfrom Chapter XXVI.). PHOSPHORUS, P4. Atomic weight =3l. Molecular weight = 124. Molecular volume 1 I I. 1 litre of 'phosphorus vapor weighs 62 criths. Sp. gr. 1.83. Fuses at 44-45° C. (111-113° F.). Boils at 290° C. (554° F.). Atomicity and v. Evidence of atomicity : Phosphorous hydride, . . . , . P"'H8. Phosphorous chloride, . . . . . P///C13. Phosphoric chloride, . . . . . . PTC1-. Phosphonic iodide, . . . PvH4L Phosphoric fluoride, . . . • • • PtF5. History.—Phosphorus was discovered in 1669 by Brand, an alche- mist of Hamburg, who obtained it by evaporating urine to dryness, and distilling the residue with sand. The process was kept secret; but in 1680 Boyle succeeded in preparing phosphorus, employing the same method. In 1769 Gahn showed that calcic phosphate is a constituent of bones, and in 1771 Scheele published a method of obtaining phos- phorus from this source. Occurrence.—Phosphorus is never found in the free state in na- ture. It generally occurs combined with oxygen and a metal to form a phosphate. The principal naturally occurring phosphates are os- PO teolite (estramadurite, sombrerite) or calcic phosphate, and apatite or calcic chlorophosphate, (PO)///3Cao"4(OCaCI). Calcic phos- phate is widely distributed in small quantities as a constituent of the primitive rocks, by the disintegration of which it passes into the soil. From the soil the phosphorus is absorbed by plants, where it accumulates chiefly in the seed. From plants it passes into the bodies of animals, in which it is found in still greater quantity. Calcic phos- phate forms the chief inorganic constituent of the bones, whilst phos- phorus in complex organic combinations is always present in the sub- stance of the nerves and brain, and in smaller quantity in the other tissues. In the slow oxidation of the living animal substance which is constantly going on, the phosphorus is eliminated in the urine as phos- phates of sodium, potassium, and magnesium. Preparation.—l. Calcined bones, which consist of calcic phosphate with a slight admixture of calcic carbonate, are digested with sufficient sulphuric acid to decompose the whole of the carbonate and two-thirds of the phosphate. In this way the tricalcic diphosphate is converted into tetrahydric calcic diphosphate : P202Cao"3 + 2S02Ho2 = P202Ho4Cao" + 2SO20ao". Tricalcic diphos- Sulphuric Tetrahydric calcic Calcic sulphate, phate (Bone-ash). acid. diphosphate. 336 INORGANIC CHEMISTRY. The tetrahydric calcic diphosphate is extracted with water from the calcic sulphate, evaporated to a syrup, mixed with charcoal, and heated to dull redness in an iron pot, stirring all the time. Under the influ- ence of heat the tetrahydric calcic diphosphate parts with water, and is converted into calcic metaphosphate, which is thus obtained intimately mixed with charcoal: P202Ho4Cao" = P204Cao" + 20H2. Tetrahydric calcic Calcic Water, diphosphate. metaphosphate. The mixture is then transferred to earthenware retorts and heated to bright redness, when the following reaction takes place: 3P204Cao" + lOC = P202Cao"3 + 10CO + P4. Calcic Tricalcic Carbonic metaphosphate. diphosphate. oxide. The phosphorus distils over, and is collected under water, whilst the carbonic oxide escapes carrying with it a small quantity of phosphorus Fig. 46. vapor, which causes it to inflame on coming in contact with the air. The apparatus employed in this distillation varies in different factories; one form is shown in Fig. 46. PHOSPHORUS. 337 The explanation of the process is as follows : Normal salts of tribasic phosphoric acid are not acted upon when heated with charcoal, but phosphoric anhydride, under these circumstances, is readily reduced. If we regard a salt as a compound of anhydride and base, it will be seen that the salts of monobasic phosphoric acid contain more anhydride in proportion than the tribasic acid. Thus: SPACao" = PACao"s + 2P,0,. The reduction takes place to the extent of the excess of anhydride above what is necessary for the formation of tricalcic diphosphate. Accord- ingly, in the above process two-thirds of the phosphorus present are reduced. Sombrerite, an impure calcic phosphate found in the West Indies, is frequently substituted for bone-ash. 2. If sand be added to the mixture in the above distillation, calcic silicate is formed, and the whole of the phosphorus is expelled (Woh- ler) : 2PACao" + 10C + 2Si02 = 2SiOCao" + 10CO + P4. Calcic Silicic Calcic Carbonic metaphosphate. • anhydride. silicate. oxide. 3. If a mixture of bone-ash and charcoal be heated to redness in a current of gaseous hydrochloric acid, the whole of the phosphorus is liberated, and calcic chloride remains (Cary-Montrand): PACao" + C 8 + 6HCI = 3CaCI2 + BCO + 3II2 + P2 Tricalcic Hydrochloric Calcic Carbonic diphosphate. acid. chloride. oxide. This process has not, however, proved successful on a manufacturing scale. The crude phosphorus is always contaminated by particles of charcoal and other impurities carried over during the distillation. From these it is freed, either by fusing it under water and pressing it through washTleather bags, or by partially oxidizing it with a mixture of po- tassic dichromate and sulphuric acid. The oxidation is attended with effervescence, which causes the impurities to rise to the surface, leaving the phosphorus pure. The purified phosphorus is cast into sticks. Properties.—Phosphorus exists in several allotropic modifications. Common or octahedral phosphorus, the modification obtained in the processes above described, is, when freshly prepared, a colorless trans- parent solid. Very frequently, however, it displays a faint yellowish tinge due to the presence of some impurity. It has a sp. gr. of 1.83. It is a non-conductor of electricity. At ordinary temperatures it may be cut with a knife like wax, but about 0° C. it becomes brittle. At a temperature of 44-45° C. (111-113° F.) it fuses to a colorless oily liquid, which readily retains its fluidity several degrees below its so- lidifying point. It boils at 290° C. (554° F.). The molecular weight of phosphorus, deduced from the vapor-density, is 124, showing that 338 INORGANIC CHEMISTRY. the molecule of phosphorus consists of four atoms, and this tetratomic molecule does not break up even at a temperature of 1040° C. (1840° F.) (Deville and Troost); but at a higher temperature, the vapor- density has a value lying between the values required for P2 and P4 re- spectively, showing that a partial dissociation has taken place (Victor Meyer), Phosphorus is a very inflammable substance, igniting in the air a few degrees above its fusing-point. For this reason it must always be pre- served and cut under water. Under the influence of air and light it becomes covered, when kept under water, with a white opaque crust, due to a partial oxidation. It ought therefore to be kept in the dark. When exposed to the air at ordinary temperatures phosphorus under- goes slow oxidation, and gives off a white vapor, which has a powerful odor of garlic. In a dark room both the phosphorus and the vapor are luminous with a greenish-white light. At a few degrees below 0° C. the oxidation and the luminosity cease. In pure oxygen under ordi- nary pressures phosphorus is not luminous at temperatures below 15° C.; but by rarefying the oxygen, or adding some inactive diluent, such as nitrogen, hydrogen, or carbonic anhydride, the phosphorus again becomes luminous. The luminosity of phosphorus in air is also pre- vented by the presence of minute traces of certain gases or vapors, such as olefiant gas, sulphuretted hydrogen, and turpentine.* When phos- phorus is exposed to the air in large quantities, the heat of oxidation is frequently sufficient to melt, and finally to ignite, the mass. The same effect is produced by exposing phosphorus to the air in a finely divided condition, so as to increase the oxidizable surface. This may be shown by pouring a solution of phosphorus in carbonic disulphide upon filter- ing paper, and allowing the liquid to evaporate. In the dark the paper becomes brightly luminous, and at last bursts into flame. Phosphorus is insoluble in water, slightly soluble in ether, turpen- tine, and benzine, readily soluble in disulphur dichloride, phosphorous chloride, and carbonic disulphide. One part by weight of the latter solvent dissolves from seventeen to eighteen parts of phosphorus. By the spontaneous evaporation of this solution it may be obtained in trans- parent crystals belonging to the regular system, generally octahedra or rhombic dodecahedra. When phosphorus is kept in the dark in sealed vacuous tubes, it*spontaneously sublimes, and is deposited on the sides of the tubes in very lustrous and perfect crystals. Phosphorus may be finely granulated by melting it under water, and agitating until it solidifies again. The addition of a small quantity of urea to the water prevents the adhesion of the granules, and by this means a higher degree of subdivision is attained. Phosphorus is an exceedingly poisonous substance. Even the fumes have a very deleterious action when inhaled, producing caries of the bones of the jaw. jßed or Amorphous Phosphorus.—This variety was discovered by Schrotter in 1845. It is formed when ordinary phosphorus is exposed * According to Chappuis, the luminosity of phosphorus depends upon the presence of ozone. Substances which destroy ozone prevent the luminosity. PHOSPHORUS. 339 to the action of the heat or light in an atmosphere devoid of oxygen. It is best prepared by heating phosphorus for some time in a closed vessel to 230-250° C. (446-482° F.). On a manufacturing scale, iron vessels are employed for this purpose, and it is not necessary to fill the apparatus with any artificial atmosphere, as the oxygen is speedily removed from the air by the combustion of. a small portion of the phosphorus. Any rise of temperature above 250° C. must be care- fully avoided, since at 260° C. (500° F.) amorphous phosphorus is re- converted into the ordinary modification, the change being accompanied with evolution of heat and taking place, in the case of large quantities, With explosive violence. Amorphous phosphorus is, however, formed when ordinary phosphorus is heated under pressure in closed iron vessels to 300° C. (572° F.), the change taking place in a few minutes. When ordinary phosphorus is heated with a small quantity of iodine or selenium, an iodide or selenide is formed, and the excess of phos- phorus is instantaneously converted into the red variety. Amorphous phosphorus, prepared by any of the above methods, in- variably contains a small quantity of white phosphorus, the presence of which renders the product dangerously inflammable. From this it may be freed by grinding the crude amorphous phosphorus under water, and subsequently treating it with carbonic disulphide, which dissolves the unaltered phosphorus, or still more advantageously by boiling with caustic soda (see Phosphoretted Plydrogen). Thus purified, amorphous phosphorus forms a reddish-brown powder of sp. gr, 2,15. It is de- void of taste and smell, is not poisonous, may be exposed to the air for any length of time without undergoing change, and is not luminous in the dark. When heated it does not fuse, and inflames in the air only at a temperature of 260° C. (500° F.), being converted at the same tune into ordinary phosphorus. It is insoluble in the solvents which dissolve ordinary phosphorus, such as carbonic disulphide and sulphur chloride. It conducts electricity feebly. Rhombohedrql Phosphorus.—This variety is obtained when phos- phorus is heated with metallic lead in sealed tubes for eight or nine hours to a temperature below redness. On dissolving the cooled lead ln dilute nitric acid, small, well-defined, violet-black rhorabohedra, having a sp. gr. of 2.34, remain. This modification may also be ob- tained by heating: amorphous phosphorus under pressure to 580° C. (1076° F.). According to some chemists red phosphorus and rhombohedral phos- phorus are identical. A fourth modification, obtained as a black mass by quickly cooling melted phosphorus, has been described; but it has been shown that this substance is produced only when metals are present, the color being due to the formation of metallic phosphides. Reactions.—Owing to its affinity for oxygen, phosphorus acts as a powerful reducing agent. Platinum, gold, silver, and copper are de- posited in the metallic state, when white phosphorus is left in contact with the solutions of their salts. When sodic carbonate is heated to redness with phosphorus, the carbonic anhydride is reduced and car- bon is set free. When dry finely divided phosphorus is mixed with 340 INORGANIC CHEMISTRY. substances which readily part with oxygen, such as potassic chlorate or metallic peroxides, very slight friction is sufficient to cause the explo- sive oxidation of the phosphorus. The other reactions of phosphorus will be described in connection with its compounds. Uses.—Phosphorus is employed chiefly in the manufacture of lucifer matches. In the commoner sorts, the matches are tipped first with sulphur, and then with a mixture of phosphorus and potassic chlorate made into a paste with glue. They ignite by friction on any rough surface. The sulphur serves to transmit the combustion from the phos- phorus to the wood. Nitre is frequently substituted for potassic chlo- rate, as the matches thus prepared ignite more quietly; whilst, in order to get rid of the disagreeable smell of burning sulphur, this substance is replaced by paraffin. In the safety matches the phosphorus is sepa- rated from the other inflammable materials. The matches are tipped with a mixture of potassic chlorate, potassic dichroraate, red lead, and antiraonious sulphide, and are ignited by friction on a prepared surface coated with amorphous phosphorus and antimonious sulphide. These matches do not readily iguite on an unprepared surface, but by rubbing them rapidly over a smooth slate, or a sheet of ground glass, they may be inflamed. COMPOUNDS OF PHOSPHORUS WITH HYDROGEN. Phosphorus forms with hydrogen three compounds. These cannot be obtained by the direct combination of their elements. Solid phosphoretted hydrogen, . . • • \ P(P'"H)"- • • 'P"2H4. Liquid “ “ Gaseous “ “ . . . . ph3. GASEOUS PHOSHPORETTED HYDROGEN. Phosphine. H PH, | H—P—H Molecular weight 34. Molecular volume I I I. 1 litre weighs 17 criths. Preparation.—l. Phosphoretted hydrogen may be obtained by heating hypophosphorous acid: 2PHHo2 = PH3 -f- POPIO3. Hypophosphorous Phosphoretted Phosphoric acid. hydrogen. acid. 2. A similar decomposition occurs when phosphorous acid is heated : GASEOUS PHOSPHORETTED HYDROGEN. 341 4PHo3 = PH3 -f 3POHo3. Phosphorous Phosplioretted Phosphoric acid. hydrogen. acid. 3. When phosphorus is heated with a solution of sodic or potassic hydrate, phosphoretted hydrogen is evolved, whilst an alkaline hypo- phosphite remains in the retort; 30NaH + P4 + 30H2 = 3PHHoNao + PH3. Sodic hydrate. Water. Sodic Phosphoretted hypophosphite. hydrogen. The gas prepared by this process contains free hydrogen and liquid phosphoretted hydrogen, the presence of this latter substance render- ing the gas spontaneously inflammable in contact with air. By employ- ing an alcoholic solution of caustic alkali, a gas is obtained which does Rot inflame spontaneously, the liquid phosphoretted hydrogen remain- ing in this case dissolved in the alcohol. 4. Phosphoretted hydrogen is evolved when calcic phosphide is treated with water: P2Ca3 -f- 60H2 2PH3 -}- 3CaHo2. Tricalcic Water. Phosphoretted Calcic diphosphide. hydrogen. hydrate. The gas is also in this case contaminated with the vapor of liquid phosphoretted hydrogen. 5. Pure phosphoretted hydrogen is most readily obtained by allow- ing concentrated caustic potash to drop very gradually upon phosphonic iodide (q.v.) contained in a flask: PHJ + OKH = PH3 + KI + OH2. Phosphonic Potassic Phosphoretted Potassic Water, iodide. hydrate. hydrogen. iodide. Properties—Phosphoretted hydrogen is a colorless gas possessing an odor resembling that of garlic. It is combustible in air or oxygen, burning with a very brilliant white light, and evolving a cloud of phos- phoric acid. When pure it is not spontaneously inflammable; but the presence of a small quantity of the vapor of liquid phosphoretted hydrogen (T"2H4) in the gas suffices to impart to it this property, of which it may again be deprived by leaving it in contact with finely di- vided charcoal, which absorbs the liquid compound, or by exposing it to the action of sunlight, by which the liquid compound is decomposed. On the other hand, the pure gas may be rendered spontaneously inflam- mable by the addition of a trace of nitrous anhydride. If the pure gas be mixed with oxygen no action is observed; but, on suddenly rarefying the mixture, combination takes place with explosion. This phenomenon is possibly allied to that of the luminosity of phos- phorus in rarefied oxygen. If the spontaneously inflammable gas be allowed to bubble through 342 INORGANIC CHEMISTRY. water, each bubble, on escaping into the air and inflaming, forms a smoke-ring of phosphoric acid. Phosphoretted hydrogen is a highly poisonous gas. When inhaled, even in a very diluted condition, it produces difficulty in breathing, and ultimately death. Reactions.—l. By combustion in oxygen it yields raetaphosphoric acid and water: PH3 + 202 = P02Ho + 0H2. Phosphoretted Metaphosphoric Water, hydrogen. acid. 2. In contact with chlorine it forms phosphoric chloride and hydro- chloric acid: PH3 + 4C12 = PC15 + 3HCI. Phosphoretted Phosphoric Hydrochloric hydrogen. chloride. acid. 3. When passed through a solution of cupric sulphate, it produces a black precipitate of cupric phosphide : 2PH3 + 3S02Cuo" = P2Cu"3 + 3S02Ho2. Phosphoretted Cupric Cupric Sulphuric hydrogen. sulphate. phosphide. acid. 4. When passed through a solution of argentic nitrate, metallic silver is deposited, whilst nitric and phosphoric acids are formed: PH3 8N02Ago + 40H2 = POHo3 Phosphoretted Argentic Water. Phosphoric hydrogen. nitrate. acid. -f 4Ag2 + 8N02Ho. Nitric acid. 5. It unites directly with hydrochloric, hydrobromic, and hydriodic acids, when the dry gases are brought together, forming compounds analogous to the haloid salts of ammonium : PH3 + HBr = PH4Br. Phosphoretted Hydrobromic Phosphonic hydrogen. acid. bromide. Phosphoretted hydrogen and hydrochloric acid unite only under the influence of pressure and cold (Ogier). Phosphonic iodide is also formed by the action of iodine on phos- phoretted hydrogen. The reaction takes place in two stages: PHS + 3I2 = PI, + SHI; Phosphoretted Phosphorus Hydriodic hydrogen. iodide. acid. and PH3 + HI = PHJ. Phosphoretted Hydriodic Phosphonic hydrogen. acid. iodide. Phosphonic iodide is, however, most conveniently prepared by the following method (A. W. Hofmann): 10 parts of phosphorus are dis- LIQUID PHOSPHORETTED HYDROGEN. 343 solved in carbonic disulphide in a retort, and 17 parts of iodine are gradually added, cooling during the operation. The carbonic disul- phide is then distilled off, a stream of dry carbonic anhydride being finally passed through the apparatus to remove the last traces of the carbonic disulphide, and 6 parts of water are very slowly added by means of a dropping-funnel. A violent reaction takes place, the heat of which volatilizes the phosphonic iodide as it is formed. Towards the close heat is applied to the retort. A slow stream of carbonic anhydride must be passed through the apparatus during the whole operation, in order to prevent the entrance of air, which might other- wise occasion an explosion. The phosphonic iodide condenses in large lustrous quadratic crystals in a wide tube attached to the neck of the retort. The following equation expresses the reaction: 13P + 91 + 210H2 = 7PHJ + 2HI + 3P203H«4. Water. Phosphonic Hydriodic Pyrophosplioric iodide. acid. acid. Phosphonic iodide is employed in the laboratory as a powerful re- ducing agent, available particularly at high temperatures. Composition.—When a series of electric sparks is passed through phosphoretted hydrogen, it is gradually decomposed into its elements. The spark should pass between carbon points, since, when platinum is employed, a fusible phosphide of platinum is formed, which melts, putting an end to the experiment. It is found that two volumes of phosphoretted hydrogen yield three volumes of hydrogen when thus treated. Expressed in litres: 2 litres of phosphoretted hydrogen weigh . . . 34 criths. Deduct weight of 3 litres of hydrogen, . . . 3 “ There remain, . 31 “ which is the weight of \ litre of phosphorus vapor. Therefore | vol- ume of phosphorus vapor in combination with 3 volumes of hydrogen yields 2 volumes of phosphoretted hydrogen, or 31 parts by weight of phosphorus combine with 3 parts by weight of hydrogen to form this compound, and its formula is, therefore, PH3. LIQUID PHOSPHORETTED HYDROGEN. H H fPR I I 'P"2H4 or { P-P H H Molecular weight = 66. Molecular volume I 11. 1 litre of the vapor weighs 33 criths. Preparation.—This compound is formed along with gaseous phos- phoretted hydrogen by the action of water at a temperature of 60—70° 344 INORGANIC CHEMISTRY. C. (140-158° F.) on calcic phosphide obtained by passing the vapor of phosphorus over lime heated to redness (see Calcic Phosphide). This latter substance probably contains, in addition to calcic pyrophosphate, a mixture of dicalcic and tricalcic diphosphide (P2Ca//3), and from these two phosphides the liquid and gaseous phosphoretted hydro- gens are respectively formed : 'P"Ca"2 + 40 U2 = T"2H4 + 2CaHo2. Dicalcic Water. Liquid phos- Calcic diphosphide. phoretted hydrogen. hydrate. (For the formation of gaseous phosphoretted hydrogen from tricalcic diphosphide, see p. 341.) The gas evolved is passed through a U-tube immersed in a freezing mixture, and in this the liquid compound con- denses. Properties.—lt is a colorless, powerfully refracting liquid which inflames instantly in contact with air. Reaction.—By exposure to sunlight, or by contact with hydrochloric acid, it is decomposed into solid and gaseous phosphoretted hydrogens : 5/po h 6PH 4- f P(P'"H)"9 2n4 OJrrl3 I ) P(P///H)/r ’ Liquid phospho- Gaseous phospho- Solid phospho- retted hydrogen, retted hydrogen. retted hydrogen. The hydrochloric acid suffers no change. A very small quantity of the acid therefore suffices to decompose a practically unlimited quantity of the phosphorus compound. SOLID PHOSPHORETTED HYDROGEN f P(P"'H)"„ t P(P///H)// • Molecular weight = 126 ? Preparation.—Solid phosphoretted hydrogen is obtained by dissolving calcic phos- phide in concentrated hydrochloric acid, or by the action of light upon the liquid phos- phoretted hydrogen. Properties.— It forms a yellow powder which turns darker on exposure to light. When strongly heated in an atmosphere of carbonic avhydride, it is decomposed into its elements. It is doubtful whether this substance has ever been prepared in a state of purity, and its exact composition is uncertain. COMPOUNDS OF PHOSPHOR US WITH THE HALOGENS. Phosphorous chloride, . . . . ...... PC13. Phosphoric chloride, . PC15. Phosphorous bromide, . . . . ..... PBr3. Phosphoric bromide, . . . . ...... PBiV Diphosphorous tetriodide, . . . J 1 PI,* Phosphorous iodide, . . . . , PIS. Phosphoric fluoride, . . . . PF5. PHOSPHOROUS CHLORIDE—PHOSPHORIC CHLORIDE. 345 PHOSPHOROUS CHLORIDE Cl PC13. I Cl—P—Cl Molecular weight 137.5. Molecular volume I I I. 1 litre of phos- phorous trichloride vapor weighs 68.75 criths. Sp.gr. 1.613. Boils at 76° C. (168.8° F.). Preparation.—This compound is obtained by heating amorphous phosphorus in a retort while a current of dry chlorine is passed over it through the tubulure. The phosphorous chloride distils off as fast as it is formed, and collects in a cooled receiver. In order to free it from pentachloride, it is redistilled over ordinary phosphorus. Properties.—Phosphorous chloride is a colorless fuming liquid with a very pungent odor. It does not solidify at —lls° C. (—l7s° F.). Reactions.—l. With water it yields hydrochloric and phosphorous acids: PC13 + 30H2 = 3.HCI + PHo3. Phosphorous Water. Hydrochloric Phosphorous chloride. acid. acid. 2. With sulphuretted hydrogen it forms hydrochloric acid and phosphorous sulphide: 2PC13 + 3SH2 = 6HCI + P2S"s. Phosphorous Sulphuretted ■ Hydrochloric Phosphorous chloride. hydrogen. acid. sulphide. PHOSPHORIC CHLORIDE Cl I PC15. ci—P—Cl y\ Cl Cl Molecular weight = 208.5. Molecular volume I I I. 1 litre of undis- sociated phosphoric chloride vapor weighs 104.25 criths. Volatilizes below 100° C. Preparation.—Phosphoric chloride Is formed by the direct union of the trichloride with chlorine. A stream of dry chlorine is passed on to the surface of the trichloride contained in a flask cooled by water. Creat heat is evolved in the reaction. The liquid ultimately solidifies to a crystalline mass. Properties.—Phosphoric chloride is a crystalline powder with a faint yellowish tinge. It fumes in contact with moist air, and possesses a 346 INORGANIC CHEMISTRY. very irritating odor. It sublimes readily, but cannot be fused under ordinary pressure. In a sealed tube, under the pressure of its own vapor, it fuses at 148° C. (298.4° F.), and on cooling, solidifies in pris- matic crystals. At higher temperatures it possesses a vapor-density only half as great as is required for the molecular weight corresponding to the formula PC15, the reason of this being that the compound under- goes dissociation into PCJ3 and Cl2 (Introduction, p. 64). This disso- ciation is only partial at lower temperatures, and its progress may be traced by means of the change of color which the vapor undergoes as the temperature rises, phosphoric chloride yielding a colorless vapor which becomes yellowish-green as the proportion of free chlorine increases. This dissociation is to a great extent checked by allowing the phosphoric chloride to volatilize in an atmosphere of phosphorous chloride vapor, and in this way Wurtz determined the vapor-density of phosphoric chloride with a result closely agreeing with the normal density required for the formula PC15. Reactions.—l. A small quantity of water converts it into phosphoric oxytrichloride with formation of hydrochloric acid : PC15 + 0H2 == POCl3 + 2HCI. Phosphoric Water. Phosphoric Hydrochloric chloride. oxytrichloride. acid. 2. With an excess of water, it yields phosphoric and hydrochloric acids: PC15 + 40H2 = POHo3 + SHCI. Phosphoric Water. Phosphoric Hydrochloric chloride. acid. acid. 3. By its action on alcohols and acids, the chlorides of the radicals of the alcohols and acids are obtained, thus: {chJho + PC1* = {olci + HCI + POC!- Ethylic Phosphoric Ethylic Hydrochloric Phosphoric alcohol. chloride. chloride, acid. oxytrichloride. {coho + PC!» = jcoci + HCI + POC,s- Acetic Phosphoric Acetylic Hydrochloric Phosphoric acid. chloride. chloride. acid. oxytrichloride. 4. When phosphoric chloride acts on organic compounds containing oxygen attached with both its bonds to the same atom of carbon, a direct exchange of one atom of oxygen for two atoms of chlorine is effected: {cOH + PCI‘ = {cOUI + POC1*- Benzaldehyde. Phosphoric Benzalchloride. Phosphoric chloride. oxytrichloride. PHOSPHORIC FLUORIDE. 347 These properties render phosphoric chloride an invaluable agent in the investigation of organic compounds. Phosphorous bromide, PBrs (molecular volume [ | |), is prepared by the action of bro- mine on amorphous phosphorus. It forms a fuming colorless liquid of sp. gr. 2.925 at 0° C., boiling at 175° C. (347° F.). Its chemical behavior is analogous to that of the chloride. Phosphoric bromide, PBr5, is obtained by the direct union of the tribromide with bromine. It is a yellow crystalline solid which melts to a red liquid, and is decom- posed at 100° C. into the tribromide and free bromine. Its reactions resemble those of the corresponding chloride. I){phosphorous tetriodide, /P//2I4 (molecular volume \ | I), is prepared by dissolving 5 parts of phosphorus in carbonic disulphide, and gradually adding 41 parts of iodine, cooling well with water during the operation. On concentrating the solution by distill- ing off the carbonic disulphide, diphosphorous tetriodide crystallizes out in orange- colored prisms fusing at 110° C. (230° F.). Water decomposes it with formation of hydriodic and phosphorous acids and liberation of phosphorus in the amorphous con- dition : 3'P"2I4 + t120H2 = 12HI + 4PHo3 + P2. Diphosphorous Water. Hydriodic Phosphorous tetriodide. acid. acid. Phosphorous iodide. Pis, is obtained in the same manner as the foregoing compound, but employing 12 parts of iodine to lof phosphorus. It forms dark-red, deliquescent crystals, fusing at 55° C. (131° F.). It cannot be distilled without decomposition. By the action of water it yields hydriodic and phosphorous acids: Pig + 30 H2 = 3HI + PHOg. Phosphorous Water. Hydriodic Phosphorous iodide. acid. acid. PHOSPHORIC FLUORIDE. F f pf5. f—p—f A F F Molecular weight = 126. Molecular volume t I I. 1 litre of 'phosphoric fluoride weighs 63 criths. Preparation.—This compound is formed when arsenious fluoride is added to phosphoric chloride: 5AsF3 ’+ 3PC15 = SAsC13 + 3PF5. Arsenious Phosphoric Arsenious Phosphoric fluoride. chloride. chloride. fluoride. Properties.—Phosphoric fluoride is a colorless gas which fumes in contact with moist air, and possesses a very irritating odor. It is not inflammable. It is not decomposed by a series of electric sparks, either when the pure gas is employed, or when it is mixed with oxygen or hydrogen. Reactions.—l. Water decomposes it, forming phosphoric and hydro- fluoric acids: 348 INORGANIC CHEMISTRY. PP5 + 40H2 = POHo3 + SHF. Phosphoric Water. Phosphoric Hydrofluoric fluoride. acid. acid. 2. It unites with dry ammonia, forming a white solid compound of the formula 2PF5,5NH3, Phosphoric fluoride is particularly interesting as an example of the union of pentadic phosphorus with five monad atoms to form a com- pound capable of existing in the gaseous state, and even of sustaining very high temperatures without dissociation. COMPOUNDS OF PHOSPHORUS WITH OXYGEN AND HYDROXYL. H I Hypophosphorous acid, . . PHHo2. H—O—P—O—H 0 O II II Phosphorous anhydride, . . P203? P—O—P? H 1 O I Phosphorous acid, .... PIIo3. H—O—P—O—II O O II II Phosphoric anhydride, . . P.,0 . P—O—P II II O O 0 Phosphoric acid (tribasic), . P0110.,. lI—O—P—O—II i o Metaphosphoric acid (mono- ) q jj basic), j 2 °* || ~ _ O COMPOUNDS OF PHOSPHORUS WITH OXYGEN AND HYDROXYL. 349 0 o Pyrophosphoric acid (tetra-Ipo H H_o_p_o_P_o-H DaSlCjj J ' o o 1 I II II Hexabasic phosphoric acid, . Pt07Ho6. Sodium salt (Fleitmann and I Henneberg) (Hexasodic te- > P407Xaof traphosphate), j 0 o o o II II II II Na—O—P—O—P—O—P—O—P—O—Na 1 I I I o o o o till Na Na Na Na Dodecabasic phosphoric acid, PlOO19110i2. Sodium salt (Fleitmann I and Henneberg) (Dode- V PlOO19Hao12. casodic decaphosphate) . j o—O ~ O 11 . li II Na—O—P O—P O—P—O—Na 1 I I 0 o o 1 1 I Na Na 8 Na Phosphorosophosphoric acid p q {Hypophosphoric acid), . f 44 8 H H I I 0 O 1 1 lI—O—P—O—F-O—P—O—P—O—II /\ /\ 0o o o 1I I 1 II lIH H O 350 INORGANIC CHEMISTRY. HYPOPHOSPHOROUS ACID. PHHo2. Molecular weight =66. Fuses at 17.4° C. (63.3° F.). Preparation.—When phosphorus is heated with a solution of baric hydrate, phosphoretted hydrogen is evolved and baric hypophosphite is formed: 3BaHo2 + 2P4 + 60H2 = 3PHHoBao// + 2PH3. Baric Water. Baric Phosphoretted hydrate. hypophosphite. hydrogen. Any phosphoric acid which is formed at the same time combines with the barium to form insoluble baric phosphate, which may be removed by filtration. To the solution of baric hypojfiiosphite a quantity of dilute sulphuric acid exactly sufficient to precipitate the barium is added, and in this way a solution of hypophosphorous acid is obtained. The clear solution is evaporated over a flame, without, however, allow- ing it to boil, until the temperature rises to 130° C. (266° F.). On cooling to 0° C. the liquid thus obtained, hypophosphorous acid is de- posited in crystals. Properties.—Hypophosphorous acid forms white laminae fusing at 17.4° C. (63.3° F.). Reactions.—1. When strongly heated, hypophosphorous acid is de- composed into phosphoric acid and phosphoretted hydrogen : 2PHHo2 = POHo3 + PH3. Hypophosphorous Phosphoric Phosphoretted acid. acid. hydrogen. 2. It readily absorbs oxygen from the air, and is ultimately con- verted into phosphoric acid: PHHo2 + 02 = POHos. Hypophosphoroua Phosphoric acid. acid. Its affinity for oxygen causes it to act as a powerful reducing agent. It precipitates many of the metals in the metallic state from the solu- tions of their salts and, when heated with sulphuric acid, reduces it to sulphurous acid, and even to sulphur. Hypophosphites.—Hypophosphorous acid is a very weak acid, and although it contains two semi-molecules of hydroxyl, its acid power is exhausted as soon as the hydrogen of one of these is replaced by a metal. It therefore acts as a monobasic acid (cf. Orthophosphates). The hypophosphites are all soluble in water, and some are crystallizable. They, exhibit the same reducing properties as the free acid, and undergo a similar decomposition on heating. PHOSPHOROUS ANHYDRIDE—PHOSPHOROUS ACID. 351 PHOSPHOROUS ANHYDRIDE. PA (?)• Molecular weight = 110 (?). Preparation.—When phosphorus is gently heated in a slow current of dry air, it burns with a greenish flame, forming a compound having the composition of an anhydride of phosphorous acid. Properties.—This compound is a white amorphous fusible powder which may be sublimed. It has an odor of garlic. Peactions.—By allowing the above compound to deliquesce, with exclusion of oxygen, carefully avoiding any rise of temperature, a yellow solution is obtained which has a neutral reaction, and may, by dialysis, be proved to contain a colloid. If the solution be now heated, a reddish substance of unknown composition separates, and the solution contains phosphorous acid, PHo3. When the so-called anhydride is dissolved in water in the ordinary way, the temperature rises so high as to bring about the above decomposition at once, and a solution of phosphorous acid is obtained with separation of the reddish substance. From the above, it is probable that the compound obtained when phosphorus is burnt in a limited supply of air is not the true anhydride of phosphorous acid, but a compound of the same composition with a higher molecular weight (compare the molecular weights of arsenious anhydride and antimonious anhydride). The hydrate which this com- pound forms is neutral, and is therefore not phosphorous acid. The colloidal condition of this hydrate also points to a higher molecular weight. Phosphorous acid is formed only when this hydrate is decom- posed by heating (Beinitzer). PHOSPHOROUS ACID. PHo3. Molecular weight 82. Fuses at 70° C. (158° F.), Preparation.—l. Phosphorous acid is formed by the action of water upon the so-called phosphorous anhydride as above described, 2. It may also be obtained by the spontaneous oxidation of phos- phorus in moist air. In this process, however, a portion of the phos- phorous acid always undergoes further oxidation to phosphoric acid. Phosphorosophosphoric acid (q.v.) is also formed. 3. It is best obtained in a state of purity by the action of water on phosphorous chloride (see p. 345). It is not necessary to prepare the phosphorous chloride separately. Phosphorus is melted under water, and a stream of chlorine is passed through the phosphorus, the phosphorous chloride being thus decomposed by the water as fast as it is formed. I he reaction must be interrupted before all the phosphorus has disap- 352 INORGANIC CHEMISTRY. peared, otherwise the excess of the chlorine in presence of water will oxidize the phosphorous acid to phosphoric acid. The solution of hydro- chloric and phosphorous acids is evaporated, gradually raising the temperature to 180°, by which means the last traces of water are ex- pelled. Properties.—Phosphorous acid is a white, crystalline, very soluble mass, fusing at 70° C. (158° F.). Reactions.—1. When heated above 180° C. (356° F.), it yields phos- phoric acid and phosphoretted hydrogen : 4PHo3 = 3POHos + PH3. Phosphorous Phosphoric Phosphoretted acid. acid. hydrogen. 2. When treated with oxidizing agents, or when exposed to the air, it yields phosphoric acid : 2PHo3 + 02 = 2POHos. Phosphorous Phosphoric acid. acid. Owing to its affinity for oxygen it acts as a powerful reducing agent. Solutions of silver salts, when warmed with it, deposit metallic silver; mercuric chloride is reduced to mercurous chloride; and cupric sulphate yields a precipitate of cuprous hydride. Phosphites.—Phosphorous acid is a tribasic acid; but only the mono- basic and dibasic salts are stable. The normal sodium salt, PNao3, is obtained by dissolving phosphorous acid in an excess of sodic hydrate and adding absolute alcohol to the solution, wdien the salt is precipi- tated as an uncrystallizable syrup. It is decomposed by water (Zira- mermann). The phosphites are decomposed on heating, with evolution of phos- phoretted hydrogen and formation of metaphosphates and pyrophos- phates. The soluble salts have a reducing action. PHOSPHORIC ANHYDRIDE. Molecular weight = 142. PA- Preparation.—Phosphoric anhydride is obtained by burning phos- phorus in an excess of dry air or oxygen. A stream of air, dried by passing through a U-tube containing pumice moistened with sulphuric acid, is drawn by means of an aspirator, attached to the tube C, through the three-necked globe (Fig. 47). Thoroughly dried phosphorus is in- troduced through the tube B into the capsule A, and is then lighted by touching it with a hot wire, the tube being then closed with a cork. As soon as one piece of phosphorus is consumed, a fresh piece is intro- duced in the same way, and is now at once ignited by the hot capsule. META PHOSPHORIC ACID. 353 The phosphoric anhydride collects in the globe, whilst any particles which are carried off by the current of air are retained in the bottle. Fig. 47. Properties.—Phosphoric anhydride is a white, voluminous, amor- phous powder, which may be sublimed at a high temperature. Reaction.—When brought in contact with water it hisses violently, evolving great heat and dissolving with formation of metaphosphoric acid: P205 + 0H2 = 2P02Ho. Phosphoric Water. Metaphosphoric anhydride. acid. W7hen exposed to the air it rapidly absorbs moisture and deliquesces. It is the most powerful desiccating agent known, and is employed in the laboratory for removing moisture from gases and liquids. Many substances containing oxygen and hydrogen are decomposed by it, as it abstracts these elements in the proportions necessary to form water. METAPHOSPHORIC ACID. P02Ho. Molecular weight 80. Preparation.—l. Metaphosphoric acid is formed by dissolving phos- phoric anhydride in cold water (see above). 2. It may be obtained by heating tribasic phosphoric acid to redness : POH63 = P02Ho + oh2. Orthophosphoric Metaphosphoric Water, acid. acid. 354 INORGANIC CHEMISTRY. Properties.—Metaphosphoric acid forms a transparent vitreous mass which is readily soluble in water. It is fusible, and at a high tempera- ture may be volatilized. Its solutions coagulate albumen. Reaction.—In aqueous solution, metaphosphoric acid is gradually converted into tribasic phosphoric acid: P02Ho + OH2 = POHos. Metaphosphoric Water. Orthophosphoric acid. acid. This change takes place rapidly on boiling. Metaphosphates.—These salts may be obtained : 1. By igniting the dihydric phosphate of a fixed base: POHo2Nao = P02Nao + 0H2. Dihydric sodic Sodic Water, phosphate. metaphosphate. 2. By igniting a monohydric phosphate which contains one atom of a volatile base: POHoNao(NvH4O) = P02Nao + NH3 + 0H2. Hydric sodic ammonic Sodic Ammonia. Water, phosphate. metaphosphate. ' 3. By igniting a dihydric pyrophosphate : P203Ho2Nao2 = 2P02Nao + OH2. Dihydric disodic Sodic Water, pyrophosphate. metaphosphate. Properties of the Metaphosphates.—The metaphosphates are remark- able as existing in several distinct modifications, referable to different polymeric varieties of metaphosphoric acid. Most of these acids form double salts, and from the relative number of atoms of the two bases contained in such a salt, the minimum molecular weight of the acid may be determined. Thus, hexametaphosphoric acid, P6012Ho6, forms a double salt of the formula ,, P^NacCao"^80 " The soluble metaphosphates are converted into dihydric tribasic phos- phates by continued boiling with water ; the insoluble metaphosphates are converted in a similar manner by boiling with dilute nitric acid. The soluble metaphosphates yield with argentic nitrate a gelatinous white precipitate of argentic metaphosphate. PYROPHOSPHORIC ACID. 355 PYROPHOSPHORIC ACID. P203Ho4. Molecular weight = 178. Preparation.—l. Pyrophosphoric acid is prepared by heating tribasic phosphoric acid for some time to 213° C. : 2POHo3 = P203Ho4 + 0H2. Phosphoric Pyrophosphoric Water, acid. acid. 2. An aqueous solution of this acid is obtained by suspending plumbic pyrophosphate (prepared by precipitating sodic pyrophosphate with a soluble lead salt) in water, and decomposing it with sulphuretted hy- drogen : P203Pbo" + 2SH2 =• 2PbS// + P2O,Ho4. Plumbic Sulphuretted Plumbic Pyrophosphoric pyrophosphate. hydrogen. sulphide. acid. Properties.—Pyrophosphoric acid forms a colorless opaque crystal- line mass. It is readily soluble in water. The solution does not coagu- late albumen. Peactions.—l. In solution, pyrophosphoric acid is converted slowly at ordinary temperatures, rapidly on boiling, into tribasic phosphoric acid: P203Ho4 + 0H2 = 2POHo3. Pyrophosphoric Water. Orthophosphoric acid. acid. 2. On heating to redness it yields metaphosphoric acid P203Ho4 = 2P02Ho -(- 0H2. Pyrophosphoric Metaphosphoric Water, acid. acid. Pyrophosphates.—These salts are prepared by heating tribasic phos- phates in which two atoms of the hydrogen of the acid are replaced by a fixed base: 2POHoNao2 = P2OsNao4 -f- 0H2. Hydric disodic Sodic Water, phosphate. pyrophosphate. 2POMgo"(NH40) = P203Mgo"2 + 2NH3 -f 0H2. Magnesic ammonic Magnesic Ammonia. Water, phosphate. pyrophosphate. Pyrophosphoric acid is a tetrabasic acid and forms four classes of salts. Only the alkaline pyrophosphates are soluble in water; but the other pyrophosphates are soluble in acids, and generally also 356 INORGANIC CHEMISTRY. in an excess of an alkaline pyrophosphate, forming, in the latter case, solu- ble double salts. With argentic nitrate the alkaline pyrophosphates yield a white granular precipitate of argentic pyrophosphate; with solu- ble salts of copper, a double salt, of the formula p2Q3^aoCuo//^uo,/, is obtained. The solutions of the pyrophosphates are perfectly stable, even when boiled. By boiling with dilute acids, however, the pyro- phosphates are converted into tribasic phosphates. PHOSPHORIC ACID, Tribasic Phosphoric Acid, Orthophosphoric Acid. POHo3. Molecular weight 98. Fuses at 38.6° C. (101.5° F.). Preparation.—l. This acid is formed when phosphoric anhydride, raetaphosphoric acid, or pyrophosphoric acid is boiled with water for some time: P2Og + 30H2 = 2POHo3. Phosphoric Water. Orthophosphoric anhydride. acid. 2, It is best prepared in a state of purity by heating amorphous phosphorus with concentrated nitric acid. The oxidation is complete when red fumes cease to be evolved on the addition of fresh nitric acid. The excess of nitric acid is then driven off by evaporation. 3. It is formed by the action of water upon phosphoric chloride (p. 346) and phosphoric oxytrichloride {q.v.). 4. It is prepared on a large scale by treating 3 parts of bone-ash or phosphorite with 2 parts of sulphuric acid and 10 parts of water, heat- ing the mixture for some days: P202Cao//3 + 3S02Ho2 + 60H2 = 2POHo3 + 3SHo4Cao". Tricalcic Sulphuric Water. Phosphoric Gypsum phosphate. acid. acid. (Tetrahydric calcic sulphate). The solution is filtered from the insoluble calcic sulphate. The phosphoric acid prepared by any of the above methods, must be heated to 150° C. (302° F.) to expel the last traces of water. Properties.—Phosphoric acid forms transparent prisms, fusing at 38.6° C. (101.5° F.). When exposed to the air, it deliquesces to a syrupy liquid. Its solution does not coagulate albumen. Phosphates.—Phosphoric acid is a tribasic acid, forming three classes of salts, of which the following are examples : Trisodic phosphate, . . . . . PONao3,20H2. Hydric disodic phosphate, . . POHoNao2,120H2. Dihydric sodic phosphate, . . . POHo2Nao,OH2. PHOSPHORIC ACID. 357 The normal salts, with the exception of those of the alkalies, are in- soluble in water. Trilithic phosphate (POLio 3) is only sparingly sol- uble. The solutions of the normal alkaline phosphates have an alkaline reaction. In solution they are decomposed by carbonic anhydride with formation of monohydric phosphates : PONaog + C02 + OH2 = POHoNao, + COHoNao. Trisodic Carbonic Water. Hydric disodic Hydric sodic phosphate. anhydride. phosphate, carbonate. Dilute acids produce this change in the insoluble normal phosphates, dissolving them with formation of monohydric phosphates. The monohydric phosphates of the alkalies are soluble in water, and have a feebly alkaline reaction. The dihydric phosphates have an acid reaction. These compounds are sometimes referred to as superphosphates. The heavy metals form, as a rule, only normal phosphates, the other phosphates existing only in solution in presence of an excess of acid. If argentic nitrate be added to a solution of any of the alkaline phos- phates, a yellow precipitate of triargentic phosphate is formed : POJsTao3 + 3NQ2Ago = POAgo3 + 3N02Nao. Trisodic Argentic Triargentic Sodic phosphate. nitrate. phosphate. nitrate. POHoNao, + 3NQ2Ago = POAgo3 + 2NOaNao + NQ2Ho. Hydric disodic Argentic Triargentic Sodic Nitric acid, phosphate. nitrate. phosphate. nitrate. POHo2Nao + 3NQ2Ago = POAgo3 -f N02Nao -f 2N02Ho. Dihydric sodic Argentic Triargentic Sodic Nitric acid, phosphate. nitrate. phosphate. nitrate. It is worthy of note that, in the second of these reactions, by the mix- ture of two solutions, one of which is neutral and the other slightly alkaline, an acid liquid is produced. The soluble phosphates also yield a white crystalline precipitate of ammonic magnesic phosphate, PO(lSi vH4O)Mgo//,60H2, when a clear solution of magnesic sulphate and ammonic chloride containing an excess of ammonia is added to their solutions ; this precipitate is insoluble in water containing free ammonia, and on ignition is converted into magnesic pyrophosphate, P203Mgo"2. With a solution of ammonic molybdate in nitric acid, they yield, especially on warming, a yellow precipitate of ammonic phosphomolybdate {q.v.). The following are some of the more important naturally occurring phosphates : Apatite (FrancoJite) . . . P,OsCao"1(°Ca").* Vivian ite, . . PAFeo" ,80H2. Wavellite, . . . . P.O('Al"'An,120H!. Pyromorpbite, . . psoJPbo''Yg]Pb"). * In this mineral, chlorine and fluorine displace each other isomorphously. 358 INORGANIC CHEMISTRY. Some of the acids of phosphorus have a tendency to exhibit a basicity lower than their hydricity. Thus, though phosphoric acid forms tri- basic salts, the last equivalent of base is so loosely attached, that in the case of the soluble tribasic phosphates, it is removed by carbonic anhy- dride. In the case of phosphorous acid, a weaker acid, the tribasic salts are decomposed even by water, whilst hypophosphorous acid, a still weaker acid, forms only salts with one equivalent of base, though its formula would show it to be dibasic. PHOSPHOROSOPHOSPHORIC ACID (Hypophosphoric Add). Molecular weight = 324. P 404H08. Preparation.—When phosphorus is allowed to oxidize spontaneously by exposure to air and in contact with water, an acid liquid is obtained, which contains phosphorous acid, phosphoric acid, and phosphorosophosphoric acid. As the latter acid, when in solu- tion, gradually undergoes decomposition, the liquid is to be removed at the end of about three days. On adding sodic acetate a crystalline precipitate of tetrahydric tetrasodicphos- phorosophosphate, P404FIo4Nao4,l 201I2, is formed, which by recrystallization may be obtained in tabular crystals. The free acid is prepared by precipitating the barium salt with sulphuric acid or the lead salt with sulphuretted hydrogen. Reactions.—Phosphorosophosphoric acid can be obtained only in solution. On evap- oration over sulphuric acid, or even on standing at ordinary temperatures, it under- goes decomposition into phosphorous and pyrophosphoric acids : P404Ho8 -(- OH, = P203Ho4 —|- 2PHo3. Phosphoroso- Water. Pyrophosphoric Phosphorous phosphoric acid. acid. acid. Phosphorosophosphates.—These salts crystallize well. Owing to the high basicity of the acid, they are generally complex. The phosphorophosphates of potassium will serve as examples: Potassic phosphorosophosphate, P404Ko8,160H2 Dihydric hexapotassic phosphorosophosphate, P404Ho2Ko6,60H2. Tetrahydric tetrapotassic phosphorosophosphate, P404Ho4Ko4,40H2, also 60H2. Pentahydric tripotassic phosphorosophosphate, P404Ho5Ko3,20H2. Hexahydric dipotassic phosphorosophosphate, P404HoGKo2. That phosphorosophosphoric acid has at least the molecular weight here ascribed to it is rendered probable by the existence of such a salt as pentahydric tripotassic phosphor- osophosphate, and by the above decomposition of the free acid into a mixture of phos- phorous and pyrophosphoric acids. PHOSPHORIC OXYTRICHLORIDE. 359 COMPOUNDS OF PHOSPHORUS WITH CHLORINE AND OXYGEN. PHOSPHORIC OXYTRICHLORIDE, Phosphorylic Chloride. Cl 1 POCI3. ci—P—Cl II o Molecular weight = 153.5. Molecular volume. I I I. 1 litre of phos- phoric oxytrichloride vapor weighs 76.75 criths. Sp. gr. 1.7. Fuses at —l.s° C. (29.3° F.). Roils at 110° C. (230° F.). Preparation.—l. Phosphoric oxytrichloride may be prepared by de- composing phosphoric chloride with a limited quantity of water: PC15 + 0H2 = POCI3 + 2HCI. Phosphoric Water. Phosphoric Hydrochloric chloride. oxytrichloride. acid. 2. It is formed when oxygen is passed through boiling phosphorous chloride: PC13 + O = POC]3. Phosphorous Phosphoric chloride. oxytrichloride. 3. It may be readily obtained by heating together in a sealed tube a mixture of phosphoric chloride and phosphoric anhydride : PA + 3PC15 = SPOCl3. Phosphoric Phosphoric Phosphoric anhydride. chloride. oxychloride. 4. It is formed by the action of phosphoric chloride on various or- ganic and inorganic compounds containing oxygen (p. 346), and is best prepared by heating dried oxalic acid or boric acid with phosphoric chloride : {cOHo + PCI» = POC1* + C°! + 00 + 2HCI' Oxalic Phosphoric Phosphoric Carbonic Carbonic Hydrochloric acid. chloride, oxytrichloride. anhydride. oxide. acid. 2BHo3 + 3PC15 = 3P0C13 + BA + 6HCI. Boric Phosphoric Phosphoric Boric Hydrochloric acid. chloride. oxytrichloride. anhydride. acid. Properties.—Phosphoric oxytrichloride is a colorless powerfully re- fracting liquid which fumes in contact with moist air. In a freezing 360 INORGANIC CHEMISTRY, mixture it solidifies at —lo° C. (14° F.) to a laminar crystalline mass fusing at —l.s° C. (29.3° F.). Reactions.—l. By contact with water it is slowly transformed into hydrochloric and phosphoric acids : POCl3 + 30H2 = POHO3 + 3HCI. Phosphoric Water. Phosphoric Hydrochloric oxytrichloride. acid. acid. 2. By distillation with the salts of acids, it yields the corresponding acid chlorides: 3S02Pbo" + 2POCl3 = poPbc,//3 + 3S02C12. Plumbic Phosphoric Plumbic Sulphuric sulphate. oxytrichloride. phosphate. oxydichloride. 3{cONao + POC1» = 3{COCI + PON^- Sodic Phosphoric Acetylic Sodic acetate. oxytrichloride. chloride. phosphate. Phosphoric oxytrichloride is itself the acid chloride of phosphoric acid. This relation, which is better expressed by the name Phosphorylic chloride, is displayed in the above decomposition of this substance with water. The corresponding bromine compound POBr3 {molecular volume | | 1) is obtained in a similar manner by the action of a limited quantity of water on phosphoric bromide. It forms a crystalline mass fusing at 45-46° C. (113-115° F.), and boiling at 195° C. (383° F.). PYROPHOSPHORYLIC CHLORIDE. 0 o II II P203C14. Cl—P—o—P—Cl 1 I Cl Cl Molecular weight = 252. Sp. gr. 1.58 at 7° C. Boils with partial decomposition at 210- 215° C. Preparation.—This compound is prepared by passing gaseous nitric peroxide into phosphorous chloride, and distilling the liquid thus obtained. The portion which passes over between 200° and 230° C. is pyrophosphorylic chloride. This product must be purified by rectification. The reaction is a very complicated one, and cannot be expressed by a single equation. The by-products are phosphoric oxytrichloride, phosphoric anhydride, nitrous oxychloride, and nitrogen. Properties.—Pyrophosphorylic chloride is a colorless fuming liquid. Reactions.—1. Water decomposes it instantaneously with formation of orthophos- phoric (not pyrophosphoric) and hydrochloric acids: PACb + 50H2 = 2POHo3 + 4HCI. Pyrophosphorylic Water. Orthophosphoric Hydrochloric chloride. acid. acid. 2. When treated with phosphoric chloride, phosphoric oxytrichloride is formed : PAG], + PC15 - 3POCI3. Pyrophosphorylic Phosphoric Phosphoric chloride. chloride. oxytrichloride. TETRAPHOSPHORUS TRISULPHIDE—PHOSPHOROUS SULPHIDE. 361 COMPOUNDS OF PHOSPHORUS WITH SULPHUR. s s s II / \ II Tetraphosphorus trisulphide, . VI(P4)VIS"3. P—P P—P. S S II II Phosphorous sulphide, . . . P—B—P. 8 8 II II Phosphoric sulphide, . . . P2S"5. P—S—P. II II S 8 8 8 II II Diphosphoric tetrasulphide, . 'P1T2S"4. P—P. II il 8 8 8 8 These compounds are all formed by the direct union of their elements. Amorphous phosphorus and sulphur are heated together in the propor- tions required by the formulae. With ordinary phosphorus, the combi- nation is apt to take place with explosive violence. TETRAPHOSPHORUS TRISULPHIDE. Molecular weight = 220. Molecular volume \ | |. 1 litre of the vapor weighs 110 criths. Fuses at 166° C. (330.8° F.). Boils between 300° and 400° C. vI(PJviS//3. Preparation.—A mixture of amorphous phosphorus and sulphur in the proportions expressed by the formula P4S3 is heated for eight hours to a temperature of 260° C. (500° P.). The substance is thus obtained as a yellow translucent mass, which is puri- fied by crystallization from carbonic disulphide. Properties.—It forms yellowish prisms with a pyramidal termination. Reaction.—Boiling with water slowly decomposes it, with formation of phosphorous acid, phosphoretted hydrogen, and sulphuretted hydrogen: Ti(P4)viS//s + 90H2 = 3PHo3 + PH3 + 3SH2. Tetraphosphorus Water. Phosphorous Phosphoretted Sulphuretted sulphide. acid. hydrogen. hydrogen. PHOSPHOROUS SULPHIDE. Molecular weight 158. Preparation.—As above. Properties.—Phosphorous sulphide forms a greyish-yellow crystalline mass melting at about 290° C. (554° F.). It has not been obtained in definite crystals, and has not been distilled. Reaction.—Water decomposes it, forming phosphorous acid and sulphuretted hydrogen: P2S". P2S" + 60ET2 = 2PHo3 < + 3SH2. Phosphorous Water. Phosphorous Sulphuretted sulphide. aeid. hydrogen. 362 INORGANIC CHEMISTRY. PHOSPHORIC SULPHIDE. P S" 2U 5" Molecular weight = 222. Molecular volume I I I. 1 litre of the vapor weighs 111 criths. Fuses at 274—276° C. (525-529° F.). Boils at 530° C. (986° F.). Preparation.—As above. The process may also be modified by dis- solving ordinary phosphorus and sulphur in the molecular proportions, P2S"5, in carbonic disulphide, and heating the solution in sealed tubes for 8-10 hours to 210° C. (410° F.). On cooling, the phosphoric sulphide is deposited on the walls of the tube in well-formed crystals. Properties.—lt forms pale-yellow crystals generally grouped in tufts. Reactions.—l. By direct combination with alkaline sulphides it forms the sulphophosphates: P2S"5 + 3SK2 = 2PS"Ks3. Phosphoric Potassic Potassic sulphide. sulphide. sulphophosphate. 2. With water phosphoric sulphide yields phosphoric acid and sul- phuretted hydrogen: P2S"5 + 80H2 = 2POHo3 + 5SH2. Phosphoric Water. Phosphoric Sulphuretted sulphide. acid. hydrogen. Phosphoric sulphide is employed in the laboratory for the purpose of replacing oxygen by sulphur in organic compounds. DIPHOSPHORIC TETRASULPHIDB. Molecular weight = 190.* Fuses at 296-298° C. 'P‘*,S"4. Preparation.—Phosphorus and sulphur in the proportions corresponding with the formula P2S4 are dissolved in carbonic disulphide and heated in sealed tubes. Properties.—It is thus obtained in the form of yellow transparent acicular crystals. It boils without decomposition. COMPOUND OF PHOSPHORUS WITH SULPHUR AND CHLORINE. PHOSPHORIC SULPHOTRICHLORIDE. Cl I PS//CIs. Cl—P—Cl. II s Molecular weight 169.5. Molecular volume | | |. 1 litre of the vapor weighs 84.75 criths. Sp. gr. of liquid 1.636 at 20° C. Boils at 126° C. (259° F.). Preparation.—1. Phosphoric sulphotrichloride is best prepared by heating together phosphoric sulphide and phosphoric chloride for a few minutes to 150° C. (302° F.): * The vapor-density of this compound has been determined with a result which would point to the formula P3S6. This anomalous result is possiby due to the employ- ment of too low a temperature in the determination. PHOSPHORUS COMPOUNDS CONTAINING NITROGEN. 363 P2S"B •+ 3PC15 = SPS"Cls. Phosphoric Phosphoric Phosphoric sulphide. chloride. sulphotnchloride. 2. It is also formed by the action of sulphuretted hydrogen upon phosphoric chloride: PCJ5 + SH2 = PS//Cls + 2HCI. Phosphoric Sulphuretted Phosphoric Hydrochloric chloride. hydrogen, sulphotnchloride. acid. Properties.—lt is a colorless fuming liquid. Reactions.—l. Water slowly decomposes it, yielding hydrochloric acid, phosphoric acid, and sulphuretted hydrogen : PS//C13 + 40H2 = POHo3 + 3HCI + SH2. Phosphoric Water. Phosphoric Hydrochloric Sulphuretted sulphotnehloride. acid. acid. hydrogen. 2. With alkalies it yields the salts of sulphophosphoric acid (PS//Ho3): PS"C13 + 60KII = PS//Ko3 + 3KCI + SOH2. Phosphoric Potassic Potassic Potassie Water, sulphotnehloride. hydrate. sulphophosphate. chloride. The corresponding bromine compound PS//Br3 is also known. PHOSPHORUS COMPOUNDS CONTAINING NITROGEN. These substances possess considerable theoretical interest as examples of a class of compounds largely represented in organic, but of rarer occurrence in inorganic, chemistry. Phospham, PN(NH)//, is prepared by passing gaseous ammonia over phosphoric chloride as long as the gas is absorbed, and then igniting the product in a current of carbonic anhydride or some other indifferent gas : PC15 + 7NH3 = PN(NH)// + 5NH4CI. Phosphoric Ammonia. Phospham. Ammonic chloride. chloride. Phospham is a white powder, insoluble in water. Phosphamimide, PO(NH)//(NH2), remains behind as a white powder when the product of the action of gaseous ammonia on phosphoric pentachloride is extracted with water: PCI5 + 7NH3 + OH2 = PO(NH)"(NHa) + 5NH4CI. Phosphoric Ammonia. Water. Phosphamimide. Ammonic chloride. chloride. Phosphoric oxytriamide, PO(NH2)3, is obtained as a white amorphous powder by the action of gaseous ammonia on phosphoric oxytrichloride : POCl3 + 6NH3 = PO(NH2)3 + 3NH4CI. Phosphoric Ammonia. Phosphoric Ammonic oxytrichioride. oxytriamide. chloride. The product is well washed with water to remove the ammonic chloride. When this, or the foregoing compound, is ignited in an atmosphere free from oxygen, am- htenia is given off, and phosphoric oxynitride, PON, remains as a white powder. Pyrophosphotriamic acid, P203(NH2)3Ho, is prepared by saturating phosphoric oxy- trichloride with gaseous ammonia without cooling, heating the product to 220° C., and nnally boiling it for a short time with water: 2P0C13 + 9NH3 + 20H2 = P203(NH2)3Ho + 6NfT4CI. Phosphoric Ammonia. Water. Pyrophosphotriamic Ammonic oxytrichioride. acid. chloride. 364 INORGANIC CHEMISTRY. boiling with water into soluble pyrophosphodiamic acid, P203(NH2)2Ho2, and pyrophos- phamic acid, P203{NH2)Ho3, this last compound being finally transformed into a mix- ture of ammonic phosphate and phosphoric acid. It forms an amorphous insoluble powder, which is successively converted by continuous VANADIUM, V 4? Atomic weight = 51.3. Probable molecular weight = 205.2. Sp.gr. 5.5. Atomicity and v. Evidence of atomicity: Vanadoiis chloride, V'^Clg. Vanadic oxytrichloride, VtOC13. History.—This rare element was discovered in 1801, by Del Rio, who obtained it from a Mexican lead-ore. He failed, however, to rec- ognize its true nature, and ultimately regarded it as impure chromium. In 1830 it was rediscovered independently by Sefstrom. Metallic vanadium was first isolated by Roscoe. Occurrence.—Vanadium occurs sparingly in various lead and iron ores. The cupric and bismuthous vanadates constitute the rare minerals volborthite and pucherite. A relatively rich source of vanadium has lately been found in the Bessemer slag of the Creusot iron works, which contains as much as 1.5 per cent, of this element. Preparation.—Metallic vanadium is obtained by heating vanadous chloride to bright redness in a current of dry hydrogen: 2VC13 + 3H2 = V 2 + 6HCI. Variations Hydrochloric chloride. acid. Properties.—As above prepared it forms a silvery, crystalline mass, of sp. gr. 5.5. It does not oxidize, either in dry or in moist air, even at 100° C. When strongly heated in air or oxygen it burns, forming vanadic anhydride, V 2Os. It does not fuse at a red heat. Hydrochloric acid is without action upon it; concentrated sulphuric acid dissolves it on heating; and nitric acid, even when dilute, attacks it energetically, dissolving it to form a blue solution. Fused with caustic alkalies it yields a vanadate of the base with evolution of hydrogen. COMPOUNDS OF VANADIUM WITH CHLORINE. Hypovanadous chloride, . . rvci2 • * ■ t vci2. Yanadous chloride, .... . . . VC13. Hypovanadic chloride, . . , . . . 'V'vci.orj^j.. Hypovanadous chloride,/V"2CI4, is obtained in apple-green micaceous plates by pass- ing the vapor of the trichloride mixed with hydrogen through a red-hot tube: 2VCI3 + H2 = + 2HCI. Vanadous Hypovanadous Hydrochloric chloride. chloride. acid. YANADICM. 365 It is hygroscopic, and dissolves in water, yielding a violet solution. Vanadous chloride, VC13, is prepared from hypovanadic chloride, which is decom- posed slowly at ordinary temperatures, rapidly at its boiling-point, into vanadous chlo- ride and free chlorine. It forms peach-blossoin-colored tabular crystals, is non-volatile, and deliquesces when exposed to the air. Hypovanadic chloride, /V’VCI+ (molecular volume j | |), is formed by the action of an excess of chlorine on metallic vanadium. It may also be obtained by repeatedly pass- ing the vapor of the oxytrichloride, mixed with chlorine, over charcoal: 2VOC13 + C 2 + Cl2 = 2/Vi7Cl4 + 2CO. Yanadic Hypovanadic Carbonic oxytrichloride. chloride. oxide. It is a dark-brown liquid, boiling at 154° C., and having a sp. gr. of 1.8584 at ,0° C. Water decomposes and dissolves it, yielding a blue liquid. The molecular formula, VC14, as deduced from the vapor-density of this compound, is anomalous. In such a compound, vanadium would be tetradic, in violation of the law regulating the variation of atomicity; otherwise, the presence of a single free bond must be assumed (see note, P- 179), COMPOUNDS OF VANADIUM WITH OXYGEN AND HYDROXYL. Hvpovanadons oxide, /V// O Vanadous oxide, .... v./v Hypovanadic oxide, .... 'v-A- Vanadic anhydride, vft. Metavanadic acid, . .... VO, Ho. Tribasic vanadic acid, .... VOHo3, Pyrovanadic acid, . .... V,03Ho4. Hypovanadous oxide, /V//202, is formed when the vapor of the oxytrichloride, mixed with hydrogen, is passed through a red-hot tube: 2VOC13 + 3H2 = 'V"A + 6HCI. Vanadic Hypovanadous Hydrochloric oxytrichloride. oxide. acid. It is a gray powder, with a metallic lustre. Acids dissolve it, yielding a lavender- colored solution, which instantly becomes brown on exposure to the air. Hypovanadous oxide w'as mistaken by Berzelius for metallic vanadium. Vanadous oxide, V 203, remains behind as a black lustrous powder when vanadic an- hydride is heated to redness in a current of hydrogen. Even at ordinary temperatures it slowly absorbs oxygen, forming hypovanadic oxide, /ViV204, and, when gently warmed in air, glows and is converted into vanadic anhydride. It is insoluble in acids. Hypovanadic oxide, /Vtv204, is formed as above by the spontaneous oxidation of vana- dous oxide. • It may also be obtained by fusing together equal molecular proportions of vanadous oxide and vanadic anhydride: VA + V 205 = 2/Vtv204. Yanadous Vanadic Hypovanadic oxide. anhydride. oxide. It is a blue powder, consisting of minute shining crystals. When exposed to moist an it is slowly converted into an olive-green hydrate. Acids dissolve it with difficulty, yielding a blue solution. Venadic anhydride, V 205.—Minerals containing vanadium are fused with nitre, and the mass is extracted with water. The solution, which contains an alkaline vanadate along with various impurities, is then almost neutralized with nitric acid and precipi- tated with baric chloride. The precipitate, consisting of baric vanadate and other barium salts, is decomposed by boiling with dilute sulphuric acid, and the solution, filtered from the baric sulphate, is neutralized with ammonia and evaporated to a small bulk, after which pieces of ammonic chloride are placed in the solution. This causes the ammonic metavanadate, which is very insoluble in a concentrated solution of ara- ttionic chloride, to be deposited in small crystals. These are washed with a solution 366 INORGANIC CHEMISTRY. of amnionic chloride, and decomposed by ignition in an open crucible, when pure va- nadic anhydride remains behind. Vanadic anhydride is a reddish-brown mass which melts at a red heat, and solidifies in a crystalline form on cooling. It is very slightly soluble in water, to which it im- parts a yellowish tinge. Both acids and alkalies dissolve it readily. The acid solu- tions yield with reducing agents first a blue, and afterwards a green coloration. Vanadates.—The various forms of vanadic acid are known only in their salts. The orthovanadates (or tribasic vanadates), the metavanadates, and the pyrovanadates are isomorphous with the corresponding compounds of phosphorus. In addition to these, a fourth series is known, the tetravanadates, of which diammonic tetravanadate, V40„(NII40)2,401I2, is an example: Ammonic metavanadate, . . . . , . . . VOs(NH40). Argentic orthovanadate, . . . . . . . VOAgOj. Argentic pyrovanadate, . . . . Vanadinite, . . . . V803Pbo//4^Q1Pb// ARSENIC, As4. Atomic weight = 75. Molecular weight = 300. Molecular volume I I I. 1 litre of arsenic vapor weighs 150 criths. Sp. gr. 5.6 to 5.9. Volatile at 180° C. (356° F.). Atomicity and T. Evidence of atomicity: Arsen iuretted hydrogen, . . . . . . As^'Hg. Arsen ions chloride, . . . As///C13. Tetrethylarsenic chloride, . . . AsvEt4Cl. History.—Arsenic was known to the alchemists, but Brand, and later Scheele, first investigated its chemical nature. Occurrence.—Arsenic is widely distributed in nature. It occurs both in the free state and in combination with various other metals in the form of ores. Of the latter the principal are: realgar, fAs"2S"2; orpi- ment, As2S"3; arsenical pyrites, /As//2(Fe2S/'2)//2; and arsenical iron, 'As"2FeiT. It is found in small quantities in other minerals, such as iron pyrites, for which -reason sulphuric acid which has been manufac- tured from pyrites is generally contaminated with arsenic. In minute traces it occurs in some mineral waters, and in the water and mud of many rivers. It is also contained in coal-smoke (derived in this case from the pyrites of the coal), and consequently in the air of towns. Preparation.—1. Arsenic is obtained by heating arsenical pyrites. The arsenic volatilizes and may be condensed, whilst ferrous sulphide remains behind: 'As"2(Fe2S"2)"2 = As2 + 4FeS''. Ferroarsenious Ferrous sulphide. sulphide. 2. It may also be prepared from arsenious anhydride, a substance produced in the roasting of many ores. The arsenious anhydride is reduced by heating with charcoal: As203 + 3C = 2As + SCO. Arsenioas Carbonic anhydride. oxide. AESENIURETTED HYDROGEN. 367 Properties.—Arsenic, like phosphorus, is known in more than one form. When arsenic is sublimed in a current of hydrogen, it is de- posited close to the heated portion of the tube in crystals, but further on, where the tube is colder, amorphous arsenic collects. The crystal- line variety forms acute rhombohedra, with a steel-gray color and a metallic lustre, possessing a sp. gr. of 5.727. In dry air it may be pre- served without change, but in presence of moisture it becomes coated with a blackish-gray crust due to oxidation. When heated under or- dinary pressure, it volatilizes without fusing; but by inclosing it in a sealed tube, so as to subject it to the pressure of its own vapor, it may be fused. The vapor is lemon-colored, and possesses an odor of garlic. The molecular weight of arsenic, as deduced from the vapor-density, is 300, showing that the molecule of this element is, like that of phos- phorus, tetratornic. At the highest temperature at which the vapor- density of arsenic has been determined (yellow heat), a partial disso- ciation is, however, found to have occurred, and the value for the vapor- density lies between those required for As2 and As4 (Victor Meyer). The amorphous variety forms a black mass with a vitreous lustre. Its sp, gr. is 4,71. When heated to 360° C. (680° F.) it is converted into the crystalline or metallic variety, great heat being liberated in the transformation. It is much more permanent in air than crystalline arsenic. Amorphous arsenic may also be obtained as a gray powder. This variety is deposited in the coldest parts of the tube during the sub- limation in hydrogen. Reactions.—l. When heated in air or oxygen arsenic burns, forming arsenious anhydride. In like manner, when arsenic is treated with oxidizing agents, arsenious anhydride and arsenic acid are produced. 2. When finely-divided arsenic is introduced into chlorine, it inflames spontaneously, yielding arsenious chloride. It also combines directly with most of the other elements. Use.—A small quantity of arsenic is added to the lead which is used m the manufacture of shot, as it is found that this addition enables the metal more readily to assume the spherical form, and at the same time renders it harder. COMPOUND OF ARSENIC WITH HYDROGEN AESENIURETTED HYDROGEN, Arsenious Hydride. AsH3. Molecular weight 78. Molecular volume I I I. 1 litre weighs 39 criths. Boils at —4o° C. (—4o° F.). Preparation.—l. This gas is obtained in the pure state by the action °f dilute sulphuric or hydrochloric acid on an alloy of arsenic and zinc: As2Zn"3 + 3S02Ho2 = 3S02Zno" + 2AsH3. Arsenious Sulphuric Zincic Arseniuretted zincide. acid. sulphate. hydrogen. 368 INORGANIC CHEMISTRY. 2. It is formed by the action of nascent hydrogen upon soluble arsenic compounds: thus by the introduction of arsenious acid into an apparatus evolving hydrogen from zinc and sulphuric acid: AsHo3 + 3H2 = AsH3 + 80H2. Arsenious Arseniuretted Water, acid. hydrogen. In this case the gas is mixed with an excess of hydrogen. Properties.—Arseniuretted hydrogen is a colorless gas of a very dis- agreeable odor. At —4o° C. it condenses to a colorless and trans- parent liquid which does not solidify at —loo° C. (—l4B° F.). Water dissolves it but slightly. It is devoid of basic properties. It is one of the most poisonous substances known. Gehlen, of Gst- tingen, lost his life by incautiously smelling a leaky joint of an appa- ratus in which he was preparing the gas, in order to detect the escape. Reactions.—l. When burnt with free access of air it forms water and arsenious anhydride: 2AsH3 + 302 As203 + 30H2. Arseniuretted Arsenious Water, hydrogen. anhydride. 2 When burnt with a limited supply of air, it yields water and free arsenic: 4AsII3 -\- 302 As4 -)~ 60H2. Arseniuretted Water, hydrogen. Thus if a piece of white porcelain be held in the flame of arseniuretted hydrogen burning in air, a black shining spot of metallic arsenic is deposited on the porcelain, 3. When exposed to a low red heat, it is decomposed into arsenic and hydrogen. This reaction, coupled with the formation of arseniu- retted hydrogen by the action of nascent hydrogen on soluble com- pounds of arsenic, is employed as a means of detecting minute traces of this element, (See Marsh’s Test, Reactions of Arsenic.) 4. W7hen passed through a solution of argentic nitrate, it yields a precipitate of metallic silver, whilst arsenious and nitric acids remain in solution: 6N02Ago -J- 30H2 -|- AsH3 = 6N02Ho -}- AsHo3 3Ag2. Argentic Water. Arseniuretted Nitric Arsenious nitrate. hydrogen. acid. acid. ARSENIOUS CHLORIDE. 369 COMPOUNDS OF ARSENIC WITH THE HALOGENS. ARSENIOUS CHLORIDE AsC]3. Atomic weight = 181.5. Molecular volume 1 I i. 1 litre of arsenious chloride vapor weighs 90.75 criths. Sp. gr. 2.205. Boils at 134° C. Preparation.—l. Arsenious chloride is obtained by the action of dry chlorine on arsenic. The product must be left in contact with arsenic, in order to free it from excess of chlorine, and then rectified. 2. It may also be prepared by distilling arsenic with corrosive sub- limate : As2 + 6HgCl2 = 3/Hg'2Cl2 + 2AsC13. Mercuric Mercurous Arsenious chloride. chloride. chloride. 3. It is most readily obtained by distilling a mixture of arsenious anhydride, sodic chloride, and concentrated sulphuric acid: As203 + 6NaCI + 6S02Ho2 = 2AsC13 Arsenious ' Sodic Sulphuric Arsenious anhydride. chloride. acid. chloride. 4 (iS02TIoNao 4- 301I2. Hydric sodic Water, sulphate. In this way hydrochloric acid, in the preparation of which arsenical sulphuric acid has been employed, always contains arsenic, 4. When a solution of arsenious anhydride in aqueous hydrochloric acid is boiled, arsenious chloride volatilizes along with the steam: As203 4- 6HCI 2AsC13 + 30H2. Arsenious Hydrochloric Arsenious Water, anhydride. acid. chloride. Properties.—Arsenious chloride is an oily liquid which does not solidify at —29° C. It fumes strongly in contact with moist air. It is extremely poisonous. Reactions.—l. A small quantity of water dissolves it, forming a clear solution, from which needle-shaped crystals of arsenious chlordihydrate are deposited on standing; AsC13 A 20H2 = AsCIHo2 + 2HCI. Arsenious Water. Arsenious Hydrochloric chloride- chlordihydrate. acid. 2. An excess of wTater decomposes it into arsenious anhydride and hydrochloric acid: 370 INORGANIC CHEMISTRY. 2AsC13 + 301T2 = 6HCI + As203. Arsenious Water. Hydrochloric Arsen ions chloride. acid. anhydride. 3. It absorbs gaseous ammonia, forming a crystalline compound of the formula AsG13,2NH3. Arsenious bromide, Asßr3 (molecular volume | | |), is prepared by adding finely powdered arsenic to a solution of bromine in carbonic disulphide. It crystallizes in colorless deliquescent prisms, fusing between 20° and 25° C. (68-77° F.). It boils at 220° C. (428° F.). Water decomposes it like the chloride. Arsenious iodide, AsI3 {molecular volume | | j), is prepared in a similar manner, and forms lustrous brick-red laminae. Arsenious fluoride, AsF3 {molecular volume | | I), is obtained by distilling a mixture of 1 part of powdered fluorspar and 1 part of arsenious anhydride with 5 parts of concentrated sulphuric acid : -A.5203 “f" 6IIF == 2AsFg -f" 30 H2. It is a colorless fuming liquid of sp. gr. 2.7, boiling at 63° C. (145° F.). Brought in contact with the skin it produces very dangerous wounds. Water decomposes it like arsenious chloride. Arsenic pentafluoride, AsF5, is known only in the form of the double com- pound, AsFS,KF, obtained by dissolving potassic arsenate in hydrofluoric acid. COMPOUNDS OF ARSENIC WITH OXYGEN AND HYDROXYL. Arsenious anhydride, .... • • • • (As203)2. Arsenic anhydride, .... As2Og. Arsenious acid, Arsenic acid, .... AsOHo3. ARSENIOUS ANHYDRIDE, Arsenic, White Arsenic, White Oxide of Arsenic. (A5203)2. Molecular weight = 396. Molecular volume I I I. 1 litre of arsenious anhydride vapor weighs 198 criths. Sp. gr. (octahedral) 3.69, (amor- phous) 3.74. Occurrence.—Arsenious anhydride is found in nature in two rare minerals : in the octahedral form as arsenic bloom and in rhombic crys- tals as claudetite. Preparation.—lt is formed when arsenic is burnt in air or oxygen. In this way, it is obtained as a by-product in the roasting of arsenical ores in various metallurgical operations. The arsenious anhydride sub- limes and is condensed in large flues. Properties.—Commercial amorphous arsenious anhydride forms, when first prepared, a colorless vitreous mass, which after a time becomes ARSENIOUS ACID. 371 white and opaque, owing to its gradual transformation into the crystal- line variety. This change is accompanied by a decrease in the sp. gr. from 3,74 to 3.69. Crystalline arsenious anhydride is soluble in 80 parts of cold water; the amorphous modification in 25 parts; hydro- chloric acid increases the solubility. The hot saturated solution in hydrochloric acid deposits octahedral crystals; if a solution of the amorphous variety be employed, the formation of each crystal is attended with a flash of light visible in the dark. When the octahedral anhy- dride is heated, it sublimes without fusing, and is again condensed in octahedral crystals; under pressure, however, it may be fused, and is thus converted into the amorphous modification. A second crystalline modification of arsenious anhydride, belonging to the rhombic system, is sometimes found in the arsenic flues. It is also deposited from a solution of an excess of the anhydride in boiling caustic potash, or from a solution of argentic arsenite in nitric acid. Arsenious anhydride has a faint, sweetish metallic taste, and, when taken internally, acts as an irritant poison. A dose of 0.06 gram has been known to prove fatal. The best antidote is freshly prepared ferric hydrate which must be administered in a large dose as soon as possible after the poison has been swallowed. The arsenious acid is oxidized by the ferric hydrate to arsenic acid, which combines with the excess of ferric hydrate to form a basic ferric arsenate, insol- uble in water and in the liquids of the stomach. By keeping, the fer- ric hydrate becomes crystalline and inactive: it is therefore prepared, when wanted, by adding calcined magnesia to a solution of ferric chlo- ride or sulphate: 'Fe'"2Cl6 + 3MgO + 30 H2 = 'Fe'"2Ho6 + 3MgCl2. Ferric Magnesia. Water. Ferric Magnesic chloride. hydrate. chloride. The magnesic chloride, which is simultaneously formed in thisreaction, serves by its aperient action to remove the various matters as speedily as possible from the stomach. In spite of the poisonous properties of arsenious anhydride, it is pos- sible by long use to train the system to support relatively large doses of this substance. In Styria, the practice of arsenic eating is stated to be not uncommon. An arsenic eater has been known to consume 0.3 gram of arsenious anhydride at once without perceptible ill-effect. The practice is asserted by the arsenic eaters to improve the complexion and to prevent shortness of breath. Uses.—Arsenious anhydride is employed in the preparation of arsen- ical pigments and in the manufacture of glass. ARSENIOUS ACID. AsHo3. Molecular weight = 126. When arsenious anhydride is dissolved in water the solution reddens litmus feebly, and contains arsenious acid. This acid cannot however 372 INORGANIC CHEMISTRY. be isolated, since on concentration the solution deposits crystals of the anhydride. Arsenites.—There are two classes of arsenites: the ortharsenites, de- rived from the acid AsHo3; and the raetarsenites, derived from the acid AsOITo. Only the arsenites of the alkali-metals are soluble in water. They yield with argentic nitrate a yellow precipitate of triar- gentic arsenite, AsAgo3. Among the arsenites important by their uses, may be mentioned dihydric potassic arsenite, AsHo2Ko, employed in medicine under the name of Fowler's solution ; and hydrie cupric arsen- ite, AsHoCuo", which forms the pigment Soheele’s green. Schweinfurt green, a double metarsenite and acetate of copper of the formula 3As202Cuo'/;Cu(C2H302)2, is prepared by boiling arsenious acid with cupric acetate. ARSENIC ANHYDRIDE. As2Os. Molecular weight = 230. Preparation.—This compound is obtained by heating arsenic acid nearly to redness: 2AsOHo3 = 30H2 + As205. Arsenic Water. Arsenic acid. anhydride. Properties.—It forms a white porous mass. Reactions.—1. It dissolves in water with formation of arsenic acid. 2. When heated to bright redness it fuses and is decomposed into arsenious anhydride and oxygen. 3. With gaseous hydrochloric acid it yields, even at ordinary tem- peratures, arsenious trichloride, chlorine, and water: As205 + 10HC1 = 2AsC13 + 2CI2 + 50H2. Arsenic Hydrochloric Arsenious Water, anhydride. acid. trichloride. ARSENIC ACID. AsOHo3. Molecular weight = 142. Preparation.—Arsenic acid is prepared by treating arsenious anhy- dride with nitric acid: As203 + 2N02Ho -J- 20H2 •= 2AsOHo3 + N203. AHSENIC ACID. 373 Properties.—When a solutien of arsenic acid is evaporated to a syrup, and then cooled below 15° C. (59° F.), crystals of the formula 2AsOHo3,0H2 are deposited. These crystals part with their water of crystallization at 100° C., and are converted into ortharsenic acid, AsOHo3. When this acid is heated to 180° C. (356° F.) it fuses and gives off water, and on cooling, hard shining prismatic crystals of pyrarsenic acid, As203Ho4, separate out. If the heating be carried to 206° C. (403° F.), the whole is converted into a white nacreous mass of metarsenic acid, As02Ho. These three acids correspond to the three varieties of phosphoric acid ; but the pyro- and metarsenic acids differ from the pyro- and meta-phosphoric acids in being capable of existing only in the solid state. In solution they are at once converted into ortharsenic acid, and the same is the case with their salts, which may be prepared in the same way as the corresponding salts of meta- and pyro- phosphoric acid (pp. 354 and 355). Reactions.—l. When arsenic acid is distilled with fuming hydro- chloric acid, arsenious trichloride, chlorine, and water distil over : AsOHo3 + SHCI = AsC13 + Cla + 40H2. Arsenic Arsenious Water, acid. trichloride. In the receiver the reverse reaction takes place, arsenic acid and hydro- chloric acid being regenerated. 2. Sulphurous acid reduces it to arsenious acid; AsOHo3 T SOHo2 = AsHo3 -f S02Ho2. Arsenic Sulphurous Arsenious Sulphuric acid. acid. acid. acid. Other reducing agents act in a similar manner. Arsenates.—The arsenates are isomorphous with the corresponding phosphates. Arsenic acid is a tribasic acid, and forms three series of salts; normal, monohydric, and dihydric. The alkaline arsenates are soluble in water; of the others only the dihydric salts are soluble, but all dissolve readily in acids. The following arsenates occur in nature : Haidingerite, .... . . . 2AsOHoCao//,OH9. Pharmacol ite,.... Cobalt bloom,.... Mimetesite (isomorphous with py-1 ,q romorphite, vanadinite, and ap- V As303Pbo///Jp,Pb//j. atite, pp. 357 and 366), . . J ' For the reactions of the arsenates, see General Reactions of Arsenic. 374 INORGANIC CHEMISTRY. COMPOUNDS OF ARSENIC WITH SULPHUR AND HYDE OS ULPHYL. Realgar, f AsS" __ „ s„ \ AsS" ~ 2’ Sulpharsenious anhydride {Ar- As.S", senious sulphide), Sulpharsenic anhydride (Arsenic As ,8" sulphide), Sulpharsenious acid, .... AsHp3. Sulpharsenic acid, AsS^Ksg. DIARSENIOUS DISULPHIDE, Realgar. 7As;/ S" or/^8^ Ab 2o 2> Ul \ sg"t Molecular weight = 214. Sp. gr. 3.5. Occurrence.—This substance occurs in nature as the mineral realgar. Preparation.—l. It may be obtained artificially by heating together sulphur and arsenic in the proportions expressed by the formula. 2. A second method consists in heating in an iron retort a mixture of arsenical pyrites and iron pyrites: /As"2(Fe2S//2)//2 + 2FeS"2 = 'As"2S"2 + 6FeS". Arsenical Ferric Diarsenious Ferrous pyrites. disulphide. disulphide. sulphide. The diarsenious disulphide distils over, whilst ferrous sulphide remains in the retort. Most of the realgar of commerce is prepared by this method. Properties.—Native realgar occurs in ruby-colored monoclinic prisms and also massive. The artificial product forms a dark-red crystalline mass. It fuses readily and may be distilled without decomposition. Reaction.—1. When heated in contact with air it burns with forma- tion of arsenious and sulphurous anhydrides: 'As"28"2 + 70 = As203 + 2S02. Diarsenious Arsenious Sulphurous disulphide. anhydride. anydride. SULPHARSENIOUS ANHYDRIDE, Arseniom Sulphide, Orpiment. As2S"3. Molecular weight = 246. Sp. gr. 3.5. Occurrence.—Sulpharsenious anhydride occurs native as the mineral orpiment. SULPHA ESENIOUS ANHYDRIDE. 375 Preparation.—lt may be obtained by precipitating a solution of ar- senious anhydride in hydrochloric acid with sulphuretted hydrogen; 2AsC13 + 3SH2 = 6HCI + As2S"3. Arsenious Sulphuretted Hydrochloric Arsenious chloride. hydrogen. acid. sulphidei Properties.—Native orpiment forms lemon-colored rhombic prisms. The substance obtained by precipitation is a yellow powder, which fuses to a reddish liquid, and may be distilled without decomposition. Reactions.—l. Arsenious sulphide dissolves in caustic alkalies, pro- ducing an arsenite and a sulpharsenite: As2S"3 + 40KH = AsHoKo2 + AsHsKs2 + OH2. Arsenious Potassic Hydric dipotassic Sulphydric Water, sulphide. hydrate. arsenite. disulphopotassic sulpharsenite. By the addition of an acid the arsenious sulphide is reprecipitated: AsHoKo2 + AsHsKs2 + 4HCI = 4KCI Hydric dipotassic Sulphydric Hydrochloric Potassic arsenite. disulphopotassic acid. chloride, sulpharsenite. H~ As2S"3 -f" 30H2. Arsenious Water, sulphide. 2. It dissolves in alkaline sulphides, forming sulpharsenites : As2S"3 + 3SK2 = 2AsKs3. Arsenious Potassic Potassic sulphide. sulphide. sulpharsenite. Sulpharsenites.—These salts correspond to the arsenites. Only the alkaline salts are soluble. On the addition of an acid to their solutions arsenious sulphide is precipitated, and sulphuretted hydrogen is evolved; 2AsKs3 + 6HCI = 6KCI + As2S"3 + 3SH2. Potassic Hydrochloric Potassic Arsenious Sulphuretted sulpharsenite. acid. chloride. sulphide. hydrogen. Colloidal Arsenious Sulphide.—On saturating a pure aqueous solution of arsenious anhydride with sulphuretted hydrogen, the liquor assumes a yellow color with a reddish fluorescence; but no precipitate is formed. In this condition the solution contains a colloidal modification of arsenious sulphide, which may be separated from unaltered arsenious anhydride by dialysis. By spontaneous evaporation this soluble sulphide is obtained in transparent amorphous masses of a yellow or reddish-yellow color with a conchoidal fracture. Acids and various metallic salts precipitate ordinary insoluble arsenious sulphide from the solution. 376 INORGANIC CHEMISTRY. SULPHARSENIO ANHYDRIDE, Arsenic Sulphide. Molecular weight = 310. As2S". Preparation.—1. Sulpharsenic anhydride maybe prepared by fusing together arsenious sulphide and sulphur in the required atomic propor- tions. 2. It is obtained as a yellow precipitate by adding hydrochloric acid to a solution of a sulpharsenate: 2AsS//Ks3 + 6HCI = 6KCI + As2S" + 3SH2. Potassic Hydrochloric Potassic Arsenic Sulphuretted sulpharsenate. acid. chloride. sulphide. hydrogen. It cannot, as was formerly supposed, be prepared by passing sulphuret- ted hydrogen through a solution of arsenic acid. The yellow precipi- tate formed under these circumstances is a mixture of arsenious sulphide and sulphur in the proportion -- S2. Sulphur senates.—These salts may be prepared by passing sulphuretted hydrogen through solutions of arsenates: AsOHoNao, + 4SH2 = AsS/rHsNas2 + 4014... Hydric disodic Sulphuretted Sulphydric Water, arsenate. hydrogen. disulphosodic sulpharsenate. General Properties and Reactions of the Compounds of Arsenic.—Owing to the frequency of cases of poisoning, both accidental and intentional, with arsenic, the detection of this substance, even when present in the minutest traces, becomes a matter of great importance. For a detailed account of the methods to be employed and of the pre- cautions to be taken, in a toxicological investigation of this kind, special works on analytical chemistry must be consulted. a. Arsenites.—From the hydrochloric acid solution of an arsenite or of arsenious anhydride, sulphuretted hydrogen precipitates yellow ar- senious sulphide. The precipitate is formed slowly In the cold, more rapidly on warming; it is soluble in ammonic sulphide, caustic alkalies, amnionic carbonate, and hydric potassic sulphite; almost insoluble in hydrochloric acid. Soluble arsenites yield in neutral solution with ar- gentic nitrate a yellow precipitate of argentic arsenite, soluble both in nitric acid and in ammonia. With a solution of arsenious anhydride the yellow precipitate only makes its appearance on the careful addition of ammonia, so as to neutralize the free nitric acid. h. Arsenates.—Only the alkaline arsenates are soluble in water. From neutral solutions argentic nitrate precipitates reddish-brown triar- gentic arsenate, soluble in ammonia. A mixture of magnesic sulphate, ammonic chloride, and ammonia gives a white crystalline precipitate of ammonic magnesic arsenate (AsOMgo"Amo,6oH2), isomorphous with the corresponding phosphorous compound (p. 357). Sulphuretted hydro- NIOBIUM AND TANTALUM, 377 gen in acid solutions first reduces the arsenic acid to arsenious acid with separation of sulphur; after which the arsenious acid is precipitated as arsenious sulphide. In the cold, the reduction of arsenic acid by sul- phuretted hydrogen takes place with extreme slowness; the action is greatly aided by keeping the liquid at a temperature of from 50° to 70° C. (122-158° F.) while passing in the sulphuretted hydrogen. Marsh’s Test.—lf any of the oxygen or halogen compounds of ar- senic be introduced into an apparatus in which hydrogen is being gen- erated from zinc and dilute sulphuric acid, the arsenic is evolved as arseniuretted hydrogen together with an excess of hydrogen. If the escaping gas be ignited and a cold surface of white porcelain be held in the flame, a black lustrous film of metallic arsenic is deposited upon the porcelain. In like manner, if the gas be passed through a strongly heated glass tube, metallic arsenic condenses as a lustrous mirror just be- yond the heated portion. These thin films of arsenic are at once dissolved by a solution of sodic hypochlorite. (Distinction from antimony.) The sulphur compounds of arsenic, and metallic arsenic itself, do not yield arseniuretted hydrogen under the above conditions. The presence of nitric acid and other oxidizing agents prevents the formation of arseni- uretted hydrogen. In applying Marsh’s test, and all similar tests, it is necessary to ascertain by a blank experiment that the reagents employed are free from arsenic. Reinsch’s Test.—lf a solution of an arsenic compound in hydrochloric acid be boiled with a piece of pure bright copper, the surface of the metal becomes covered with a dark-gray coating of arsenide of copper. If this coating be separated, dried, and then heated in a small glass tube, a portion of the arsenic is oxidized to arsenious anhydride, which forms a sublimate of minute transparent octahedra in the tube. To this sublimate the above confirmatory tests may be applied. This test ought never to be trusted when the mixture contains a chlorate or a nitrate, as a portion of the copper will then be dissolved, and the traces of arsenic which are generally present in the purest copper will be precipitated on the remaining copper. All compounds of arsenic, when heated in a narrow bulb-tube with a mixture of sodic carbonate and potassic cyanide, are reduced to me- tallic arsenic, which sublimes and collects as a mirror in the colder part of the tube. When heated with sodic carbonate on charcoal in the reducing flame of the blowpipe, the arsenic compounds evolve a char- acteristic odor of garlic. NIOBIUM, Nb, and TANTALUM, Ta. Atomic weights :Nb = 94, Ta ~ 182. Atomicity and T. _ Occurrence.—These two very rare elements generally occur together as tantalates and niobates. Preparation.—Very little is known of them in the free state. They may he obtained as black powders by heating potassic niobofluoride and potassic tantalofluoride with potassium or sodium. The following are some of the principal compounds of these elements; 378 INORGANIC CHEMISTRY. Niobic chloride, . . . NbCl5. Niobic oxytrichloride, . . . . . . NbOCIs. Niobic fluoride, . . . NbF2. Potassic niobofluoride, . . . . . . NbK2F7. Hyponiobous oxide, .... . . . /Nb//202. Hyponiobic oxide, .... . . . /Nb,'204. Niobic anhydride, .... . . . NbiOg. COMPOUNDS OF NIOBIUM. COMPOUNDS OF TANTALUM. These correspond with the above compounds of niobium, with the single exception that hypotantalous oxide has not been prepared. ANTIMONY, Sb4 ? Atomic weight = 120. Probable molecular weight = 480. Sp. gr. {crys- talline) 6.7, (amorphous) 5.78. Fuses at 430° C. (806° F.). Ato- micity and v. of atomicity: Antimonious chloride, Sb"'Cl8. Antimonic tetretho-chloride (tetr ethyl' 1 SbvEt4CI. stibonic chloride), j Antimonic chloride, SbvCl5. History.—Many of the compounds of antimony have been known from very early times. In the fifteenth century Basil Valentine de- scribed the preparation of metallic antimony. Occurrence.—Antimony is rarely found native. Sometimes it occurs alloyed with other metals in various minerals. Combined with oxygen, it occurs as valentinite or antimonious oxide, and as antimony ochre or diantimonious tetroxide. The chief source of antimony is gray anti- mony ore or stibnite, which consists of antimonious sulphide. Sulph- antimonites also occur, such as miargyrite, an argentic raeta-sulphanti- monite. «. Crystalline Antimony. Preparation.—l. Antimony may be obtained by fusing the powdered native sulphide with iron filings: Sb2S"3 + 3Fe = 3FeS" + 2Sb. Antimonious Ferrous sulphide. sulphide. 2. In preparing antimony on a large scale, the sulphide is first roasted in contact with air, when it is partially converted into antimo- nious oxide: 2Sb2S"3 + 902 = 2Sb2G3 + 6SQ2. Antimonious Antimonious Sulphurous sulphide. oxide. anhydride. ANTIMONY. 379 The roasted mineral is then fused with charcoal and sodic carbonate. The reaction takes place in two stages: first,, the remaining sulphide is converted into oxide by the sodic carbonate, and subsequently the oxide 18 reduced by the carbon: 1. Sb2S"3 + 3CONao2 = 3C02 + 3SNa2 + Sb203. Antimonious Sodic Carbonic Sodic Antimonious sulphide. carbonate. anhydride. sulphide. oxide. 2. Sb203 + 3C = 3CO + Sb2 Antimonious Carbonic oxide. oxide. 3. Pure antimony may be obtained by reducing with charcoal the oxide formed by the action of nitric acid upon crude antimony. Properties.—Antimony is a bluish-white lustrous metal, with a crys- talline fracture. By slow cooling it may be obtained in rhombohedra, closely approximating to the cube. It fuses at 480° C., and may be dis- tilled at a white heat. /?. Amorphous Antimony. Preparation.—This variety, discovered by Gore {Phil. Trans., 1858, p. 185), is obtained by the electrolysis of a solution of tartar emetic in antimonious chloride. Properties.—Amorphous antimony has the appearance and lustre of polished steel, with a peculiar mammillated surface, and an amorphous fracture. It contains sor 6 per cent, of antimonious chloride derived from the electrolyte. When heated or struck it undergoes a molecular change, which spreads rapidly through the entire mass and is attended 'vith a rise of temperature from 15° to 250° C. At the same time fumes of antimonious chloride are evolved. After this change the metal 18 found to possess an increased density and to have acquired a granular fracture. Reactions.—l. When antimony is heated to redness in air it burns, forming antimonious oxide. If a small quantity of antimony be heated on charcoal to its point of ignition, and then thrown on to a large sheet of paper folded into the form of a tray, the metal breaks up into a number of globules, which dance about on the surface of the paper, burning brilliantly, and leaving black intermittent streaks behind them. 2. With hot concentrated sulphuric acid it yields antimonious sul- phate with evolution of sulphurous anhydride: 2Sb + 6S02Ho2 = S306{Sb206)Ti + 3S02 + 60H2. Sulphuric Antimonious Sulphurous Water, acid. sulphate. anhydride. Uses.—Metallic antimony is employed only in the form of its alloys, to which it imparts the valuable property of expanding on solidification. This renders them especially suitable for taking sharp casts. The most important alloys containing antimony are type metal and Britannia metal {q.v.). 380 INORGANIC CHEMISTRY. COMPOUND OF ANTIMONY WITH HYDROGEN ANTIMONIURETTED HYDROGEN, Antimonious Hydride. Molecular weight = 123. SbH3. This compound is unknown in the pure condition. Preparation.—l. It is prepared by the action of hydrochloric acid upon an alloy of zinc and antimony: Sb2Zn"3 -f 6HCI = 2SbH3 + 3ZnCI2. Antimonious Hydrochloric Antimonious Zincic zincide. acid. hydride. chloride. 2. It is formed by the action of nascent hydrogen, evolved from zinc and sulphuric acid, upon soluble antimony compounds: SbCI3 + 3H2 = SbH3 + 3HCI. Antimonious Antimonious chloride. hydride. In both these reactions the antimonious hydride is always mixed with much hydrogen. Properties.—lt is a colorless gas, possessing a most offensive odor. It burns with a bluish flame. Reactions.—1. When burnt in air or oxygen it yields water and an- timonious oxide: 2SbH3 + 302 = Sb2Os + 30H2. Antimonious Antimonious Water, hydride. oxide. 2. When burnt with a limited supply of air, the hydrogen alone is oxidized, the antimony being deposited: 4SbH3 + 302 = Sb4 + 60H2. Thus, if a cold surface of porcelain be held in the flame of antimo- nious hydride, a dull black spot of metallic antimony is formed. 3. When passed through a red-hot tube, it is decomposed, like ar- senious hydride, into its elements. 4. When transmitted through a solution of argentic nitrate, it pro- duces a black precipitate of antimonious argentine, thus differing from arsenious hydride (p. 368): 3N02Ago + SbHg = 3NOaHo + SbAg3. Argentic Antimonious Nitric Antimonious nitrate. hydride. acid. argentide. ANTIMONIOUS CHLORIDE. 381 From the composition of this compound, and from that of some of its analogues, the composition of antimonious hydride is inferred. Antimonious hydride, . SbH.j.. Antimonions bromide, . SbBr3. Antimonious argentide, . SbAg3. Antimonious zincide, . Sb.,Zn",. Antimonious ethide (Triethylstibine), . Antimonious amylide (Triamylstibine), . SbEt3. . SbAy3. COMPOUNDS OF ANTIMONY WITH THE HALOGENS. Antimonious chloride, .... .... SbCl,. Antimonic chloride, .... SbCl5. Antimonious bromide, .... Antimonious iodide, .... Sbl, Antimonious fluoride, .... .... SbF3. Antimonic fluoride, .... SbF, ANTIMONIOUS CHLORIDE. SbCI3. Molecular weight 226.5. Molecular volume i I I. 1 litre of antimo- nious chloride vapor weighs 113.25 criths. Fuses at 72° C. 1161.6° F.). Boils at 223° C. (433.4° F.). over excess of metallic antimony or antimonious sulphide; Preparation.—l. This compound is formed when chlorine is passed 2Sb + 3C12 == 2SbCl3. Antimonious chloride. 2Sb2S"3 + 9C12 = 4SbCl3 + 3'S'2C12. Antimonious Antimonious Disulphur sulphide. chloride. dichloride. The product must be purified by distillation. 2. It may also be prepared by dissolving antimonious sulphide in hydrochloric acid, or antimony in aqua-regia, evaporating, and distill- ing the product: Sb2S"3 + 6HCI = 3SH2 + 2SbCl3. Antimonious Hydrochloric Sulphuretted Antimonions sulphide. acid. hydrogen. chloride. 2Sb + 6HCI + 6N02Ho = 2SbCl3 + 60H2 + 3'Niv20,. Hydrochloric Nitric Antimonious Water. Nitric acid. acid. chloride. peroxide. 382 INORGANIC CHEMISTRY. The receiver must be changed as soon as the distillate begins to solidify, and the product which is collected above this point must be purified by repeated rectification. 3. It may be conveniently obtained by the distillation of a mixture of 1 part of powdered antimony with 3 parts of mercuric chloride, or of 3 parts of antimonious sulphide with 7 parts of mercuric chloride: Sb4 + 4HgCl2 == 2SbCl3 + 'Sb"2Hg" + 'Hg^Clj. Mercuric Antimonious Dimercuric Mercurous chloride. chloride. diantimonide. chloride. Sb2S" + 3HgCl2 = 2SbCl3 + 3HgS. Antimonious Mercuric Antimonious Mercuric sulphide. chloride. chloride. sulphide. 4. Another method consists in distilling together antimonious sul- phate and sodic chloride: S306Sbo'"2 + GNaCI = 2SbCl3 + 3S02Nao2. Antimonious Sodic Antimonious Sodic sulphate. chloride. chloride. sulphate. Properties.—Antimonious chloride is a soft, colorless, laminated crys- talline mass. From its consistency and fusibility, it was formerly known as butter of antimony. It is deliquescent and powerfully cor- rosive. Reaction.—With water it produces antimonious oxychloride, which is thus obtained as a white powder: SbCl3 + 0H2 = 2HCI + SbOCl. Antiraonious Water. Hydrochloric Antimonious chloride. acid. oxychloride. Long-continued action of water transforms this compound into anti- monious oxide; 2SbOCI + 0H2 = 2HCI + Sb203. Antiraonious Water. Hydrochloric Antimonious oxychloride. acid. oxide. ANTIMONIC CHLORIDE. SbCl5. Molecular weight = 297.5. Fuses at 0° C. Preparation.—It is obtained by acting upon antimony with excess of chlorine, or by passing this gas over antimonious chloride, when the latter liquefies, producing antimonic chloride: OXIDES AND ACIDS OF ANTIMONY. 383 2Sb + 5C12 = 2SbCl5. Antimonic chloride. SbCl3 + Cl2 = SbCl5. Antimonious Antimonic chloride. chloride. Properties.—Antimonic chloride is a colorless, fuming liquid. It is readily decomposed on heating into antimonious chloride and free chlorine, and thus behaves towards many substances as a chlorinating agent. Reactions.—1. With a small quantity of water, it forms antimonic oxytrichloride, analogous to phosphoric oxytrichloride: SbCl5 + OH2 = SbOCl3 + 2HCI. Antimonic Water. Antimonic Hydrochloric chloride. oxytrichloride. acid. 2. An excess of water transforms antimonic chloride into orthanti- monic acid, or pyrantimonic acid corresponding to pyrophosphoric acid : SbC]5 + 40H2 = SbOHog + SHCI? Antimonic Water. Orthantimonic Hydrochloric chloride. acid. acid. or 2SbCl5 + 70H2 = Sb203Ho4 + 10HC1. Antimonic Water. Pyrantimonic Hydrochloric chloride. acid. acid. 3. By the action of sulphuretted hydrogen, antimonic sulphotri- chloride is formed: SbCl5 + SH2 = SbS"Cl3 + 2HCI. Antimonic Sulphuretted Antimonic Hydrochloric chloride. hydrogen, sulphotrichloride. acid. Antimonious bromide, Sbßr3, resembles antimonious chloride. It crystallizes from carbonic disulphide in colorless octahedra. It fuses at 90° C. (194° F.), boils at 275° C. (527° F.), and by the action of water is converted into the oxybromide, SbOßr. Antimonious iodide, Sbl3, crystallizes in red hexagonal plates, and, when acted upon by water, forms the oxyiodide, SbOI. Antimonious fluoride, SbF3, is obtained by evaporating a solution of antimonious oxide in excess of hydrofluoric acid. It crystallizes in colorless prisms or scales, and deliquesces with formation of the oxy fluoride, SbOF. Antimonic fluoride, SbF5, is left behind as a gummy mass when a solution of anti- monic acid in hydrofluoric acid is evaporated in vacuo. OXIDES AND ACIDS OF ANTIMONY. Antimonious oxide or anhydride, . . . (Sb203)2. Diantiraonic tetroxide, . . . . 'Sb^O^. Antimonic anhydride, . . Sb2Og. Metantiraonious acid, . . SbOHo. Orthantimonic acid, . . SbOHog ? Metantimonic acid, . . SbO„Ho. Pyrantimonic acid, ...... . • Sb203Ho4. 384 INORGANIC CHEMISTRY. ANTIMONIOUS OXIDE, or ANHYDRIDE. (5b203)2. Molecular weight 576. Molecular volume I I I. 1 litre of antimo- nious oxide vapor weighs 288 criths. Sp. gr. {octahedral) 5.25, {rhombic) 5.55. Occurrence.—Antimonious oxide is found in nature in two rare minerals: in the rhombic form as valentinite, and in the octahedral form as senarmontite. Preparation.—1. It is formed when antimony is burnt in air or oxygen. 2. It is most readily obtained by pouring a solution of antimonious chloride in dilute hydrochloric acid into a boiling solution of sodic carbonate: 2SbCl3 + 3CONao2 = Sb203 + 3C02 + 6NaCI. Antimonious Sodic Antimonions Carbonic Sodic chloride. carbonate. oxide. anhydride. chloride. 3. When raetantimonious acid is heated to 100° C., it is converted into antimonious anhydride, with elimination of the elements of water: 2SbOHo = 0H2 + Sb203. Metantiraonious Water., Antimonious acid. oxide. Properties.—Antimonious anhydride may be obtained in two dis- tinct crystalline forms—in rhombic prisms and in regular octahedra— corresponding with the two forms of arsenious anhydride, with which substance it is therefore isodimorphous. When antimony is heated in a slow current of air, rhombic prisms of the oxide are formed in the immediate neighborhood of the metal; further on a mixture of prisms and octahedra is deposited ; whilst in the colder parts of the tube the crystals consist of octahedra alone. Antimonious oxide in both its forms is colorless, but when heated, assumes a yellow tint, which disap- pears again on cooling. When air is excluded, it may be fused and sublimed. Water does not dissolve it. Reactions.—1. When heated to redness in air, it burns like tinder, forming diantimonic tetroxide: Sb2Os + O = 'Sbiv204. Antimonious Diantimonic oxide. tetroxide. 2. It is really reduced to the metallic state by ignition with charcoal or hydrogen: METANTIMONIOUS ACID DIANTIMONIC TETROXIDE. 385 Sb203 + 3C = 2Sb + 3CO. Antimonious Carbonic oxide. oxide. Sb2Oa + 3H2 = 2Sb + 80H2. Antimonious Water. oxide. 3. It is readily dissolved by a hot solution of hydric potassic tartrate (cream of tartar), forming potassic antimonylic tartrate (tartar emetic) : f coho rco(Sb'"o2) 0 ] CHHo , no j CHHo , att 2 I CHHo + Sb2°3 2 ] CHHo + OH2” 1 COKo [OOKo Hydric potassic Antimonious Potassic antimony lie Water, tartrate (cream oxide. tartrate (tartar of tartar), emetic). 4. Hydrochloric acid dissolves it with formation of antimonious chloride: Sb2Os + 6HCI = 2SbC]3 + 30H2. Antimonious Hydrochloric Antimonious Water, oxide. acid. chloride. METANTIMONIOUS ACID. Molecular weight 153. SbOHo. Preparation.—Metantimonious acid is obtained by pouring a solution of antimonious chloride into a cold solution of sodic carbonate: 2SbCl3 + 3CONao2 + OH, = 2SbOHo + 3C02 + 6NaCI. Antimonious Sodic Water. Metantimonious Carbonic Sodic chloride. carbonate. acid. anhydride. chloride. Properties.—lt forms a white amorphous powder, which is insoluble in water. Reactions.—l. It is decomposed by heat (p. 384). 2. It is dissolved by an excess of alkaline hydrates, producing ill-defined anti- Wonites. It also possesses weak basic properties and forms salts in which the monad group (SbO) replaces the hydrogen of the acid. Potassic antimonylic tartrate is an ex- ample. DIANTIMONIC TETROXIDE, Antimonylic Antimonate. {sibo! = 'Sb‘'A, or Sb'02(Sb'"02). Molecular weight = 304. Occurrence.—Diantimonic tetroxide is found native as cervantite. Preparation.—l. It is obtained by igniting antimonic anhydride, or the white solid produced by the action of nitric acid upon metallic antimony; 386 INORGANIC CHEMISTRY. 2SbA = 2/Sb!v204 + 02. Antimonic Diantimonic anhydride. tetroxide. 2. When antimonious oxide is heated in contact with air, it is con- verted into diantimonic tetroxide (p. 384). Properties.—Diantimonic tetroxide is a white, infusible and non- volatile powder. When heated, it turns yellow, but becomes white again on cooling. Reaction.—When boiled with a solution of hydric potassic tartrate, it is decomposed, potassic antimonylic tartrate and metantimonic acid being formed: f coho r co(Sb'^o2) ] OHHo + = j + Sb02Ho. L COKo L COKo Hydric potassic Antiraonylic Potassic antimonylic Metantimonic tartrate. antimonate. tartrate. acid. This reaction seems to indicate that this oxide is in reality an anti- monylic antimonate as formulated in the above equation. ANTIMONIC ANHYDRIDE. Sb205. Molecular weight 320. Sp. gr. 6.6. Preparation.—lt is obtained by heating the corresponding acids to 280° C. (536° F.): 2SbOHo3 == Sb205 + 30H2. Orthantimonic Antimonic Water, acid. anhydride. 2Sb02Ho = Sb2Os + OH2. Metantimonic Antimonic Water, acid. anhydride. Sb203Ho4 = SbA + 20H2 Pyrantimonic Antimonic Water, acid, anhydride. Properties.—Antimonic anhydride is a pale yellow amorphous sub- stance, insoluble in water. Reactions.—1. When heated it is decomposed into antimonylic anti- monate and oxygen {supra). This decomposition begins at 300° C. 2. Fused with potassic carbonate, it produces potassic metanti- monate: Sb205 + COKoa = 2Sb02Ko + C02. Antimonic Potassic Potassic Carbonic anhydride. carbonate. metantimonate. anhydride. COMPOUNDS OF ANTIMONY WITH SULPHUR. 387 ORTHANTIMONIC ACID. SbOHog ? Preparation.—This acid is said to be formed by the action of water upon antimonic chloride (p. 383). METANTIMONIC ACID. Preparation.—l. It is obtained by the action of nitric acid containing a little hydro- chloric acid on metallic antimony : Sb02Ho. Sb2 + 4N02Ho = 2SbOsHo + N203 + 'll",©, + OH2. Nitric Metantimonic Nitrous Nitric Water, acid. acid. anhydride. oxide. 2. It is also formed by the spontaneous dehydration of orthantimonic acid, or of pyrantimonic acid SbOHog = OH2 + Sb02Ho. Orthantimonic Water. Metantimonic acid. acid. Sb203Ho4 = OII2 -j- 2Sb02Ho. Pyrantimonic Water. Metantimonic acid. acid. Properties.—lt is a soft white powder, sparingly soluble in water. The solution reddens litmus. Reaction.—By the action of alkaline hydrates, it produces either metantimonates or orthantimonates: Sb02Ho + = Sb02Ko + OH2. Metantimonic Potassie Potassic Water, acid. hydrate. metantimonate. Sb02Ho + OKH = SbOHo2Ko. Metantimonic Potassic Dihydrie potassie acid. hydrate. orthantimonate. PYRANTIMONIC ACID, Parantimonic Acid (Metantimonic Acid of Fremy). Sb2OgHo4. Preparation.—l. It is formed by the action of water upon antimonie chloride (p. 383). 2. It is also obtained by acidifying solutions of pyrantimonates: Sb203Ho2Ko2 + 2HCI = Sb20,Ho4 + 2KCI. Dihydrie dipotassic Hydrochloric Pyrantimonic Potassic pyrantimonate. acid. acid. chloride. Dihydrie dipotassic pyrantimonate is prepared by fusing antimonic anhydride with an excess of potassic hydrate, and extracting the mass with water, when an alkaline solution containing dihydrie dipotassic pyrantimonate, Sb203Ho2Ko2, is obtained. This solution produces precipitates in solutions of sodium salts, the sodic pyranti- monate thus formed having the formula Sb203Ho2Nao2,60H2. COMPOUNDS OF ANTIMONY WITH SULPHUR. Antimonious sulphide, . . . . Sb2S"3. Antimonic sulphide, . - , . . Sb2S"5. 388 INORGANIC CHEMISTRY. ANTIMONIOUS SULPHIDE, B'ulphantimonious Anhydride. Molecular weight = 336. Sb2S"3. Occurrence.—Antimonious sulphide is found in nature as stibnite or gray antimony ore. Preparation.—l. It may be obtained by heating together antimony and sulphur, or antimonious oxide and sulphur, in the proper molec- ular proportions: Sb4 + 3S2 = 2Sb2S"s. Antimonious sulphide. 2Sb2Oa + 9S = 2Sb2S"3 + 3S02. Antimonious Antimonious Sulphurous oxide. sulphide. anhydride. 2. It is precipitated when sulphuretted hydrogen is passed through a solution of antimonious chloride: 2SbCl3 + 3SH2 = Sb2S"3 + 6HCI. Antimonious Sulphuretted Antimonious Hydrochloric chloride. hydrogen. sulphide. acid. Properties.—The native sulphide occurs in dark-gray radiating crys- talline masses, with a metallic lustre—less frequently in rhombic prisms. The precipitated substance is an orange-red amorphous powder, con- taining water of hydration which may be expelled by heating. Anti- monious sulphide is readily fusible, and may be sublimed. Reactions.—1. Hot hydrochloric acid decomposes it, forming anti- monious chloride and sulphuretted hydrogen (see p. 381). 2. It dissolves with decomposition in solutions of alkaline hydrates, yielding a mixture of antimonite and sulphantimonite: Sb2S"3 + 6KHO = SbKs3 + SbKo3 + 30H2. Antimonious Potassic Trisulphopotassic Tripotassic Water, sulphide. hydrate. sulphantimonite. antimonite. The addition of an acid reproduces and precipitates the antimonious sulphide: SbKs3 + SbKo3 + 6HCI = Sb2S"3 Trisulphopotassic Tripotassic Hydrochloric Antimonious sulphantimonite. antimonite. acid. sulphide. + 6KCI + 30 H2. Potassic chloride. Water. 3. It dissolves in a solution of an alkaline sulphhydrate, forming a sulphantimonite: Sb2S"3 + 6KHs == 2SbKs3 + 3SH2. Antimonious Potassic Trisulphopotassic Sulphuretted sulphide. sulphhydrate. sulphantimonite. hydrogen. SULPHANTIMONITES—ANTIMONIC SULPHIDE. 389 SULPHANTIMONITES. Many sulphantimonites occur in nature : Orthosulphantimonites. Dark-red silver. Trisulphargentic sulphantimonite, . SbAgss. Boulangerite. Trisulphoplumbic sulphantimonite, . . Sb2Pbs//3. Bonrnonite. Disulphoplumbic sulphocuprous sulphanti- monite, Sb2Pbs//2(Cu2S//2)//. General formulae: SbMs3 and Sb2Ms//3. Metasulphantimonites. Miargyrite. Sulphargentic metasulphantimonite, . . SbS//Ags. Zinkenite. Sulphoplumbic metasulphantimonite, . . Sb2S//2Pbs//, Antimony copper glance. Sulphocuprous metasulph- antimonite, Sb-S".(Cu.S".)". Berthierite. Sulphqferrous metasulphantimonite, . . Sb2S//2Fes//. General formulas: SbS//Ms and Sb2S//2Ms//. JPyrosulphantimonites. General formulae: Sb2S//Ms1 and Sb2S//Ms//2. Feather ore- Sulphoplumbic pyrosulphantimonite, . . Sb2S//Pbs//2. Fahl ore. Sulphocuprosoferrous pyrosulphantimonite,. Sb2S//(Cu2FeS//2)//, A soluble colloidal modification of antimonious sulphide corresponding with colloidal arsenious sulphide (p. 375) is also known. ANTIMONIC SULPHIDE, Sulphantimonic Anhydride. Sb2S", Molecular weight 400 Preparation.—l. It is precipitated as a yellowish-red powder when sulphuretted hydrogen is passed through a solution of antimonic chloride : 2SbClg + 5SH2 = Sb2S" + 10HC1. Antimonic Sulphuretted Antimonic Hydrochloric chloride. hydrogen. sulphide. acid. 2. The same precipitate is formed by the addition of an acid to a solution of a sulphantimonate: 2SbS"Nas3 + 6HCI = Sb2S" + 6NaCI + 3SH2. Trisulphosodic Hydrochloric Antimonic Sodic Sulphuretted sulphantimonate. acid. sulphide. chloride. hydrogen. Reactions.—l. When heated it is decomposed into antimonious sul- phide and free sulphur. 2. Boiling hydrochloric acid decomposes it into antimonious chloride, sulphuretted hydrogen, and sulphur : 390 INORGANIC CHEMISTRY. Sb2S" + 6HCI = 2SbCl3 + 3SH2 + S2. Antimonic Hydrochloric Antimonious Sulphuretted sulphide. acid. chloride. hydrogen. 3. It dissolves in a solution of an alkaline sulphide, forming a sulph- antirnonate: Sb2S"5 + 3SK2 = 2SbS//Ks3. Antimonic Potassic Trisulphopotassic sulphide. sulphide. sulphantimonate. 4. It is soluble in a solution of an alkaline hydrate, a mixture of antimonate and sulphantimonate being formed : 4Sb2S" + 240KH = 3SbOKos + SSbS"Ks3 + 120H2. Antimonic Potassic Tripotassic Trisulphopotassic Water, sulphide. hydrate. antimonate. sulphantimonate. General Properties and Reactions of the Compounds of Antimony : Antimonious Compounds.—Solutions of antimonious oxide in acids became milky on dilution with water. The milkiness disappears on addition of tartaric acid. (Distinction from bismuth compounds.) Sul- phuretted hydrogen precipitates from acid solutions orange-colored anti- monious sulphide, soluble in concentrated hydrochloric acid, in caustic alkalies, and in ammonic sulphide, almost insoluble in amnionic car- bonate, insoluble in hydric potassic sulphite. If a hydrochloric acid solution of the sulphide or of any other compound of antimony be brought into a platinum dish along with a piece of zinc, the antimony is deposited by voltaic action as a black coating adhering to the plati- num, whilst any tin which may be present is precipitated as a gray powder on the zinc. The hydrochloric acid solution of an antimonious compound precipitates gold in the metallic form from its solutions. Antimonic Compounds.—These yield in acid solution with sulphuretted hydrogen a yellowish-red precipitate of antimonic sulphide which is soluble in the same reagents as the antimonious compound. The compounds of antimony when introduced into Marsh’s apparatus (p. 377) evolve antimoniuretted hydrogen. The flame of this gas de- posits, upon a cold surface of porcelain, a stain of metallic antimony, which is blacker and less lustrous than that of arsenic. A mirror of metallic antimony is also formed when the gas is passed through a heated tube. These coatings maybe distinguished from those of arsenic by their almost total insolubility in sodic hypochlorite. When heated with potassic cyanide upon charcoal in the reducing flame of the blow- pipe, compounds of antimony yield a brittle metallic regulus, and the charcoal becomes covered with a white incrustation; but no odor of garlic is perceptible as in the case of arsenic. BISMUTH—BISMUTHOUS CHLORIDE. 391 BISMUTH, Bi4? Atomic weight 208.2. Sp. gr. 9.83. Fuses at 265° C. Atomicity and v. Evidence of atomicity: Bismuthous chloride, .... . . . Bi"'Cl3. Bismuthons oxide, • • • Bi'"203. Bismuthous ethide, . . . Bi'"Et3. Bismuthons dichlorethide, . . . . . . Bi'"EtCl2. Bismuthic anhydride, .... . . . Bi'A- Metabismuthic acid, . . . BE02Ho. Occurrence.-—Bismuth is found principally in the metallic state, but it also occurs in combination with oxygen, sulphur, and tellurium. Preparation.—l. The method of extraction from the ores formerly consisted in heating the crude native bismuth in sloping iron tubes placed in a furnace. The metal fused and ran off, whilst the impuri- ties were left in the tubes. The bismuth thus obtained was contami- nated with sulphur, arsenic, iron, and other metals. 2. At the present day large quantities of bismuth are obtained as a by-product in the manufacture of smalt (q.v.). The crude bismuth is purified by fusing at the lowest possible temperature, when the more fusible bismuth runs off, leaving the iron, nickel, and other impurities behind. 3. It may be obtained in the pure state by dissolving commercial bis- muth in nitric acid, precipitating the basic nitrate by the addition of water, and reducing the precipitate by ignition with charcoal. Properties.—Bismuth is a grayish-white metal with a slight reddish tinge. It crystallizes in rhombohedra which approximate closely to the cube. At a very high temperature it volatilizes. It is not oxid- ized by exposure to the air at ordinary temperatures, but, when strongly heated, burns, forming bismuthous oxide. Uses.—Metallic bismuth is employed in the preparation of fusible alloys, such as Rose’s metal and Wood’s metal {q.v.). No compound of bismuth with hydrogen is known. HALOGEN AND OXYHALOGEN COMPOUNDS OF BISMUTH. BISMUTHOUS CHLORIDE. BiCl3. Molecular weight 314.7. Molecular volume I 11. 1 litre of bis- muthous chloride vapor weighs 157.35 criths. Preparation.—l. It is formed when dry chlorine is passed over metallic bismuth; 392 INORGANIC CHEMISTRY. Bi2 + 3€I2 = 2BiCl3. Bismuthous chloi’ide. 2. It may be obtained by evaporating a solution of bismuth in hydro- chloric acid containing a little nitric acid, and distilling the residue. 3. Another method consists in distilling bismuth with mercuric chloride : Bi2 + 6HgCl2 = 2BiCl3 + S'Hg'.Cl,. Mercuric Bismuthous Mercurous chloride. chloride. chloride. Properties.—lt forms a white fusible deliquescent mass which may be distilled. Reaction.—Water decomposes it, precipitating bismuthous oxychloride as a white powder: BiClg + 0H2 = BiOCl + 2HCI. Bismuthous Water. Bismuthous Hydrochloric chloride. oxychloride. ecid. f BiCl Dibismuthous tetrachloride, < is obtained as a black amorphous mass by heating bismuthous chloride with bismuth. Bismuthous bromide, Bißr3, forms yellow prisms fusing at 200° C. Water converts it into bismuthous oxybromide, BiOßr. Bismuthous iodide, Bil3, is obtained by heating a mixture of bismuth and iodine. It sublimes in lustrous, dark-gray hexagonal plates. By boiling with water it is de- composed into hydriodic acid and copper-colored bismuthous oxyiodide, BiOI. Bismuthous fluoride,BBis, is obtained as a white powder by evaporating a solu- tion of bismuthous oxide with an excess of hydrofluoric acid: Bi203 + 6HF = 2BiF3 + 30H2. Bismuthous Hydrofluoric Bismuthous Water, oxide. • acid. fluoride. COMPOUNDS OF BISMUTH WITH OXYGEN AND HYDROXYL. Dibismuthous dioxide, BiO BiO' Bismuthous oxide, ' Bi203. Dibismuthic tetroxide, /Biiy204. Bismuthic anhydride, Bi205. Bismuthous oxyhydrate, or metabismuthous acid, BiOHo. Metabismuthic acid, Bi02Ho. DIBISMUTHOUS DIOXIDE. / BiO \ BiO’ Molecular weight 448.4. Preparation.—When a mixture of stannous and bismuthous chlorides is poured into an excess of dilute caustic potash, a black precipitate of dibismuthous dioxide is formed. The reaction takes place in two stages. In the first, dipotassic stannite and tripotassic bisrnuthite are formed; these salts then react upon each other: BISMUTHOUS OXIDE. 393 2BiKo3 + SnKo2 + SOH2 = /B1"2Oj + SnOKo2 + 60KH. Tripotassic Dipotassic Water. Dibismuthous Dipotassic Potassic hismuthite. staimite. dioxide. stannate. hydrate. The precipitate of dibismuthous dioxide must be filtered and washed out of contact with air and then dried by heating in a current of carbonic anhydride. It is thus obtained as a gray crystalline powder. Reactions.—1. The moist substance when exposed to the air oxidizes spontaneously to bismuthous oxide: /Bi//202 + O - Bi203. Dibismuthous Bismuthous dioxide. oxide. In the same way when the dried compound is heated in the air, it glows like tinder and is converted into bismuthous oxide. 2. Hydrochloric acid decomposes it into bismuthous chloride and bismuth: 3/Bi//202 + 12HC1 = 4BiCls + Bi2 + 60H2. Dibismuthous Hydrochloric Bismuthous Water, dioxide. acid. chloride. BISMUTHOUS OXIDE Bi203. Molecular weight = 464.4. Sp. gr. 8.2 Occurrence.—This substance is found in nature as the rare mineral bismuth ochre. Preparation.—l. It is formed when bismuth is burnt in air or oxygen. 2. It is left behind when the nitrate, carbonate, or hydrate is heated: 2N306Bio/// = Bi203 -f- 3N203 -f- 302. Bismuthous Bismuthous Nitrous nitrate. oxide. anhydride. 2N02(Bi'//Ho20) = Bi203 + 20H2 + N2Os + Oa. Bismuthous nitrate Bismuthous Water. Nitrous dihydrate. oxide. anhydride. 0O(BiO2)2 Bi203 ~b 0O2. Eismuthylic Bismuthous Carbonic carbonate. oxide. anhydride. 2BiOHo = Bi2Os -(- OH2. Bismuthous Bismuthous Water, oxyhydrate. oxide. 3. When bismuthous hydrate is dissolved in a solution of potassic hydrate and boiled, it parts with the elements of water, and is precipi- tated as bismuthous oxide. Properties.—Bismuthous oxide is a yellow insoluble powder, which becomes darker on heating, and then fuses. The oxide obtained by boiling the solution of the hydrate in caustic alkali is crystalline. 394 INORGANIC CHEMISTRY. Reaction.—lt is dissolved by hydrochloric, nitric, and sulphuric acids, forming the bismuthous chloride, nitrate, and sulphate: Bids- N306Bio/". S306Bio'"2. Bismuthous Bisrauthous Bismuthous chloride. nitrate. sulphate. Salts of Bismuthous Oxide.—These salts are soluble only in water containing an excess of acid. Pure water decomposes them into basic salts and free acid; NgOgßio'" + 20H2 = N02(Bi"'Ho20) + 2N02Ho. Bismuthous Water. Bismuthous nitrate Nitric nitrate. dihydrate. acid. BISMUTHOUS OXYHYDRATE, Metabismuthous Acid. BiOHo. Preparation.—By pouring a solution of bismuthous nitrate in dilute nitric acid into dilute ammonia or potassic hydrate, a precipitate is formed, which probably contains orthobismuthous acid. On drying this precipitate at 100° C., metabismuthous acid is obtained as a white amorphous mass: N306Bio'" + 30KH = BiHo3 + 3N02Ko. Bismuthous Potassic Orthobismuthous Potassic nitrate. hydrate. acid. nitrate. BiHog = BiOHo + OH2. Orthobismuthous Metabismuthous Water, acid. acid. Reaction.—By heat or by boiling with caustic alkali, water is expelled and bismuthous oxide is formed (see p. 393). An unstable metabismuthite is produced by fusing bismuthous oxide with sodic carbonate; Bi203 + CONao2 = 2BiONao + C02. Bismuthous Sodic Sodic Carbonic oxide. carbonate. metabismuthite. anhydride. BISMUTHIC ANHYDRIDE. BiA- Preparation.—This compound is obtained as a brown powder by beating bismuthic acid to 130° C. Reactions.—1. When heated to the boiling point of mercury it loses oxygen, being converted either into bismutbous oxide or into dibismuthic tetroxide: METABISMUTHIC ACID—DIBISMUTHOUS DISULPHIDE. 395 BL205 = Bi203 + O2 j Bismuthic Bismuthous anhydride. oxide. 2Bi205 = 2/Bi*,204 + 02 Bismuthic Dibismuthic andydride. tetroxide. 2. When heated in a current of hydrogen, it is readily reduced to bismuthous oxide. 3. Heated with hydrochloric acid it evolves chlorine, producing bismuthous chloride and water: Bi205 + 10HC1 = 2BiCl3 + 50 H2 + 2C12. Bismuthic Hydrochloric Bismuthous Water, anhydride. acid. chloride. 4. Sulphurous acid converts it into bismuthous sulphate: 3Bi205 + 6SOHo2 = 2S306Bio///2 + Bi203 + 60H2. Bismuthic Sulphurous Bismuthous Bismuthous Water, anhydride. acid. sulphate. oxide. 5. When heated with sulphuric or nitric acid it evolves oxygen, producing bis- muthous sulphate or nitrate: Bi205 + 3S02Ho2 = S306Bio///2 + SOH2 + 02; Bismuthic Sulphuric Bismuthous Water, anhydride. acid. sulphate. Bi205 + 6N02Ho = 2N306Bio/// + 30H2 + 02. Bismuthic Nitric Bismuthous Water, anhydride. acid. nitrate. METABISMUTHIC ACID. Bi02Ho. Preparation.—Metabismuthic acid is obtained as a red deposit by passing chlorine through a solution of potassic hydrate containing bismuthous oxide in suspension; 4CKH ■+• 2C12 + Bi203 = 2Bi02Ho + 4KCI + OH2. Potassic Bismuthous Metabismuthic Potassic Water, hydrate. oxide. acid. chloride. Reaction.—It dissolves in a hot solution of potassic hydrate. By the addition of an acid to the liquid a salt, said to have the composition Bi204HoKo, is precipitated. COMPOUNDS OF BISMUTH WITH SULPHUR. Dibismuthous disulphide, . . • . 'Bi"2S"2. Bismuthous sulphide, . . . . . . Bi2S;/3. DIBISMUTHOUS DISULPHIDE {SP or /Bi//2S//2. Sp. gr. 7.8, Preparation.—Dibismuthous disulphide is obtained as a mass of gray, metallic acicular crystals by fusing together bismuth and sulphur in the proper molecular proportions. 396 INORGANIC CHEMISTRY. BISMUTHOUS SULPHIDE. Bi2S"3. Sp. gr. 6.4. Occurrence.—Bisrauthous sulphide is found native as the rare mineral bismuth glance. It forms rhombic crystals and foliated or fibrous masses with a metallic lustre. Preparation.—l. It maybe obtained by fusing together bismuth and sulphur in the proper molecular proportions. 2. It is also obtained as a blackish-brown powder by precipitating bismuth solutions by sulphuretted hydrogen : 2BiCl3 + 3SH2 = Bi2S"3 + 6HCI. Bisrauthous Sulphuretted Bisrauthous Hydrochloric chloride. hydrogen. sulphide. acid. Reaction.—This compound is not dissolved by alkaline hydrates or sulphydrates. A few sulphobismuthites are found in nature: Kobellite. Trisulphoplmibic ) -o- pi ~ sulphobismuthite j 2 S3* Needle ore. Disulphoplumhie- 1 Pbs" f'Cu' SV7 dieuprous sulphobismuthite j 2 22' ' Bismuthous telluride, Bi2Te"3, occurs native as telluric bismuth or tet- radymite. It forms gray, metallic, rhombohedral crystals or foliated masses. A portion of the tellurium in this mineral is generally replaced by sulphur. General Properties and Reactions oe the Compounds of Bismuth.—The salts of bismuth with colorless acids are colorless. Their solutions have an acid reaction. Dilution with water causes the solutions to become milky, owing to the separation of a basic salt. Mineral acids redissolve this basic salt; but the presence of tartaric acid does not prevent or remove the milkiness as in the case of anti- mony. Caustic alkalies and ammonia precipitate white bismuthous hydrate, insoluble in excess. Sulphuretted hydrogen gives a brown precipitate of bismuthous sulphide, insoluble in ammonic sulphide and in caustic alkalies, soluble in hot nitric acid. Potassic chromate pre- cipitates yellow bismuthous chromate, soluble in nitric acid, insoluble in caustic alkalies. (Distinction from plumbic chromate.) Heated on charcoal in the reducing flame the bismuth compounds yield a brittle metallic bead, whilst the charcoal becomes covered with a yellow incrus- tation. THE METALS. CHAPTER XXXI. DISTINGUISHING CHARACTERISTICS OF THE METALLIC ELEMENTS. There are certain broad differences which prevail between metallic and non-metallic elements, so that as a rule members of the one class may be readily distinguished from those of the other. The most ob- vious of these differences are physical. Thus the power of reflecting light is much more marked in the metals than in the non-metals. This power, when intensified by the perfect or almost perfect opacity of the reflecting substance—a property possessed in the highest degree by the metals—constitutes the phe- nomenon of metallic lustre. The non-metals are generally either trans- parent or translucent: they admit light into their interior, where it is either transmitted further, or absorbed and dispersed, and they cannot therefore possess the high reflecting power—the power of giving back the whole or nearly the whole of the light which falls upon them— necessary to the production of the metallic lustre. Smoothness of sur- face is, however, a necessary condition of metallic lustre, and for this reason finely-divided metals do not possess this property. Gold, silver, platinum, and other metals may be obtained in this condition by pre- cipitation from their solutions; but these non-lustrous powders assume a lustre under the burnisher. Again, the metals are much better conductors of heat and of elec- tricity than the non-metals. The above broad physical differences have their counterparts in the chemical characters of the elements; thus a metal uniting with oxygen generally yields a base or alkali, whilst the compounds of the non- metals with oxygen generally possess acid properties. But nature abhors classification, and renders futile all our attempts to form exclusive families of her productions. The animal and vege- table kingdoms merge into each other, so that it is impossible to predicate definitely of the intermediate members to which class they belong—whether they are to be regarded as plants or as animals. In like manner the metals and the non-metals gradually approach and overlap each other in respect of nearly all the so-called distinctive properties just enumerated. Thus, as regards lustre, we find that various non-metals possess a lustre which is distinctly metallic in character—for example, graphite, the popular name for winch, black-feud, is derived from this property. lodine is another instance : the crystals of this substance have a lustre 398 INORGANIC CHEMISTRY. resembling that of graphite, and not much inferior to that of metallic arsenic when sublimed in a glass tube. Again, as regards opacity, which was stated to be a general property of the metals, we find that this rule is not absolute. Gold in very thin leaves transmits a green light, silver a blue light, whilst, on the other hand, graphite is opaque, and iodine nearly so. Again, as regards the power of conducting heat and electricity, carbon in the form of graphite shares this power with the metals. As to chemical character also, the classification above given does not always hold. Thus some metals yield acids with oxygen—chromic acid, manganic acid, molybdic acid, and others. But no non-metal yields a decided base with oxygen. Tellurium and arsenic yield no base, and the basic properties of antimony and bismuth are very weak. Although, therefore, the division of the elements into metals and non-metals cannot lay claim to rigid accuracy, it may, in the present state of the science, be regarded as a good practical classification. With the few exceptions just enumerated, it is no more difficult to distinguish a metal from a non-metal than to distinguish an animal from a plant. Relations of the Metals to Heat. Expansion hy Heat.—Metals as a rule expand more on heating than non-metals. The following table gives the length to which the unit length of a number of substances, measured at 0° C., expands when the substance is heated to 100° C. (212° F.). This value, diminished by unity, is therefore the coefficient of linear expansion for a rise of 100° C.: Expansion of Solids by Heat. One part by length measured at 0° C. measures at 100° C.: English flint glass, .... 1.000811 French glass tube, .... 1.000861 Platinum, Palladium, . . . . 1.000844 .... 1.001000 Untempered steel, .... 1.001079 Antimony, . . . . 1.001083 Iron, . . . . 1.001182 Bismuth, .... 1.001392 Gold, .... 1.001466 Copper, . . . . 1.001718 Brass, Silver, . . . . 1.001866 Tin (East India), . . . . 1.001937 Lead, .... 1.002848 Zinc, .... 1.002942 Fusibility.—Another important property of metals is their degree of fusibility. This is almost as varied in the different metals as the range of temperature at our command. On the one hand mercury fuses at —39.5° C. (—39.1° F.), and gallium with the heat of the hand, whilst RELATIONS OF METALS TO LIGHT. 399 iridium scarcely melts in the oxyhydrogen flame, requiring the voltaic arc for its complete liquefaction. Ruthenium is still more infusible, and osmium has never been melted. The following table contains the fusing-points of some of the metals: Name of metal. Fusing-point. Mercury,.... . . . —39.5° C. (—39.1° F.). Gallium, .... . . . +30.1 “ (+86.2 “ )• Potassium, . . . . 62.5 “ ( 144.5 “ )• Sodium, .... . . . 95.6 “ ( 204.1 “ )• Lithium, .... . . . 180 “ ( 356 u )• Tin, . . . 228 “ ( 443.4 " ). Bismuth, . . . 268 “ ( 514.4 " )• Thallium, . . . 294 “ ( 561.2 “ )• Cadmium, . . . 320 “ ( 608 w )• Lead, . . . 326 “ ( 618.8 “ )• Zinc, . . . 420 , “ ( 792 a )• Antimony, . . . . 430 “ ( 810 )• Silver, .... . . . 1040 “ (1872 “ )• The fusing-point of alloys is always lower than the mean fusing-point of their constituents—taking the relative proportion of the constituents into account in calculating this mean ; and sometimes lower than the lowest fusing-point of any of the constituents. Thus Wood’s fusible metal, which is an alloy of 4 parts of bismuth, 2 of lead, 1 of tin, and 1 of cadmium, fuses at 60.5° C. (140.9° F.). The alloy of potassium and sodium is liquid at ordinary temperatures. Volatility.—All metals are volatile, but usually only at very high temperatures. Mercury boils at 360° C. (680° F.), but is volatile at ordinary temperatures, as may be shown by suspending a piece of gold- leaf from the stopper of a bottle containing mercury : in course of time, the gold-leaf becomes white, owing to the absorption of the vapor of mercury. Arsenic volatilizes below redness without first assuming the liquid form. Cadmium boils at 860° C. (1580° F.); zinc at 1040° C. (1904° F.). Potassium and sodium are distilled in their manufacture. Lead is volatilized in the process of lead smelting, and means are em- ployed to condense the lead which would thus otherwise escape. Even copper is perceptibly volatile at the temperature of the smelting furnace. Relations of Metals to Light. Colors of Solid Metals.—Most metals appear nearly colorless when polished. Some, however, exhibit, even when viewed in the ordinary way, specific colors: thus copper is red ; and gold, calcium, and barium display shades of yellow. By causing the light to be reflected several times from their surfaces, some metals, which under ordinary conditions appear colorless, may be made to exhibit color, whilst in the case of the colored metals the particular shade is intensified or altered. Thus by multiple reflection the following metals display the annexed colors: 400 INORGANIC CHEMISTRY, Copper, .... Gold, Silver, Zinc, Iron, At large angles of incidence—that is, when the light falls very ob- liquely upon the surface—all metals reflect white light. But their specific reflective power for the different rays varies more as the inci- dent light becomes more perpendicular. Colors of Ignited Liquid Metals.—At high temperatures, metals in the liquid state generally emit white light; but molten copper gives out a ruddy glow, and molten gold emits a beautiful green light. Colors of Ignited Vaporous Metals.—All metallic vapors exhibit at very high temperatures characteristic phenomena of color, and some possess, even at relatively low temperatures, colors more or less marked. Thus tin gives a blue vapor; copper a green; silver a green of a different shade; gold a blue; and sodium a yellow. The nature of the colors which metallic and other vapors display at high temperatures forms the subject matter of Spectrum Analysis. Spectrum Analysis. The study of the colors of the vaporous elements at high tempera- tures has developed, in the hands of chemists, into an invaluable method of analysis, surpassing in scope and delicacy all other known methods. This method is known as spectrum analysis, and the instru- ment by means of which the discrimination of the colors of the vapors is effected is the spectroscope. Although this method has been em- ployed by chemists only since 1860, it has already been the means of enriching chemistry with several new metals. It has further demon- strated that some elements which were formerly believed to have been obtained in a state of purity have in reality been contaminated with foreign matter: a state of things which has rendered necessary a revi- sion of some of the atomic weights. But the achievement of spectrum analysis which appeals most powerfully to the imagination is the crea- tion of an entirely new branch of chemical science, that of celestial chemistry, in which, by the application of the spectroscope to the exami- nation of the light emitted by solar and stellar matter, chemists have been enabled to prove the presence of many of our terrestrial elements in the sun and stars. In addition to this, the spectroscope has fur- nished us with information concerning the physical constitution of these luminaries, and even concerning their rate of motion, which would formerly have been deemed unattainable. The form of spectroscope most generally employed for chemical pur- poses is represented in Fig. 48. The rays of light to be examined pass through a vertical slit situated at the end of the tube A, and turned from the spectator in the diagram. After being rendered parallel by means of a lens, they fall upon the prism P. The spec- trum is viewed directly by means of the telescope B. In this way it is SPECTRUM ANALYSIS. 401 not only magnified, but is made to exhibit a greater degree of sharpness of detail than it would possess if thrown upon a screen. The tube C carries a transparent horizontal graduated scale, which is illuminated by a small luminous gas-flame placed at the end of the tube and not Fig. 48. represented in the diagram. In looking through the telescope this scale is seen reflected in the face P of the prism. In this way it is viewed simultaneously with the spectrum, the various parts of which may thus be referred to the divisions of the graduated scale. In order to compare light from two different sources, one half of the slit, which is represented on a larger scale in Fig, 49, is covered by a small prism ah. The light from one source, F, Fig. 48, situated in front of the Fig. 49. slit passes directly through the uncovered half of the slit; the light from the second source,/, which must be placed to the side of the slit, passes through the covered portion of the slit by total reflection in the small prism. Various arrangements of the small prism are employed for this purpose; one of the simplest is that represented in Fig. 50. The light from /, Fig. 50, enters the equilateral prism ode perpen- dicularly to the face de, is totally reflected at r from cd, and emerging from the prism perpendicularly to ce, enters the slit s. As the direction of the rays of light on entering and on quitting the prism is perpen- 402 INORGANIC CHEMISTRY. dicular to the faces of the prism, no refraction occurs with this prism. At the same time the light from F cannot enter the slit through the prism, and can pass only through the uncovered portion of the slit. In this way the two spectra from the two sources of light may be viewed simultaneously, one above the other, and as in both spectra the light ’ Fig. 50. passes through the same slit and is refracted by the same prism P, there will be perfect correspondence of the similar parts of each : rays of the same wave-length will be found in the same vertical line in the two spectra, and thus coincidences may be observed and studied. In order to understand the principles upon which spectrum analysis is based, it will be necessary to consider what is the precise nature of the phenomena observed when bodies are heated to the temperature at which they become self-luminous. If a liquid or a solid be thus gradually heated, and at the same time examined with the spectroscope, the red end of the spectrum will be observed first. The body is then at a low red heat. As the temperature rises, the orange rays will be added; then the yellow, and so on from the less refrangible to the more refrangible rays, until the entire visible spectrum from red to violet can be seen. The body is then white-hot, and the white light which it emits is thus seen to be compounded of every wave-length in the visible spectrum.* The spectra of glowing solids and liquids are there- fore continuous. The molecules of solids and liquids are hampered by cohesion, and are not free to take up those vibrations which are peculiar to them. We may conceive that in different parts of the mass cohesion is overcome to a varying extent at the same time, and that molecular groupings of every possible degree of variety and complexity are vi- brating, each with its specific rate of vibration. We should thus have the simultaneous emission of light of every wave-length—of every degree of refrangibility. Gases or vapors behave otherwise. Their molecules are free to os- cillate unimpeded by each other; and the molecules of any one element, being all of the same kind, execute at a given temperature vibrations * In addition to the visible spectrum, there is an invisble region of rays of lower refrangibility than the red—the infra-red rays—and a second invisible region of rays of higher refrangibility than the violet—the ultra-violet rays. These two invisible por- tions, which lie on either side of the visible spectrum, have by the aid of photography been rendered accessible to spectroscopic study. SPECTRUM ANALYSIS. 403 identical in nature and in velocity, and consequently, if heated to the temperature at which they become self-luminous, emit light of definite wave-lengths and therefore of definite color, not a mixture of light of all wave-lengths, or white light. Every element, therefore, in the state of self-luminous vapor, and at a temperature sufficiently high, displays a spectrum peculiar to itself, and consisting of definite lines or bands. The dark spaces between these lines or bands correspond to those wave- lengths of light which the atoms or molecules of the element do not excite. In this way the spectroscopic examination of the elements in the con- dition of self-luminous vapor affords a means of distinguishing between them—a means more expeditious, less liable to misinterpretation, and, as we shall see presently, more delicate than the ordinary chemical tests. The identification of the elements by means of the spectroscope is greatly facilitated by the arrangement for comparing spectra already described (p. 401). For example, if the spectrum of a substance under examination appears to be that of barium, it is only necessary to ex- amine, simultaneously with this spectrum, the spectrum of an actual specimen of barium by means of the comparing prism : the coincidence or non-coincidence of the lines in the two spectra will at once inform us whether our surmise is correct, or the reverse. The metallic vapors for examination may in many cases be obtained by heating the metal or its compounds in the Bunsen flame. Some- times, however, a higher temperature is necessary, in which case the electric arc or the induction spark may be employed as the source of heat. In the case of metals, it is sufficient to pass the spark between poles of the metal, when a sufficient quantity is volatilized to give a spectrum. It is to be borne in mind, however, that in this method the spectra of the gases through which the spark passes (oxygen, nitrogen, etc.) will also be visible. As regards the certainty of identification of the elements by spectro- scopic means, a noteworthy point is the ease with which metallic vapors, the colors of which appear to the eye almost or entirely identical, may be discriminated with the aid of the spectroscope. The red colors which lithium and strontium compounds respectively impart to the Bunsen flame, though distinguishable by a trained eye, are yet extremely sim- ilar; but the flame spectrum of lithium consists of a bright red line and a very weak line in the yellow, whilst that of strontium contains several lines in the red, one in the orange, and one in the blue. The flame colors of the compounds of potassium, cmsiura, and rubidium are to the eye absolutely identical, and there are, moreover, no character- istic qualitative tests by which the compounds of these elements may be distinguished, but their spectra present the most marked differences. So similar are these elements, that it is probable that by chemical means alone csesiumand rubidium could never have been discovered. Indeed csesium had, previous to its spectroscopic recognition as a distinct ele- ment, been confounded with potassium (see Csesiura). The delicacy of the spectroscopic tests for the elements is due to the minuteness of the quantity of self-luminous vapor necessary to impart to the luminiferous ether a perceptible impulse. The highest degree 404 INORGANIC CHEMISTRY. of delicacy is manifested in the case of sodium, a quantity of which less than the ttwoijWouo a gram may be detected. This almost inconceivable delicacy is due to two causes : in the first place the spec- trum of sodium consists of one double line in the yellow, hence the entire effort of the atoms is concentrated upon one part, and that the most luminous of the spectrum; and, secondly, the atomic weight of sodium is low, so that a smaller quantity is required to produce an effect. Thallium also gives only one line, but it is in the green —a portion of the spectrum which affects the eye less powerfully; and the atomic weight of thallium is high; hence the reaction is in this case less delicate. In the case of lithium soiydtto'ij °f a gram may be detected. With the induction spark 7 5 answer® a gram of copper gives a brilliant spectrum, and 0.2 of a milligram of copper keeps up this spectrum for six hours. In identifying an element by means of its spectrum, it is not neces- sary that every line in the spectrum should be perceived. In almost all spectra there are certain lines brighter than the rest, and these are frequently visible when the quantity of substance vaporized is insuffi- cient for the perceptible production of the fainter lines. The presence of one of these prominent or characteristic lines is sufficient for the identification of an element. Nearly all metallic compounds are decomposed into their elements at a temperature below that at which their vapors become luminous. On this account the spectra of the compounds of the metals with the non- metals are frequently the same as those of the metals themselves.* But this is not always the case, especially at comparatively low tempera- tures. Thus copper and cuprous chloride give the same spectrum in the electric arc, but not in the Bunsen flame. In many such cases there is a temperature at which a compound gives its own peculiar spectrum plus that of each of its elements. When no chemical combination occurs, spectra of any number of elements can co-exist side by side without confusion. In this way the qualitative analysis of mixed materials may be safely made. It is only necessary to identify in the mixed spectrum the more characteristic lines of the various elements. Gases which under ordinary pressures give a line spectrum behave otherwise under high pressures. As the pressure increases, the lines gradually broaden, until ultimately the spectrum becomes con- tinuous. This is again a case in which the freedom of atomic vibration is interfered with by the too great proximity of the atoms to each other. All bodies capable of vibration possess the power of taking up or absorbing those waves which they would cause by their own vibration. Thus a finger-glass may be made to sound by singing its own note close ta it. The same law holds with regard to the vibrating atoms * The non-metals require a higher temperature than many of the metals in order that they may exhibit their characteristic spectra. Thus in the case of the decomposed compound of a metal with a non-metal, it frequently happens, as above stated, that the spectrum of the metal alone is visible. SPECTRUM ANALYSIS. 405 and molecules of a gas. If we examine with a spectroscope a source of while light yielding a continuous spectrum—a white-hot solid or liquid —and then introduce between the slit of the spectroscope and the source of white light, a layer of sodium vapor, then according to the relative temperatures of the source of white light and the sodium vapor, one of three things will happen: either the sodium vapor is hotter—i.e., pos- sesses greater energy of atomic vibration—than the white-hot solid or liquid, in which case it will emit more yellow light than it receives from the latter, and a bright yellow sodium line will be visible in the otherwise continuous spectrum; or it is of the same temperature, when it will emit just as much as it receives, and only the continuous spec- trum will be seen ; or, finally, it is colder, in which case it will absorb more than it emits, and a dark sodium line will be visible on the back- ground of the continuous spectrum. This is in accordance with the law of exchanges. Its chief importance in connection with the present subject lies in the explanation which it affords of the phenomena ob- served in the spectroscopic study of the heavenly bodies. Solar and Stellar Spectra.—lf the light from the sun be examined spectroscopically, the phenomena observed do not correspond either with those of an incandescent gas, or with those of an incandescent solid or liquid. The visible solar spectrum consists of a band of colored light stretching from the red to the violet; but this colored spectrum is crossed by a vast number of fine dark lines. These lines were first observed by Wollaston. They were afterwards mapped by Fraunhofer, a German optician, for which reason they are known as the Fraunhofer lines. If we examine simultaneously by means of the comparing prism the solar spectrum and the spectrum of a metallic element, we find that in the case of many metallic elements, such, for example, as iron or calcium, every bright line in the spectrum of the metallic element corresponds in position, breadth, and intensity with a dark line in the solar spectrum. We have already seen that the bright line of sodium maybe reversed and converted into a dark line. The dark lines in the solar spectrum have a similar origin. In the sun, we have in the first place an incan- descent nucleus, solid or liquid, the source of light, and capable of yield- ing a continuous spectrum. Owing to the high temperature of the sun, the elements, of which the mass of this luminary is composed, are in part volatilized, and we have thus an atmosphere of incandescent vapor surrounding the incandescent nucleus. Through this atmosphere all light from the nucleus must pass. The temperature of the solar atmos- phere is necessarily lower than that of the nucleus; hence metallic vapors contained in this atmosphere absorb more light than they emit, and the lines of their spectra consequently appear dark on the continu- ous spectrum of the nucleus. The nucleus of the sun is distinguished as the 'photosphere ; its atmosphere, in which this selective absorption occurs, as the chromosphere. Under certain conditions it is possible to submit the light from the chromosphere alone to spectroscopic exami- nation, and in this case a spectrum of bright lines on a dark ground, corresponding with that of a glowing gas, is obtained. The origin of the dark lines in the solar spectrum was first satisfac- 406 INORGANIC CHEMISTRY. torily explained by Kirchhoff, who verified his theory by an elaborate series of observations. The same explanation had, however, been pre- viously suggested by Stokes. The alternative hypothesis, that the coincidence of the bright lines of the spectra of the metallic elements with the dark lines of the solar spectrum is due to chance, and not to the presence of these elements in the solar atmosphere, is untenable. In the case of the spark spectrum of iron, Angstrom has counted no fewer than 460 lines, each of which coincides with a dark line in the solar spectrum. The probability of 460 chance coincidences in the spectrum of one metal is inconceiv- ably small; and, when we take into account the fact already men- tioned, that the coincidence of the lines is one not merely of position, but in every case one also of breadth and intensity, this small proba- bility becomes still further diminished. We must therefore conclude that the various elements which yield these lines are really present in the solar atmosphere. The following is a list of the metallic elements which have thus been detected in the atmosphere of the sun : H, Ka, K, Rb, Cs, Li, Ba, Sr, Ca, Mg, Al, Cr, Be, Ce, La, Yt, Zn, Mn, Ni, Co, Fe, U, V, Pb, Bi, Cu, Cd, Pd, Ir, Sn, Mo, Ti. The spectroscopic study of the stars has afforded much information concerning the constitution of these bodies. The moon and planets exhibit the same spectrum as the sun, which is in accordance with the fact that they shine by the reflected light of that luminary. The fixed stars are found to be bodies constituted like our sun, although differing greatly both from the latter and from each other. The spectra of the greater number display dark lines. Many terrestrial elements have already been detected in the stars. Thus Aldebaran contains hydrogen, sodium, magnesium, calcium, iron, tellurium, antimony, bismuth, and mercury; whilst in Sirius sodium, magnesium, and hydrogen have been detected. The spectra of the irresolvable nebulae, on the other hand, display bright lines. This shows that these nebulae consist of masses of incan- descent gas, without a solid or liquid nucleus—a discovery which affords powerful support to the Kant-Laplace hypothesis of the origin of the solar system. Relations of the Metals to Gravity. Specific Gravity of Metals.—A table of specific gravities of substances consists of a series of numbers indicating the relative quantities of matter contained in equal volumes of these substances. The measure of the quantity of matter is, cceteris paribus, the weight. Since in the case of solids and liquids the specific gravity of water at 4° C. is taken as unity, we may put it that the number expressing the specific gravity of a solid or liquid substance indicates the number of times that a given volume of this substance is heavier (or lighter) than an equal volume of water at 4° C. For an account of the methods by which the specific gravity is determined a work on physics must be consulted; but the following relation, which is useful to remember, may be mentioned COHESIVE POWER. 407 here: The number expressing the specific gravity of a solid or liquid also expresses the weight in grams of one cubic centimetre of the substance measured at the temperature at which the specific gravity was deter- mined. This is due, in the first place, to the fact that, in the metric system, the unit of weight is the weight of the unit of volume of water at 4° C. (1 cubic centimetre of water at 4° C. weighs 1 gram); and, secondly, to the fact above mentioned, that the specific gravities of solids and liquids are referred to that of water at 4° C. as unity. The metals exhibit a very wide range in their specific gravities, vary- ing from 0.594 in the case of lithium, the lightest of known solids, to 22.477 in the case of osmium, the heaviest. The following table contains the specific gravities of some of the more important metals: Name of Metal. Sp. gr. Osmium, . . . 22.477 Iridium, . . . 22.40 Platinum, . . . . 21.50 Gold, . . . 19.26 Mercury, . . . 13.596 Lead, . . . 11.37 Silver, . . . 10.47 Copper, Cadmium, . . . . 8.95 . . . 8.66 Iron, . . . 7.79 Tin, . . . 7.29 Zinc, . . . 6.92 Aluminium, . . . . 2.67 Magnesium, . . . . . 1.74 Sodium, . . . 0.974 Potassium, . . . . . . . 0.865 Lithium, . . . 0.594 Cohesive Power. The properties of matter which are dependent upon cohesion, that is to say, upon the mutual attraction of the molecules of a substance, are tenacity, hardness, brittleness, malleability, and ductility. These very important properties are possessed by the various metals in very differ- ent degrees. Upon them depends the value or otherwise of the metals for the purposes of art and manufacture. The tenacity of a substance is the resistance which that substance opposes to the separation of its parts. This separation may be sought to be effected either by strain or by crushing weight. The tenacity of a metal towards strain may be determined by suspending weights by a wire of the metal, and noting the weight sufficient to cause rupture. By repeating this operation with wires of different metals, care being taken that the wires are, in every case, of equal cross-section, a table of relative tenacities may constructed. In the following table the tenacity of lead is taken as unity: 408 INORGANIC CHEMISTRY. Relative Tenacity of Metals. Lead, ... 1 Tin, ... 1.3 Zinc, ... 2 Palladium, . . . 11.5 Gold, . . . 12 Silver, Platinum, .... . . . 12.5 ... 15 Copper, .... ... IS Iron, . . . 27.5 Nickel, . . . 41.2 Steel, ... 42 This means that if a lead wire of given thickness willl support, as maximum load, say 1 kilogram, a steel wire of the same thickness will support 42 kilograms. The tenacity of cobalt is greater than that of iron. The tenacity of most metals is diminished by annealing; that is, by heating the metal and allowing it to cool slowly. Resistance to strain and to crushing weight are distinct properties. Thus the three kinds of iron range as follows in regard to their order of tenacity: Strain. Crushing weight. Wrought iron. White cast iron. Gray cast iron. Gray cast iron. White cast iron. Wrought iron. Hardness is the resistance which a substance opposes to penetration, or to change of form generally. It is not easy to determine hardness with quantitative accuracy; but we may readily ascertain which of two substances is the harder by endeavoring to scratch the one with the other. In this way a scale of standard substances has been prepared, each of which is harder than its predecessor: 1. Talc. 6. Felspar. 2. Gypsum or rock salt. 7. Quartz. 3. Calcite. 8. Topaz. 4. Fluorspar. 9. Corundum 5. Apatite. 10. Diamond. Scale of Hardness. (Mohs.) Thus, a substance which scratches fluorspar but is scratched by apatite, has a hardness lying between 4 and 5. The numerals denote simply order, not degree of hardness. This scale is much employed by min- eralogists. Among the metals, titanium, manganese, chromium, and ruthenium are so hard as to scratch glass, whilst sodium may be moulded with the fingers. The native alloy of osmium and iridium is exceedingly hard, and is employed on this account in the manufacture of the nibs of gold pens. Brittleness is the incapacity of a substance to undergo change of form —by bending, hammering, or otherwise—without rupture. Among the MALLEABILITY AND DUCTILITY. 409 metals, brittleness is generally associated with a crystalline structure; the crystalline metals, antimony, arsenic, and bismuth, fly into fragments under the hammer. Tenacious metals frequently possess a fibrous struc- ture. Thus, the highly tenacious metals, wrought iron and wrought copper, are fibrous, as may be seen by fracturing a bar of the metal by repeated bending and observing the surface of fracture; whereas, cast iron and slowly deposited electrolytic copper are crystalline and brittle. Fibrous wrought iron, when kept in a state of vibration for a great length of time, undergoes a slow molecular rearrangement whereby the fibrous structure becomes crystalline. To this cause is sometimes due the snapping of the axles of railway carriages and of the shafts of screw steamers. Malleability and Ductility.—Malleability is the property of being reducible to thin leaves, either by hammering or by passing between rollers. The most malleable of the metals is gold; it has been beaten into leaves of an inch in thickness. 1 square decimetre of this leaf weighs less than 20 milligrams. Silver and copper may also be hammered into thin leaf. The remaining metals in the accompany- ing table may be reduced to thin foil by rolling, but not by hammering: Order of Malleability. 1. Gold. 2. Silver. 3. Copper. 4. Tin. 5. Platinum. 6. Lead. 7. Zinc. 8. Iron. Ductility is the capability of being drawn into wire. The metal is first formed into rods; these are then drawn through holes in a steel draw-plate. The holes, through which the wire passes, diminish in size by regular gradation. The process of drawing is continued until the requisite degree of tenuity is attained. Sometimes it is necessary to anneal the wire from time to time during the process of drawing. Very fine gold and silver wire is drawn through an aperture in a ruby. Most malleable metals are ductile, but in an order somewhat different from that of their malleability : Order of Ductility. 1. Gold. 2. Silver. 3. Platinum. 4. Iron. 5. Copper. 6. Palladium. 7. Aluminium. 8. Zinc. 9. Tin. 10. Lead. 410 INORGANIC CHEMISTRY. Thus iron, by virtue of its superior tenacity, is more ductile than some of the more malleable metals. A non-malleable metal cannot be duc- tile. Gold wire has been drawn of an inch in diameter. Wires of gold and platinum have been obtained by Wollaston of an inch in diameter. This extraordinary degree of tenuity was attained by placing a wire of gold or platinum in the axis of a cylinder of silver, then drawing the compound wire in the ordinary way and dissolving off the silver with nitric acid. Soft metals, such as sodium and potas- sium, may be obtained in the form of wire by forcing them through an aperture in a steel die. This has of course nothing to do with the true ductility of these metals: the wires are pressed, not drawn. True ductility, as above stated, is dependent to a great extent upon te- nacity. The properties of malleability and ductility vary in each metal with the temperature. Copper is tough and malleable at ordinary tempera- tures; but at a temperature approaching its fusing-point it becomes so brittle that it maybe reduced to powder. In reference to this property copper is said to be “ hot short.” The behavior of zinc in this respect is peculiar: at ordinary temperatures it is moderately brittle: between 100° and 150° C. (212°-262° F.) it is so malleable and ductile that it may be wrought with facility: whilst at 205° C. (401° F.) it is more brittle than at ordinary temperatures, and may be pulverized in a mortar. Many metals, when fused along with others, unite with these to form a homogeneous metallic mass known as an alloy. In some such cases chemical combination appears to take place: thus the union of sodium with mercury is accompanied with evolution of heat and light; in others the combination is merely one of mutual solution. The chem- ical compounds which are formed are difficult to isolate, as they are generally soluble in all proportions in an excess of any of the constit- uents. The best characterized chemical compounds are always those which result from the union of elements differing most widely in their properties—thus of the most positive with the most negative elements; and in such compounds the properties of the constituent elements are obliterated. The metals, on the other hand, standing, as they do, near to each other in the electrochemical scale, form compounds which are devoid of sharply-defined characteristics, and in which the properties of the constituent metals are preserved. Thus all alloys possess metallic lustre, and are good conductors of heat and electricity. Very few pure metals possess properties which fit them, as such, for use in the arts. Thus pure copper is soft, and cannot be worked on the lathe. By alloying it with zinc it is converted into the hard and work- able brass. In like manner before gold and silver can be coined, these metals must be alloyed with a certain percentage of copper in order to impart to them the necessary hardness and durability. Thus the prop- erties—sometimes even the defects—of one metal are employed to cor- rect or modify those of another in the preparation of alloys. Alloys of metals with mercury are known as amalgams (q.v.). Alloys. POTASSIUM. 411 The properties of the various alloys will be treated of later on in connection with one or other of their constituent metals. The law regulating the fusing-point of alloys has already been referred to (p. 399). CHAPTER XXXII. MONAD ELEMENTS. Section 111. POTASSIUM, K2? Atomic weight = 39. Probable molecular weight = 78. Sp. gr. 0.865. Fuses at 62.5° C. (144.5° F.). Boils at a low red heat. Atomicity '. Evidence of atom icity: Potassic chloride, KC1. Potassic iodide, KI. Potassic hydrate, KHo. Potassic sulphide, SK, History.—Potassium was first isolated in 1807 by Davy, who obtained it by the electrolysis of potassic hydrate. Occurrence.—The salts of potassium are widely distributed in na- ture. Double silicates of potassium with aluminium and other metals form a variety of important minerals, which are among the proximate constituents of the igneous rocks. By the disintegration of these rocks soils are produced. From the soils the potassium is absorbed by plants, in the juices of which it occurs as the potassium salts of organic acids. From plants it passes into the bodies of animals. Further, in the inorganic world, the chloride, bromide, and iodide of potassium are found in sea-water, in mineral springs, and in solid saline deposits, whilst the nitrate occurs in tropical climates as an efflorescence on the soil. Preparation.—l. When a piece of solid potassic hydrate, slighly moistened in order to increase its conducting power, is placed between the poles of a powerful voltaic battery, decomposition takes place ac- cording to the following equation : 20KH = K2 + H2 + 02. Potassic hydrate. Potassium and hydrogen are liberated at the negative pole. The po- tassium forms metallic globules which inflame in contact with air, and must be removed and preserved under petroleum. This was the method of preparation originally employed by Davy. 412 INORGANIC CHEMI&TRY. 2. Potassium may also be obtained by the action of metallic iron on potassic hydrate at a strong white heat: 40KH + 3Fe = iT(Fe3)viiio4 + 2K2 + 2H2. Potassic Magnetic iron hydrate. oxide. 3. The most convenient method of preparing potassium consists in heating potassic carbonate to a white heat with charcoal. Hydric po- tassic tartrate (cream of tartar) is first ignited in a closed crucible, when the following decomposition takes place : f COKo 2-{ = COKo2 -(- 50H2 -(- 400 + 3C. [ COHo Hydric potassic Potassic Water. Carbonic tartrate. carbonate. oxide. The residue, consisting of potassic carbonate and finely divided carbon, is mixed with charcoal and distilled at a white heat from an iron retort E (Fig. 51): OOKo2 + 2C = 300 + K2 Potassic Carbonic carbonate. oxide. The vapor of potassium condenses in a copper receiver r, from which air is excluded. If the neck of the retort becomes choked during the pro- Fig. 51. cess, it may generally be cleared by means of an iron rod h, introduced through the lateral tube of the receiver. Should this fail, the fire-bars COMPOUND OF POTASSIUM WITH HYDROGEN. 413 B, which are movable, must be withdrawn, so as to allow the fire to fall on to the hearth. The potassium obtained by the above process is contaminated with carbonic oxide, from which it must be freed by redistillation. A neglect of this precaution may lead to dangerous accidents, as when the crude potassium is preserved, even under petroleum, a black pow- der is formed which explodes violently on the slightest friction. Properties.—Potassium is a silvery-white metal, brittle and crystal- line at 0° C, but at ordinary temperatures soft like wax. The freshly cut surface of the metal has a brilliant lustre, which it almost instantly loses when exposed to air, owing to the formation of oxide. For this reason it is necessary to keep the metal immersed in some liquid devoid of oxygen, such as petroleum. When heated in air it inflames and burns with a violet light, forming a mixture of peroxides of potassium. By melting potassium in a sealed tube filled with coal-gas, allowing the metal partially to solidify, and then pouring off the liquid portion, well formed crystals of potassium may be obtained. Reactions.—l. Potassium decomposes water, even at its freezing point, with great energy, the heat evolved being sufficient to cause the ignition of the liberated hydrogen : K2 + 20H2 = 2KHo + H2. Water. Potassic hydrate. 2. It inflames spontaneously in an atmosphere of chlorine. It also inflames when brought in contact with bromine, the reaction taking place with explosive violence. In these cases potassic chloride (KCI) and bromide (KBr) are formed. 3. When potassium is ignited in a stream of carbonic anhydride, a portion of the latter is reduced, with liberation of carbon: 2K2 + 3C02 = 20OKo2 + C. Carbonic Potassic anhydride. carbonate. Uses.—Owing to its powerful affinity for electro-negative elements, potassium is employed in the laboratory to expel elements, less strongly electro-positive than itself, from their combinations with electro-negative elements. Thus, by its means, boron and silicon may be prepared from their oxides, and aluminium, magnesium, and other metals from their chlorides. The more readily obtainable sodium has, however, almost totally superseded it for these purposes. COMPOUND OF POTASSIUM WITH HYDROGEN. Potassic hydride, K4H2.—When potassium is heated in a current of pure hydrogen, the gas is absorbed by the metal, and potassic hydride is formed. The absorption begins at 200° C. (392° F.), and is most rapid about 300° C. (572° F.). The hydride is a brittle crystalline mass, with a silvery metallic lustre. It may be fused in an atmos- phere of hydrogen. Under ordinary pressures it may be heated to 410° C. (770° F.) without change, but in a vacuum it begins to dissociate at 200° C. (392° F.). It in- flames spontaneously in conact with air. 414 INORGANIC CHEMISTRY. COMPOUNDS OF POTASSIUM WITH THE HALOGENS. Potassic chloride, KCI, occurs native in saline deposits as the mineral sylvine. In smaller quantities it is found in sea-water and in brine-springs. It crystallizes in colorless cubes, and possesses a saline taste. It dissolves in 3 parts of water at ordinary temperatures, and is more soluble at higher temperatures. Alcohol does not dissolve it. It forms molecular compounds—double salts—with various other me- tallic chlorides. Potassic platinic chloride (potassic chloroplatinate), PtCl4,2KC1, is obtained as a granular precipitate, consisting of minute octahedra, when solutions of the two chlorides are mixed. This salt is almost insoluble in cold water, and is used in the quantitative determi- nation of potassium. Potassic bromide, KBr, forms colorless cubes of sp. gr. 2.69, readily soluble in water, slightly soluble in alcohol. Potassic iodide, KI, is prepared by digesting iron filings, water, and iodine together, filtering the colorless solution, and precipitating the iron by potassic carbonate : Fe + I 2 Fel2; Ferrous iodide. Fel2 + COKo2 = 2KI + COFeo". Ferrous Potassic Potassic Ferrous iodide. carbonate. iodide. carbonate. It crystallizes in cubes of sp. gr. 2.9. It dissolves at ordinary tem- peratures in 0.7 part of water and in 40 parts of alcohol. The aqueous solution dissolves large quantities of iodine. Potassic iodide forms molecular compounds with many other metallic iodides. Potassic fluoride, KF, is obtained by neutralizing hydrofluoric acid with potassic carbonate. At ordinary temperatures it is deposited from its solutions in crystals of the formula KF,2OH2, but above 35° C. (95° F.) it crystallizes in anhydrous cubes. It is deliquescent. The solution attacks glass. It forms numerous double fluorides: the so-called acid fluoride has the formula KF,HF. Potassic silicqfluoride, SiF6K2 ( = SiF4,2KF), which is formed as a gelatinous precipitate when hydrofluosilicic acid is added to the solution of a potash salt, may also be regarded as belonging to this class. COMPOUNDS OF POTASSIUM WITH OXYGEN. Potassic oxide, ok2. k—o—k. Potassic dioxide, <; [ok- k-o-o-k. ) r ok Potassic tetroxide,...... -t M 1 o 1 Q 1 o >2, oc 1 , OK COMPOUND OF POTASSIUM WITH HYDROXYL. 415 Potassic oxide, 0K2, is formed by the spontaneous oxidation of po- tassium at ordinary temperatures in dry air. It may also be obtained by heating potassic hydrate with potassium: 2KHo + K2 = 20K2 + H2; Potassic Potassic hydrate. oxide. or by fusing together, in a current of nitrogen, potassic peroxide and potassium. Potassic oxide is white, fusible, and, at high temperatures, volatile. It is very deliquescent, and combines violently with water to form potassic hydrate. When moistened with water it becomes incandescent. Potassic dioxide, K202, is formed with evolution of oxygen when the tetroxide is dissolved in water. Potassic tetroxide, Potassic peroxide, K204, is prepared by fusing potassium in a current of oxygen. It is a chrome-yellow powder. Water decomposes it as above (see Potassic dioxide). COMPOUND OF POTASSIUM WITH HYDROXYL. Potassic hydrate, Caustic potash, Potash, KHo or OKH, is pre- pa red by boiling in an iron vessel a solution of potassic carbonate with calcic hydrate: COKo2 + CaHo2 = 2KHo + COCao". Potassic Calcic Potassic Calcic carbonate. hydrate. hydrate. carbonate. 1 part of potassic carbonate is dissolved in 12 parts of water, and milk of lime is added till a sample of the filtered liquid no longer effer- vesces when treated with an acid. (With a concentrated solution of the carbonate, the reaction does not take place; in fact a concentrated solu- tion of potassic hydrate decomposes calcic carbonate with formation of potassic carbonate and calcic hydrate.) The clear solution of potassic hydrate is decanted from the precipitate of calcic carbonate, and is concentrated, first in a covered iron pot, and afterwards in a silver basin, until all the water has been driven off and the fused oily hydrate remains. This solidifies on cooling to a crystalline mass. It is also formed by the action of potassium upon water (see p. 413) and by dissolving potassic oxide in water: OK2 + OH2 = 20KH. ‘ Properties.—Potassic hydrate is a hard white brittle substance, with a slightly fibrous fracture. It fuses below a red heat, and at higher temperatures volatilizes without decomposition. It is very deliquescent. It dissolves in about half its weight of water, yielding a highly caustic solution, which, when exposed to the air, rapidly absorbs carbonic anhy- 416 INORGANIC CHEMISTRY. dride. Hot concentrated solutions deposit on cooling quadratic plates, or octahedra, of the formula KH0,20H2, readily soluble in alcohol. Reactions.—By contact with acids, potassic hydrate produces potas- sium salts: KHo HCI = KCI + OH2. . Potassic Hydrochloric Potassic Water, hydrate. acid. chloride. KHo + S02Ho2 == SQ2HoKo + 0H2. Potassic Sulphuric Hydric potassic Water, hydrate. acid. sulphate. 2KHo + S02Ho2 = S02Ko2 + 0H2. Potassic Sulphuric Potassic Water, hydrate. acid. sulphate. OXY-SALTS OF POTASSIUM. Potassic nitrate, Nitre, Saltpetre, N02Ko. (Occurrence,formation, nitre plantations, see p. 214.) Nitre is manufactured in large quantities from Chili saltpetre (sodic nitrate) by the double decomposition of the latter salt with potassic chloride. Equal molecular proportions of the two salts are dissolved in hot water until the specific gravity of the solution attains to 1.5. Sodic chloride, which is almost equally soluble in hot and in cold water, separates out, whilst the solution deposits potassic nitrate on cooling. The product is technically known as “ con- verted nitre.” Potassic nitrate is dimorphous. It crystallizes most frequently in longitudinally striated six-sided prisms belonging to the rhombic system, but may also be obtained in minute rhombohedra, isomorphous with those of sodic nitrate. It has a cooling saline taste. It dissolves in four times its weight of cold water, and in a third of its weight of boiling water, but is insoluble in alcohol. It fuses at 339° C., and at a red heat is decomposed with evolution of oxygen and forma- tion of potassic nitrite. At a very high temperature it is converted into potassic oxide. Owing to its property of thus parting with oxygen, it oxidizes most of the elements when heated with them, frequently with explosive violence. Gunpowder.—Gunpowder is a mixture of 75 parts of nitre, 10 parts of sulphur, and 15 parts of charcoal. The composition varies, however, in different countries, and also according to the purpose for which the powder is intended. The separate ingre- dients are finely powdered, then intimately mixed, adding a small quantity of water; the mixture is pressed by hydraulic power into a hard cake, which is then granulated. The grains are sorted according to size, polished, and finally dried. The principal pro- ducts of the combustion of gunpowder are nitrogen, carbonic anhydride (with traces of carbonic oxide), potassic sulphate, and potassic carbonate. The explosive force of gun- powder is due to the sudden evolution of gases occupying a volume several hundred times greater than that of the original substance. Potassic nitrite, NOKo, is prepared by fusing the nitrate, either alone or with lead, the oxidizable metal serving to remove the oxvgen from the nitrate. The mass is ex- tracted with water, and the solution evaporated and allowed to crystallize. The un- changed nitrate separates out first, whilst the nitrite remains in the mother liquor, from which it may be obtained by further evaporation in small prismatic deliquescent crystals. It is insoluble in absolute alcohol. COMPOUNDS OF POTASSIUM. 417 Potassic chlorate, (Preparation, p. 182.) This salt forms lustrous tabular crystals belonging to the monoclinic system, soluble in 16 parts of cold, and in 2 parts of boiling water. It fuses at 334° C. (633° F.), and is decomposed at 352° C. (666° F.) into oxygen, potassic chloride, and potassic perchlorate. At a still higher temperature it parts with the whole of its oxygen, and is converted into potassic chloride (pp. 184 and 161). It is a powerful oxidizing agent, and, along with sulphur or anti- monious sulphide, forms detonating mixtures which explode by percus- sion or friction, owing to the sudden combustion of the oxidizable ingredient at the expense of the oxygen of the potassic chlorate. roci Potassic perchlorate, < O (Preparation, p. 184), crystal- [ OKo lizes in rhombic prisms, soluble in 70 parts of cold, in 6 parts of boiling water, insoluble in alcohol. When heated to about 400° C. (752° F.) it is decomposed into oxygen and potassic chloride. f OBr Potassic bromate, qj-0) is best prepared by passing chlorine into an aqueous solution of 1 mol. of bromine with 6 molecules of potassic carbonate: 6COKo2 + 5C12 + Br2 = 2{°|rQ + 10KC1 + 6C02. Potassic Potassie Potassic Carbonic carbonate. bromate. chloride. anhydride. (See also p. 319.) It crystallizes in rhombohedra, sparingly soluble in water. It resembles in its properties the chlorate. Potassic iodate, | qj^()-—Chlorine is passed into water, in which iodine is sus- pended, until all the iodine dissolves. Potassic chlorate is then added, when potassic iodate is formed with evolution of chlorine; + {gg. - {oko + Hypiodous Potassic Potassic chloride. chlorate. iodate. (See also p. 303.) It forms small, lustrous, regular crystals, soluble in 13 parts of cold water. It decomposes on heating into oxygen and potassic iodide. {Hyperacid iodates, p. 303.) fOl Potassic periodate, ( O ,is prepared like the sodium salt (q.v.). It forms small ( OKo rhombic crystals which require 300 times their weight of cold water for solution. Between 250° and 300° C. (482-572° F.) it undergoes decomposition into oxygen and potassic iodate; at a higher temperature it parts with all its oxygen, and is converted into potassic iodide. (For the formulae of the more complex periodates, see p. 306.) Potassic carbonate, COKo2, is obtained from the ashes of land plants. Wood ashes, when lixiviated, yield a solution of potassic car- bonate, contaminated with small quantities of sodic carbonate, potassic and sodic chlorides, and potassic sulphate. When the solution is evap- orated, the impurities crystallize out first, leaving the more soluble potassic carbonate in the mother liquor, from which it may be obtained in the crystallized form by further evaporation. On a large scale it is 418 INORGANIC CHEMISTRY. prepared from native potassic chloride by a process similar to that by which sodic carbonate is obtained from sodic chloride (see Leblanc’s process). Very pure potassic carbonate may be obtained by igniting hydric potassic tartrate (cream of tartar) and extracting with water the mixture of potassic carbonate and carbon (see p. 412). It crystallizes from its aqueous solution in colorless, long, pointed monoclinic prisms of the formula 2COKo2,30H2. This salt, when dried at 100° C., has the formula COKo2,0H; at a higher temperature it becomes anhydrous. The anhydrous salt is fusible, and, at a bright red heat, volatile. It is deliquescent and very soluble in water, but insoluble in alcohol. The solution has a strong alkaline reaction.—Hydric potassic carbonate, COHoKo, is formed when carbonic anhydride is passed into a concen- trated solution of the normal carbonate : coko2 + co2 + on2 = 2COHOKO. Potassic Carbonic Water. Hydric potassic carbonate. anhydride. carbonate. It crystallizes in anhydrous monoclinic prisms, which are soluble in 3-4 parts of cold water. When the dry salt is heated, or when its aqueous solution is warmed to 80° C. (176° F.), it is decomposed into normal carbonate, carbonic anhydride, and water. Potassic sulphate, S02Ko2, is obtained in large quantities as a by- product in many manufacturing processes. It forms anhydrous, color- less, rhombic crystals, with a bitter, saline taste, which are soluble in 10 parts of cold, in 4 parts of boiling water. It decrepitates on heating, and fuses at a bright red heat.—Hydric potassic sulphate, S02HoKo, is obtained as a by-product in the preparation of nitric acid (p. 215), and may be prepared by heating 1 molecule of the normal salt with 1 mole- cule of sulphuric acid. From solutions containing an excess of acid, it crystallizes in tabular rhombic crystals. It fuses about 200° C. (392° F.), and may be obtained in monoclinic crystals by the slow solidification of the fused salt. It is readily soluble in water, but an excess of this sol- vent decomposes it into the normal salt and free sulphuric acid. For this reason, only the normal salt is deposited from dilute solutions. Heated above its fusing-point, it parts with the elements of water and r S02Ko is converted into potassicpyrosulphate, < O , which, at a temperature (SO2Ko of 600° C. (1112° F.), breaks up into normal sulphate and sulphuric anhydride (cf. p. 266). Potassic sulphite, SOKo2,20H2, is prepared bypassing sulphurous anhydride into a solution of potassic carbonate until the carbonic anhy- dride is expelled. It forms monoclinic octahedra, which are very soluble in water and somewhat deliquescent. The solution possesses an alkaline reaction and a bitter taste. When heated, the salt is decomposed, yielding potassic sulphate, potassic sulphide, and potassic hydrate.— Hydric potassic sulphite, SOHOKO, is obtained by saturating a concen- trated solution of potassic carbonate with sulphurous anhydride. It forms very soluble monoclinic prisms. The addition of alcohol to the COMPOUNDS OF POTASSIUM. 419 aqueous solution causes the salt to be deposited as a mass of acicular crystals. It has an alkaline reaction and emits an odor of sulphurous anhydride. Exposed to the air in solution, it is gradually oxidized to sulphate. fSOKo Potassic pyrosulphite, \ O ,is formed when sulphurous anhydride is passed into iSOKo a warm concentrated solution of potassic carbonate until effervescence ceases and the liquid assumes a greenish tinge. On cooling, it is deposited in granular crystals. \ SO Ko Potassic dithionate, | q(/{£q> prepared hy exactly precipitating the barium salt (q.v.) with potassic sulphate. It forms hexagonal crystals, soluble in 16 parts of cold, in 14 parts of boiling water. On heating, it is decomposed into potassic sulphate and sulphurous anhydride. Potassic thiosulphate, 2S02K0K5,30H2.—This is prepared like the sodium salt (q.v.). The salt of the above formula is deposited from its aqueous solution at ordi- nary temperatures, and crystallizes in rhombic octahedra. At temperatures above 30° C. (86° F.), the solution deposits thin four-sided prisms of the formula 3S02KoKs,OH2- At 200° C. (392° F.) the water of crystallization is expelled, and at a still higher temperature the salt is decomposed into a mixture of potassic sulphate and penta- sulphide’; 4S02KoKs = 3S02Ko2 + K2S5. Potassic Potassic Potassic thiosulphate. sulphate. pentasulphide. Potassic selenate, Se02Ko2, is prepared by fusing selenious anhydride with nitre, ex- tracting with water, and evaporating. It crystallizes in forms exactly resembling those of potassic sulphate. It may be distinguished from this salt by evolving chlorine when heated with hydrochloric acid, at the same time undergoing reduction to potassic sele- nite. The selenite, SeOKo2, forms granular, very soluble deliquescent crystals. Potassic tellur ate, Te02Ko2. Hydric potassic tellurate, 2Te02H0K0,30H2. These salts are obtained by adding the requisite quantities of telluric acid to solutions of potassic carbonate. The neutral salt is very soluble, the acid salt sparingly soluble, in cold water. Other more complex tellurates of potassium are known (see pp. 289, 290). Potassic Phosphates.—a. Potassic orthophosphate, POKo3, is prepared by igniting 2 molecules of phosphoric anhydride with 3 mole- cules of potassic carbonate, dissolving in water and evaporating. It forms colorless, very soluble needles.—Hydric dipotassic orthophos- phate, POHoKo2, may be obtained in solution by adding potassic car- bonate to a solution of phosphoric acid till a slight alkaline reaction is produced. It is uncrystallizable.—Dihydric potassic orthophosphate, POHo2Ko, is prepared by adding phosphoric acid to a solution of potassic carbonate till the liquid has a strongly acid reaction. On evaporating, large colorless quadratic crystals, very soluble in water, are obtained. ( POKo2 b. Potassic pyrophosphate, < O ,30H2, is prepared by igniting ( POKo2 hydric dipotassic phosphate (cf. p. 355). It may also be obtained by almost neutralizing a solution of phosphoric acid with alcoholic potash, then adding alcohol as long as milkiness is produced, and sepa- rating, drying and igniting the syrupy precipitate. The mass is extracted with water and evaporated to the point of crystallization. It forms a radio-crystalline mass, very soluble in water. One molecule of water of crystallization is driven off at 100° C., but a temperature of 300° C. (572° F.) is required to render the salt anhydrous. In the 420 INORGANIC CHEMISTRY. anhydrous state it is deliquescent.—Dihydric dipotassic pyrophosphate, P203Ho2Ko2, is obtained by precipitating with alcohol the solution of the neutral salt in acetic acid. The syrupy mass is washed with alco- hol to remove the potassic'acetate and dried over sulphuric acid. It forms a white deliquescent mass. c. Potassic metaphosphate, POJvo, is prepared by igniting dihydric potassic phosphate (cf. p. 354). It is thus obtained as a translucent mass, almost insoluble in water, readily soluble in dilute acids. Meta- phosphates of complex constitution are also known (p. 354). Potassic phosphite, PHoKo2.—This salt is obtained by neutralizing the aqueous acid with potassic hydrate or carbonate and evaporating in vacuo. It is deliquescent and very soluble, and can only with difficulty be obtained in a crystalline form. Potassic arsenates.—These are prepared like the corresponding phosphates, with which they are isomorphous, and which they closely resemble in their other properties. Po- tassic arsenate, AsOKo3, forms deliquescent needles; hydric dipotassic arsenate, AsOHoKo2, is uncrystallizable and deliquescent; dihydric potassic arsenate, AsOHo2Ko, which is most readily obtained by fusing arsenious acid with nitre, extracting with water and evaporating, forms large soluble quadratic crystals. Very little is known concerning the arsenites of potassium. Potassic aMimonates.—When a mixture of 1 part of powdered antimony with 4 parts of nitre is deflagrated, and the mass extracted with tepid water, potassic metantimonate, Sb02Ko, remains as a white powder, almost insoluble in cold water. When this sub- stance is boiled with water it gradually dissolves, taking up the elements of water and forming dihydric potassic antimonate, which, on evaporating the solution to a syrup, sepa- rates out in granular crystals of the formula 2SbOHo2K0,301T2. By fusing antimonic acid or either of the above antimonates with a large excess of potash, dissolving the mass in water and evaporating, warty crystals of tetrapotassic pyrantimonate, Sb203Ko4, are obtained. This salt is stable in solution only in presence of an excess of caustic potash; pure water decomposes it into free potash and dihydric dipotassic pyrantimonate (metantimonate of Frerny), Sb203Ho2Ko2,60H2, a granular, almost insoluble powder, which is converted by long boiling with water into soluble dihydric potassic antimo- nate (see above.) . Potassic borate.—The metaborate, BOKo, is prepared by fusing together equal mole- cules of boric anhydride and potassic carbonate. It is very soluble, and crytallizes with difficulty. Exposed to the air in solution, it absorbs carbonic anhydride and is converted into hydric potassic diborate, B202H0K0,20H2. A dipotassic tetraborate, B405Ko2,60H2, is obtained by mixing hot concentrated solutions of 1 molecule of po- tassic carbonate and 2 molecules of boric anhydride, and cooling to 6° C. (42.8° F.). The salt crystallizes in hard, transparent, prismatic crystals, with a vitreous lustre. When a boiling solution of potassic carbonate is acidified with boric acid, it deposits on cooling prismatic crystals of hydric potassic hexaborate, B608H0K0,40H2. Potassic silicate is formed when silicic acid or amorphous silicic anhydride is dissolved in potassic hydrate. It is generally prepared by fusing together potassic carbonate and white quartz sand. No compound of definite composition is known. Potassic silicate, under the name of soluble glass, is employed as a cement. COMPOUNDS OF POTASSIUM WITH SULPHUR. The following have been obtained: Dipotassic sulphide, . . . sk2. K—S—K Dipotassic disulphide, . k2s9. K—S—S—K Dipotassic trisulphide, . . . k2s; K—S—S—S—K Dipotassic tetrasulphide, . K2s4. K-S-S-S-S-K Dipotassic pentasulphide, • K2S5. K—S—S—S—S—S—K Dipotassic heptasulphide, . k2s7? K—S—S—S—S—S—S—S—K COMPOUNDS OF POTASSIUM. 421 Dipotassic sulphide, SK2, is formed when potassic sulphate is reduced by ignition with carbon or in a current of hydrogen: SO2KOj + 4H2 = SK2 + 40 H2. Potassic Dipotassic Water, sulphate. sulphide. It is a reddish crystalline mass, which deliquesces when exposed to the air. A solution of dipotassic sulphide maybe obtained by dividing a con- centrated aqueous solution of potassic hydrate into two equal parts, sat- urating one part with sulphuretted hydrogen so as to form potassic sulphhydrate (q.v.), and then adding the other part: KHo + KHs = SK2 + 0H2. Potassic Potassic Dipotassic Water, hydrate. sulphhydrate. sulphide. The concentrated solution deposits deliquescent prismatic or tabular crystals of the formula SK2,50H2. Dipotassic disulphide, K2S2, is formed when the sulphhydrate is oxidized by exposure to air: 2KHs + O = K.2S2 + OH2. Potassic Dipotassic Water, sulphhydrate. disulphide. By evaporation in vacuo the disulphide is obtained as an orange-colored mass. The other polysulphides of potassium are prepared by fusing dipotassic sulphide with sulphur. Below 600° C. (1112° F.) the pentasulphide is formed ; between 600° and 800° C. (1112-1472° F.) the tetrasulphide; and at 900° C. (1652° F.) the trisvl- phide. They are brownish-yellow solids with an alkalide reaction. Exposed to moist air they emit an odor of sulphuretted hydrogen. Solutions of these polysulphides are formed when solutions of dipotassic sulphide are boiled with the requisite quantities of flowers of sulphur. In this way crystallized aquates of some of these sulphides may be obtained, for example K2S4,20H2, which forms orange-colored laminae. COMPOUND OF POTASSIUM WITH HYDROSULPHYL. Potajssic sulphhydrate, KHs, is obtained by heating potassium in a current of sulphuretted hydrogen; 2SH2 + K2 = 2SKH + H2, Sulphuretted Potassic hydrogen. sulphhydrate. or by passing sulphuretted hydrogen over potassic carbonate heated to low redness: COKo2 + 2SH2 = 2SKH + 0O2 + 0H2. Potassic Sulphuretted Potassic Carbonic Water, carbonate. hydrogen. sulphhydrate. anhydride. 422 INORGANIC CHEMISTRY. It forms a flesh-colored crystalline mass, which melts at low redness to a yellow liquid. A solution of potassic sulphhydrate may be obtained by saturating an aqueous solution of potassic hydrate with sulphuretted hydrogen : KHo + SH2 = KHs + 0H2. Potassic Sulphuretted Potassic Water, hydrate. hydrogen, sulphhydrate. The solution, when concentrated in vacuo, deposits colorless rhom- bohedra of the formula 2KHs,OET2, Reactions of potassic sulphhydrate, dipotassic sulphide and the higher potassic sulphides.—l. Potassic sulphhydrate and dipotassic sulphide, when acted upon by acids, yield sulphuretted hydrogen: KHs + HCI = KCI + SH2. Potassic Hydrochloric Potassic Sulphuretted sulphhydrate. acid. chloride. hydrogen. SK2 + 2HCI = 2KCI + SH2. Dipotassic Hydrochloric Potassic Sulphuretted sulphide. acid. chloride. hydrogen. 2. The higher sulphides, when similarly treated, yield sulphuretted hydrogen with precipitation of sulphur: K2S3 + 2HCI = 2KCI + SH2 + Sa. Dipotassic Hydrochloric Potassic Sulphuretted trisulphide. acid. chloride- hydrogen. 3. When dipotassic sulphide is exposed in aqueous solution to the action of the air, it absorbs oxygen and is converted into a mixture of potassic thiosulphate and potassic hydrate : 2SK2 + 0H2 + 202 = S02KoKs + 2KHo. Di potassic Water. Potassic Potassic sulphide. thiosulphate. hydrate. 4. A mixture of the higher potassic sulphides with potassic thio- sulphate, known under the name of hepar sulphuris or liver of sulphur, may be obtained as a brown mass by heating potassic carbonate with sulphur : 3COKo2 + 4S2 = 2K2S3 + S02KoKs + 3C02.. Potassic Dipotassic Potassic Carbonic carbonate. trisulphide. thiosulphate. anhydride. 30OKo2 + 6S2 = 2K2S5 + S02KoKs + 3C02. Potassic Dipotassic Potassic Carbonic carbonate. pentasulphide. thiosulphate. anhydride. 5. The last mixture, when acted upon by acids, undergoes suc- cessively the following decomposition: COMPOUNDS OF POTASSIUM. 423 2K2S5 + S02KoKs + 6HCI = 6KCI Dipotassic Potassic Hydrochloric Potassic pentasulphide. thiosulphate. acid. chloride. + S02HoHs + 2SH2 + 4S2; Thiosulphuric Sulphuretted acid. hydrogen. then— S02HoHs = S02 + S + 0H2' Thiosulphuric Sulphurous Water, acid. anhydride. and finally— rso.Ho 5SH2 + 5S02 = { S"3 + 40H2 + SS. ( S02Ho Sulphuretted Sulphurous Pentathionic Water, hydrogen. anhydride. acid. SULPHO-SALTS OF POTASSIUM. Potassic sulpharsenate, AsS//Ks3, is prepared by dissolving arsenic sulphide, or arsenious sulphide together with sulphur, in a solution of potassic sulphide or potassic sulph hydrate: Aa,S"# + 3SK2 = 2AsS//Kss. Arsenic Potassic Tripotassic sulphide. sulphide. sulpharsenate. Ab,S"B + 3SK2 + S2 = AsS//Ks3. Arsenious Potassic Tripotassic sulphide. sulphide. sulpharsenate. It is also formed when a solution of tripotassic arsenate is saturated with sul- phuretted hydrogen: AsOKo3 + 3SH2 = AsS//Ks3 + 80 H2. Tripotassic Sulphuretted Tripotassic Water, arsenate. hydrogen. sulpharsenate. It is obtained as a deliquescent crystalline mass of the formula AsS//Ks3,OH2 (per- haps AslloHsKss). Potassic sulphantimonate, SbS//Ks3, may he obtained in the same manner as the sulpharsenate, employing the corresponding sulphides of antimony. In practice, it is prepared by heating together finely powdered antimonious sulphide, sulphur, potassic carbonate, slaked lime and water, filtering and evaporating. It forms yellow deli- quescent crystals of the formula 2SbS//Ks3,90H2. Treated with dilute acids in the cold, the alkaline sulph arsenates and sulphantimo- nates yield the corresponding acids AsS//Hs3 and SbS//Hs3. On boiling the solutions these acids are decomposed into arsenic and antimonic sulphides respectively, and sul- phuretted hydrogen: 2AsS//Hs3 = As2S//5 + 3SH2. Sulpharsenie Arsenic Sulphuretted acid. sulphide. hydrogen. COMPOUND OF POTASSIUM WITH NITROGEN AND HYDROGEN. Potassic amide, NKH2, is obtained by heating potassium gently in a current of dry gaseous ammonia. The potassium fuses in the gas to a blue liquid, which solidifies on cooling to a flesh-colored mass. Water decomposes it with violence into ammonia and potassic hydrate; 424 INORGANIC CHEMISTRY. nktt2 -f- ott2 NH3 -f" Oku. Potassic Water. Ammonia. Potassic amide. hydrate. When strongly heated in an atmosphere free from oxygen, it is decomposed into ammonia and potassic nitride: 3NKH2 = 2NHs + NK3. Potassic Ammonia. Potassic amide. nitride. Potassic nitride is a greenish-black substance which, in contact with air, sponta- neously inflames. General Properties and Reactions of the Compounds of Potassium.—The salts of potassium with colorless acids are colorless. Platinic chloride precipitates from hydrochloric acid solutions of potash salts a yellow crystalline powder of potassic platinic chloride (PtCl4,- 2KCI), very sparingly soluble in water, insoluble in alcohol and ether; this salt, when heated to redness, is decomposed with evolution of chlo- rine, leaving potassic chloride and metallic platinum. Hydrofluosilicic acid gives a gelatinous precipitate of potassic slicofluoride, SiK2F6. Tartaric acid in excess precipitates from moderately concentrated solu- tions hydric potassic tartrate, as a white crystalline powder. The compounds of potassium impart to a non-lumiuous flame a violet coloration which, when viewed through blue cobalt glass or a solution of indigo, appears red. The spectrum of potassium contains two characteristic lines—K« in the red and K/3 in the violet—both coin- cident with lines of the solar spectrum. SODIUM, Na2? Atomic weight = 23. Probable molecular weight = 46. Sp. gr. 0.97. Fuses at 95.6° C. (172° F.). Foils at a red heat. Atomicity Evi- dence of atomicity: Sodic chloride, . . . , Sodic hydrate, . , . . . NaCI. , . . . ONaH. Sodic oxide, . . . . ONa2. History.—Metallic sodium was first obtained by Davy, in 1807, by the electrolysis of sodic hydrate. Occurrence.—Sodium is an abundant and widely distributed element. It does not occur in the free state. In combination with silicic acid it is found in many minerals and rocks, and in soils. As nitrate, or Chili saltpetre, it forms large beds on the surface of the ground in dry districts in Chili and Peru. As carbonate and as iodide it occurs in the ashes of sea plants. The chloride is, however, the form in which it is found in the greatest abundance—thus, as rock salt, in sea water, and in the water of salt springs. The borate and sulphate also occur in nature. Preparation.—l. Davy obtained sodium by electrolyzing, between SODIUM. 425 the poles of a powerful battery, solid sodic hydrate moistened with water (see Potassium, p. 411): 20NaH = Na2 + H2 + Oa. 2. Sodium is also liberated from the hydrate by acting upon it with metallic iron at a strong white heat. The reaction is the same as in the case of potassium (p. 412). 3. On a manufacturing scale, sodium is prepared by distilling from a cylindrical iron retort a mixture of dry sodic carbonate and charcoal, to which a small quantity of chalk is added to prevent the fusion of the mass and the consequent separation of the charcoal: OONaoa + 20 = Na2 + 300. Sodic Carbonic carbonate. oxide. Properties.—Sodium resembles potassium in its properties. It is a lustrous, silver-white metal, which almost instantaneously tarnishes from oxidation when exposed to the air. At a temperature of —2o° C. (—4° F.) it is hard, but at ordinary temperatures it is of the consistence of wax. When heated in air it burns with a yellow flame, forming oxides of sodium. By fusing it in a tube filled with coal-gas, allowing it partially to solidify, and pouring off the still liquid portion, it may be obtained in crystals. Reactions.—The reactions of sodium are similar to those of potassium, but less energetic. Thus, sodium decomposes water with evolution of hydrogen, the metal moving rapidly on the surface with a hissing noise, but the heat developed is not sufficient to inflame the hydrogen. If, however, the water be previously heated above 60° C. (140° F.), or if, by rendering the water viscid with glue, or by placing the metal on wet blotting paper, the sodium be prevented from moving, and therefore from too rapidly cooling, the hydrogen will inflame. Under these cir- cumstances, the reaction is, however, sometimes attended with a violent explosion. Sodium is not acted upon by dry chlorine or bromine, even when gently heated with these reagents; in presence of moisture, how- ever, chloride and bromide of sodium are formed. Uses.—Sodium, like potassium, is employed in the preparation of various metals and metalloids from their oxides or chlorides. It acts by combining with the oxygen or chlorine, and liberating the element which it is desired to isolate. On account of its greater cheapness and lower atomic weight, it is generally preferred for this purpose to potas- sium (see p. 413). It is thus used in the arts, in the preparation of aluminium and magnesium from their chlorides. In the laboratory it is also employed as a source of nascent hydrogen. The substance to be submitted to the hydrogenating action is brought, along with water or alcohol, in contact with the sodium (preferably in the form of an amalgam, or alloy of the metal with mercury—the mercury being added in order to moderate the violence of the reaction), and in this way the hydrogen from the water or alcohol, instead of being liberated, combines with the substance. 426 INORGANIC CHEMISTRY. COMPOUND OF SODIUM WITH HYDROGEN.' Sodic hydride, N:i4TI2. Sodium when heated to a temperature between 300° and 420° C. (572°-788° F.) in a current of dry hydrogen, absorbs the gas with formation of sodic hydride, a silvery metallic mass of sp.gr. 0.959, which is soft at ordinary tem- peratures, but at lower temperatures brittle. It fuses at a somewhat lower temperature than sodium. It is more permanent in air than the corresponding potassium com- pound. It begins to dissociate under ordinary pressures at 420° C. (788° F.); in vacuo, at 300° C. (572° F.). COMPOUNDS OF SODIUM WITH THE HALOGENS. Sodic chloride (Common salt), NaCl.—This important compound occurs in sea-water (2.5 to 3 per cent.), in salt springs, and as rock salt. The most celebrated salt mines are those of Wieliczka, in Galicia, in which the salt deposit is 500 miles long, 20 miles broad, and J2OO feet thick. When the salt is pure, as is sometimes the case with rock salt, it is obtained direct by ordinary mining operations. Generally, how- ever, it is contaminated with earthy matters, from which it must be freed by dissolving in water and recrystallizing. Salt is also obtained from sea-water; in warm climates, by allowing the water to evaporate in shallow basins; in cold climates, by letting it freeze and removing the ice, the salt remaining in the liquid! Chloride of sodium is formed when sodium is burnt in chlorine. It crystallizes in large colorless anhydrous cubes belonging to the regular system; from solutions con- taining urea it is deposited in octahedra. Below —lo° C. it crystallizes from water in monoclinic plates of the formula NaCI,2OH2, which at ordinary temperatures part with their water of crystallization and fall to pieces, being converted into a number of minute cubes. It is almost equally soluble in hot and cold water: at 0° C. water takes up 36 parts, at 100° O. 39 parts. Alcohol does not dissolve it. At a red heat it is fusible and volatile. Sodic bromide, Naßr, is prepared by neutralizing hydrobroraic acid with sodic carbonate, or by decomposing ferrous bromide (Feßr2) with a solution of sodic carbonate (see Potassic iodide, p. 414). It crys- tallizes from its aqueous solution above 30° C, in anhydrous cubes; below this temperature in monoclinic prisms of the formula Naßr,2oH2. It is readily soluble both in water and in alcohol. Sodic iodide, Nal, is prepared like the bromide, which it also resem- bles in its crystallographical characteristics. Above 20° C. it crystal- lizes from water in anhydrous cubes; at lower temperatures in mono- clinic forms with 2 molecules of water of crystallization. Both water and alcohol dissolve it freely. Like potassic iodide it forms double compounds with the iodides of the heavy metals. Sodic fluoride, NaF, is obtained by neutralizing hydrofluoric acid with sodic carbonate. It crystallizes in anhydrous cubes, which are soluble in 25 parts of cold, very slightly more soluble in boiling water. It forms numerous double compounds with other fluorides and with hydrofluoric acid. The mineral cryolite is an aluminio-sodic fluoride of the formula Al2F6,6NaF. Sodic silicofluoride., SiFeNa2 (=SiF4,2NaF), forms small lustrous hexagonal crystals, sparingly soluble in water. 427 COMPOUNDS OF SODIUM. COMPOUNDS OF SODIUM WITH OXYGEN AND HYDROXYL. Sodic oxide, ONa2.—When sodium burns in air a mixture of sodic oxide with disodic dioxide (Na2o2) is formed. By heating this mixture to a very high temperature, the dioxide parts with half its oxygen, and is converted into sodic oxide, which is thus obtained as a gray mass with a conchoidal fracture. Water converts it, with evolution of great heat, into the hydrate. f ONa Disodic dioxide, < Q\pp ]'s obtained by heating sodium in oxygen gas till the weight becomes constant. It is a white substance, which becomes yellow on heating, but turns white again on cooling. In con- tact with water, it evolves great heat, and parts with some of its oxygen. Sodic hydrate (Caustic soda), NaHo.—This compound is formed by the action of water upon sodium or upon sodic oxide. It is prepared by acting upon a boiling solution of sodic carbonate with milk of lime: CONao2 + CaHo2 = 2NaHo + COCao". Sodic Calcic Sodic Calcic carbonate. hydrate. hydrate. carbonate. The solution of sodic hydrate is decanted from the insoluble calcic car- bonate and concentrated, first in an iron and lastly in a silver basin. Most of the sodic hydrate of commerce is obtained in the manufacture of sodic carbonate (see Leblanc’s process), the calcic oxide, which is formed in roasting the black ash, acting upon a portion of the sodic carbonate when the mass is treated with water. The caustic soda remains in the mother liquors after the separation of the other salts— carbonate and sulphate. A small quantity of sodic nitrate is added in order to oxidize the sodic sulphide to sulphate.—Sodic hydrate is an opaque white fibrous substance of sp. gr. 2.00, resembling potassic hydrate in nearly all its properties. It fuses below redness, and at a higher temperature volatilizes. When exposed to the air in large masses, it does not deliquesce, but merely becomes moist on the surface, after which a coating of the non-deliquescent carbonate is formed, which pro- tects it from further action. It is very soluble, both in water and in alcohol, yielding powerfully caustic solutions. The concentrated aqueous solution, when exposed to a low temperature, deposits crystals of the formula 2Na110,70H2, which fuse at 6° C. (43° F.) to a liquid of sp. gr. 1.405. Its solutions absorb carbonic anhydride from the air. With acids it yields the corresponding sodium salts: NaHo + N02Ho = N02Nao + 0H2. Sodic Nitric acid. Sodic nitrate. Water, hydrate. OXY-SALTS OF SODIUM. Sodic nitrate (Chili saltpetre), N02Nao, occurs, more or less con- taminated with other salts, in enormous deposits in Chili and Peru. It can be readily purified by crystallization, and forms rhomboh is formed in the same manner as the potassium salt (p. 182), but, owing to its solubility and the impossibility of separating it from the chloride which is formed simultaneously, cannot be so prepared. It is most readily obtained by neutralizing a solution of chloric acid with sodic carbonate and evaporating. It forms large transparent crystals belonging to the regular system, and exhibiting hemi- hedral faces, which in some crystals are positive, in others negative. These crystals possess a corresponding action on the ray of polarized light, the positive crystals being dextrorotatory, the negative Isevorotatory. It is soluble in its own weight of water at ordinary temperatures, and in half its weight at 100° C. In other respects it resembles the potassium salt. roci Sodic perchlorate, O ,is prepared by neutralizing perchloric acid with sodic hy- ( ONao drate or carbonate. It is a deliquescent salt, readily soluble in water, soluble also in alcohol. lustrous crystals, soluble in about 3 parts of water at ordinary temperatures. Below —4° C. (25° F.) it crystallizes in four-sided prisms containing water of crystallization. Sodic iodate, | QNao ’*s °^ta^ned in the same manner as the potassium salt (p. 417). It crystallizes at ordinary temperatures with one molecule of water of crystallization in silky needles. It is soluble in 11-12 parts of water. Below 5° C. (41° F.) it is deposited in transparent rhombic prisms with 5 molecules of water of crystallization. It forms well-crystallized double salts with the chloride, bromide, and iodide of sodium. The compound with sodic chloride has the formula Sodic bromate, | q^!"10 ,is prepared like the potassium salt (p. 417). It forms small 2{SNao'3NaCl.9°H!- fOl Sodic periodate, ■( O ,30H2. When chlorine is passed into a solution of sodic ( ONao iodate in caustic soda, a sparingly soluble basic salt of the formula I0SHNa2,OH2 is deposited, which, when dissolved in dilute nitric acid and evaporated, is converted into the normal salt I0315ra0,30H2. (On the formulation of the periodates, see p. 305.) The normal salt forms colorless hexagonal crystals, soluble in 12 parts of water at ordinary temperatures. The crystals part with their water of crystallization at 100° C. Heated to 275° C. (527° F.) the anhydrous salt gives off oxygen, and is con- verted into iodate. Sodic carbonate, CONao2, occurs in the soda lakes of Egypt and Hungary, and in the volcanic springs of Iceland. It constitutes the greater part of the ash of sea plants, from which source it was formerly obtained. The two methods at present employed in its preparation are: the process of Leblanc and the ammonia-soda process, both of which start from sodic chloride. 1. Leblanc’s Process.—This process consists of two parts : the con- version of the sodic chloride into sodic sulphate or salt cake, known as the salt-cake process ; and the manufacture of sodic carbonate or soda COMPOUNDS OP SODIUM. 429 ash from the sulphate, known as the soda-ash process. In the first of these processes the sodic chloride is treated, in a large hemispherical cast-iron pan heated over a furnace, with the requisite quantity of sul- phuric acid. The hydrochloric acid which is evolved passes through towers filled with coke, over which a stream of water trickles, and is thus absorbed. After heating for some time, the mixture of acid and salt solidifies, upon which it is transferred from the iron pan to the bed of a reverberatory furnace, where the decomposition is completed. In the soda-ash processs, the sodic sulphate or salt cake, as it is tech- nically termed, is mixed with crushed chalk or limestone and small coal, and gradually heated in a reverberatory furnace. The action takes place in the two following stages: S02Nao2 + 4C = SNa2 + 4CO, Sodic Sodic Carbonic sulphate. sulphide,* oxide, and SXa2 + COCao" = COXao2 + OaS" Sodic Calcic Sodic Calcic sulphide. carbonate. carbonate. sulphide, the calcic sulphide combining with the excess of calcic oxide (formed from the chalk), and yielding insoluble calcic oxysulphide. These reactions take place simultaneously in the above operation. When the change is complete, the mass, which is known as black ash, is allowed to cool, and is then extracted with water, which dissolves the sodic carbonate, leaving behind the insoluble oxysulphide. On evaporating, the sodic carbonate crystallizes out, and may be purified by recrystallization. A portion of the chalk is converted by the heat into quicklime, and this gives rise to the formation of sodic hydrate when the mass is treated with water. This sodic hydrate may be recovered from the mother liquors of the carbonate (p. 427). 2. Ammonia-soda Process.—By the action of hydric ammonic car- bonate on a concentrated solution of sodic chloride, hydric sodic car- bonate and ammonic chloride are produced : 0OHo(NvH4O) + NaCl = COHoXao + NH4CI. Hydric ammonic Sodic Hydric sodic Amnionic carbonate. chloride. carbonate. chloride. The sparingly soluble hydric sodic carbonate separates out, whilst the ammonic chloride remains in solution. By heating the hydric sodic carbonate, it is converted, with evolution of carbonic anhydride, into the normal salt: 2COHoNao = OONao, + C02 + OH2. Hydric sodic Sodic Carbonic Water, carbonate. carbonate. anhydride. The carbonic anhydride is employed in reconverting into hydric am- monic carbonate, the ammonia recovered from the ammonic chloride. 430 INORGANIC CHEMISTRY. Sodic carbonate crystallizes at ordinary temperatures in efflorescent monoclinic crystals of the formula CONao2,100H2, fusing at 50° C. (122° F.) to a clear liquid, which gives off water, and deposits a pulverulent salt, with one molecule of water of crystallization. At temperatures between 30° and 50° C. (86-122° F.) it is deposited in rhombic crystals with 70H2. It is readily soluble in water, with a maximum solubility at 38° C. (100° F.). 100 parts of water dissolve: At 0° C. (32° F.), . . 7 parts of anhydrous salt. At 15° C. (59° F.), • . 16 parts of anhydrous salt. At 38° C. (K)0( 3 F.), • . 51 parts of anhydrous salt. At 104° C. (219' 3 F.), • . 45 parts of anhydrous salt. Anhydrous sodic carbonate fuses at a bright red heat, and may be volatilized at a white heat. The chief consumption of sodic carbonate is in the manufacture of glass, in soap-making, and in bleaching calico. Hydric sodic carbonate, COHoNao, occurs naturally in many mineral waters. It is formed when a concentrated solution of the normal car- bonate is saturated with carbonic anhydride. The crystallized normal carbonate also absorbs carbonic anhydride with evolution of heat, a property which is taken advantage of in the preparation of the salt on a large scale. The acid carbonate can be readily separated from the normal carbonate by its more sparing solubility. Hydric sodic carbon- ate is also obtained in the preparation of sodic carbonate by the ammo- nia-soda process (p. 429). It forms monoclinic prisms, soluble in 10-11 parts of water at ordinary temperatures. When its solution is heated, the salt parts with a portion of its carbonic acid, yielding the so-called sesquicarbonate, CONao2,2C0H0Na0,20H3, which may be ob- tained in crystals by cooling the solution. The sesquicarbonate also occurs in large deposits in Africa and South America, the natural pro- duct being known as trona or urao. If the solution of hydric sodic catbonate be boiled for a sufficient length of time, it is entirely decom- posed into normal carbonate, carbonic anhydride, and water. The same decomposition takes place when the dry salt is heated. Potassic sodic carbonate, COKoNao,6O Ha.—This salt crystallizes from the solution of a mixture of potassic and sodic carbonates. It forms efflorescent monoclinic crystals. It cannot be recrystallized from water without decomposition. The anhydrous salt fuses at a red heat more readily than either potassic or sodic carbonate. On account of this property it is employed in mineral analysis for the decomposition of silicates by fusion. Some sulphate [Glaubers salt), S02Nao2, occurs in nature in the anhydrous form as the mineral thenardite, and with ten molecules of water of crystallization as Glauber’s salt Glauberite is a native sodic calcic sulphate of the formula Sodic sulphate often oc- curs in sea-water and in the water of brine springs. It is prepared in enormous quantities under the name of salt cake as a preliminary step in the manufacture of sodic carbonate by Leblanc’s process (p. 429). 431 COMPOUNDS OP SODIUM. It crystallizes at ordinary temperatures in large colorless efflorescent monoclinic prisms of the formula SO2Nao2,100H2, which fuse at 33° C. in their water of crystallization. It is very soluble in water, with a maximum solubility at 33° C. 100 parts of water dissolve: At 0° a, . . . . 5 parts of anhydrous salt. At 20° a, . . . . 20 . parts of anhydrous salt. At 33° c., . . . . 50.6 parts of anhydrous salt. At 103° a, . , . . 42.65 parts of anhydrous salt. (See also p. 127). A solution saturated at 33° C. deposits, when heated above this temperature, small rhombic octahedra of the formula SO2Nao2,0H2 (formerly supposed to be anhydrous; see, however, Thompson, Ber. d. deutsch. diem. Ges., 11, 2042). This monaquate is always deposited from solutions at temperatures above 40° C. (101° F.). When a solution, saturated at 33° C. (91° F.) is cooled, it does not, if protected from the air, deposit crystals, and in hermetically sealed ves- sels, may be preserved for an indefinite period in this supersaturated condition ; but the introduction of a fragment of the solid salt, or even contact with dust from the air, which probably always contains the solid salt, is sufficient to determine the solidification of the liquid to a magma of crystals, this process being accompanied by a rise of temper- ature. When the supersaturated solution is evaporated in vacuo over sulphuric acid, it deposits crystals of a salt having the formula so2n ao2,70H2, this probably being the form in which the substance is present in the supersaturated solution. Crystallized sodic sulphate dissolves in concentrated hydrochloric acid with great absorption of heat. A useful freezing mixture is obtained by pouring 5 parts of the acid upon 8 of the sulphate.—Hydric sodic sulphate, S02HoNao, is prepared like the potash salt (p. 418). It crystallizes at ordinary temperatures in monoclinic prisms with 1 aq.;* above 50° C,, in anhydrous triclinic forms. It is readily fusible. Heated above its fusing point it parts with the elements of water, yielding sodic pyrosulphate, S205Nao2; at a still higher temperature sulphuric anhydride is expelled and the nor- mal sulphate remains. Tripotassic sodic disulphate, is obtained in hexagonal plates when mixed solutions of sodic and potassie sulphate are allowed to crystallize. At the mo- ment of crystallizing, the salt emits flashes of light, visible in the dark, the phenomenon being most striking when the temperature of the solution is about 40° C. Sodic sulphite, SONao2,70H2, forms monoclinic crystals, readily soluble in water and possessing an alkaline reaction. When the solu- tion is heated, it deposits an anhydrous salt, which dissolves again on cooling. Hydric sodic sulphite, SOHoISTao, crystallizes in small lus- trous prisms, readily soluble in water, and possessing an acid reaction. The salt evolves sulphurous anhydride when exposed to the air, and is spontaneously oxidized to sulphate. Sodic pyrosidphite, S2OsNao2, is * In the aquates the symbol “ aq.” is frequently employed to denote a molecule of water of crystallization. 432 INORGANIC CHEMISTRY. also known. The sulphites of sodium are prepared like the correspond- ing potassium salts (p. 418). Sodic dithionate, | gQ2^®,2OH2, is prepared like the potassium salt (p. 419). It forms transparent rhombic prisms, readily soluble in water. Sodic thiosulphate (Sodic hyposulphite), S02NaoNas,soH2. (Preparation, see p. 277.) This salt is obtained on a large scale from soda waste, the insoluble matter which remains after the extraction of the sodic carbonate from the black ash in Leblanc’s process. By exposing this residue in a moist condition to the air, the calcic sulphide (or oxy- sulphide) which it contains is oxidized to calcic thiosulphate, calcic hydrate being formed at the same time: 2CaS" + 202 + 0H2 = S02/gCa) + CaIIo2. Calcic Water. Calcic Calcic sulphide. thiosulphate. hydrate. The calcic thiosulphate is extracted with water and decomposed with sodic sulphate, thus yielding sodic thiosulphate and insoluble calcic sulphate. Sodic thiosulphate forms large, well defined monoclinic crys- tals, readily soluble in water and somewhat deliquescent. It fuses at 56° C. (133° F.) in its water of crystallization. When the dry salt is heated it is decomposed like the potassium salt (p. 419) into a mixture of sulphate and sulphide. The aqueous solution dissolves the chloride, bromide, and iodide of silver, a property which has caused the salt to be employed in fixing photographic prints. Sodic thiosulphate is also used as an antichlore, to destroy the excess of the chlorine employed in bleaching vegetable fibre. Sodic selenate, SeO2Nao2,IOOH2, is prepared liks the potassium salt (p. 419). It closely resembles sodic sulphate in its properties. Sodic tellurate, TeChNaoz, resembles the potassium salt. Sodic Phosphates : a. Sodic phosphate, PONao3,120H2, is prepared by fusing 2 mole- cules of hydric disodic phosphate with 1 molecule of sodic carbonate, dissolving in water and crystallizing; or by evaporating a solution of hydric disodic phosphate in caustic soda. The salt crystallizes in thin six-sided prisms, readily soluble in wuter, efflorescent in dry air. The solution, which has a strong alkaline reaction, absorbs carbonic anhy- dride from the air, the third atom of sodium being thus abstracted to form carbonate, whilst hydric disodic phosphate remains.—Hydric disodic phosphate (“phosphate of soda”) POHoNao2,120H2, is obtained by adding sodic carbonate or sodic hydrate to orthophosphoric acid until the liquid has a slight alkaline reaction, and then evaporating to the crystallizing point. On a large scale the orthophosphoric acid for this purpose is obtained by decomposing bone-ash with the requisite quan- tity of dilute sulphuric acid and filtering from the insoluble calcic sul- phate. The salt forms efflorescent monoclinic prisms, soluble in 4.5-5 COMPOUNDS OF SODIUM. 433 parts of water at ordinary temperatures. The solution has a weak alkaline reaction. At 37° C. (99° F.) the crystals fuse in their water of crystallization. At temperatures above 30° C. (86° F.) the solution deposits non-efflorescent crystals of a salt with 7 aq. When heated to redness hydric disodic phosphate parts with the elements of water, forming tetrasodie pyrophosphate, P2O,Nao4. Hydric sodic phosphate was formerly much used in calico-printing under the name of “dung substitute,” but is now superseded by the cheaper sodic arsenate.— Dihydric sodic phosphate, POHo2Na0,0H2, is obtained by adding phosphoric acid to the disodic salt until the solution no longer yields a precipitate with baric chloride, and then evaporating. It crys- tallizes in rhombic prisms, readily soluble in water, yielding an acid solution. Hydric potassic sodic phosphate, P()IToKoNao,7OH2, is prepared by adding potassic carbonate to a solution of dihydric sodic phosphate until the liquid has a slight alkaline reaction. It forms soluble monoclinic crystals. h. Sodic pyrophosphate, P203lSrao4,100H2, is prepared by heating hydric disodic phosphate to redness, dissolving the mass in water and allowing to crystallize (p. 355). It is thus obtained in large monoclinic crystals, soluble in 10-12 parts of water at ordinary temperatures, and in their own weight of water at 100° C. The aqueous solution may be boiled without alteration, but when boiled with hydrochloric, nitric, or even acetic acid, the salt takes up the elements of water, at the same time parting with a portion of its base to the acid, and is converted into dihydric sodic phosphate.—Dihydric disodic pyrophosphate, P20311o2]Srao2, separates as a crystalline powder when alcohol is added to a solution of the normal pyrophosphate in acetic acid. It may be boiled with water without decomposition. Dipotassic disodic pyrophosphate, P203Ko2lSrao2, is obtained by neutralizing a solution of the acid sodium salt with potassic carbonate. It forms soluble acicular crystals. c. Sodic metaphosphate, P()2Nao, is prepared by igniting either dihy- dric sodic phosphate, or hydric ammonic sodic phosphate, or dihydric disodic pyrophosphate (see Metaphosphates, p. 354). According to the temperature to which the substance has been heated and the rate of cooling, products differing widely in their properties, but all possessing the same percentage composition, are obtained. When the substance is heated to redness and rapidly cooled, the product is a vitreous deliques- cent mass, which dissolves readily in water and remains behind on evaporation in the form of an uncrystallizable gum. If the cooling has been effected more slowly, there is formed, in addition to the uncrystal lizable salt, a compound which is deposited from the solution in monoclinic prisms of the formula PO2Na0,20H2. A third modifi- cation is obtained by limiting the temperature to 315° C. (599° F.). On extracting with water, an almost insoluble metaphosphate remains as a white powder. These differences are supposed to depend upon polymeric modification (see p. 354.) 434 INORGANIC CHEMISTRY. Sodic arsenates.—The sodic arsenates are prepared like the phosphates, which they also resemble in properties. Sodic arsenate, AsONaO3,I2OH2, is very soluble in water, and is converted by the carbonic anhydride of the air into the monohydric salt. Hydric disodic arsenate, AsOHoNaos, 120 FT2, closely resembles the corresponding phosphate, crystallizing in large efflorescent monoclinic prisms. Like the phosphate, it may be obtained from hot solutions in non-efflorescent crystals with 7 aq. At a red heat it parts with the elements of water, yielding sodicpyrarsenate, which, however, cannot exist in solution, but, in contact with water, at once regenerates hydric disodic arsenate. Dihydric sodic arsenate, AsOMo2Nao,OIl2, obtained by adding arsenic acid to sodic carbonate till the solution no longer precipitates baric chloride, forms large soluble rhombic prisms. Sodic antimonate.—When a solution of dihydric dipotassic pyrantimonate is added to the solution of a sodium salt, a granular crystalline precipitate of dihydric disodic pyran- timonute, 60112, is produced. This salt is insoluble in water. Sodic antimonite.—A solution of antimonious anhydride in caustic soda deposits lus- trous rhombic octahedra of sodic metantimonite, SbONao,3OH2, almost insoluble in cold, sparingly soluble in boiling water. Very concentrated solutions sometimes deposit rhombic prisms of dihydric sodic trimetantimonite, SbsOsllozNao. Sodic borate.—The metaborate, 80Na0,40H2, is prepared by fusing together equal molecules of boric anhydride and sodic carbonate, or by boiling a solution of borax with the necessary quantity of sodic hydrate, evaporating to a syrup and allowing to crystallize over sulphuric acid. It forms large triclinic crystals, readily soluble in water. The solution has an alkaline reaction, and absorbs carbonic anhydride from the air. A metaborate with 2 aq, is obtained in long acicular crystals by fusing the above salt in its water of crystallization and then allowing it to crystal- lize, or by crystallizing in presence of a large excess of sodic hydrate.— Sodic tetraborate [borax), B405Nao2,l()OFI2. This salt occurs in the water of some lakes in Thibet, from which it is obtained by evaporation and crystallization. The natural product, known as tineal, formed at one time the chief supply of this salt; but at present most of the borax of commerce is prepared from the boric acid obtained from the lagoons of Tuscany (p. 191.) The boric acid is either added to a boiling solution of sodic carbonate, or boric acid is heated with half its weight of anhy- drous sodic carbonate, in a reverberatory furnace, and the mass, after cooling, extracted with water. The salt crystallizes in monoclinic prisms, soluble in 14 parts of water at ordinary temperatures and in half their weight of water at 100° C. The solution has an alkaline reaction. At temperatures above 60° C. (140° F.) borax crystallizes from concen- trated solutions in regular octahedra, with 5 aq. (octahedral borax). When heated borax parts with its water of crystallization, inturaescing and forming a porous white mass, which, at a higher temperature, fuses to a clear glass. In a state of fusion, it dissolves metallic oxides, with many of which it yields characteristically colored fluxes. This prop- erty, which depends upon the presence of an excess of boric anhydride in the salt, is utilized in the employment of borax as a blowpipe reagent. It is also used in soldering oxidizable metals, to dissolve the oxide in order to expose clean metallic surfaces. Further applications are : in various metallurgical operations as a flux, in the preparation of enamels, and in fixing colors on porcelain. Sodic silicate, SiONao'i.SOTL, is prepared by dissolving 1 molecule of amorphous silicic anhydride in a solution of 2 molecules of sodic hydrate, evaporating to a syrup, and cooling by means of a freezing mixture, stirring at the same time. The salt, after LITHIUM. 435 being purified by recrystallization, forms large monoclinic crystals, very soluble in water. Both in solution and in the dry state it absorbs carbonic anhydride from the air, undergoing decomposition, with separation of amorphous silicic acid. Soluble soda glass may be obtained in the same manner as the potassium compound. On a large scale it is prepared by heating together 100 parts of quartz sand, 60 parts of anhydrous sodic sulphate, and 15 to 20 parts of charcoal dust. The charcoal, by taking up part of the oxygen of the sulphate, facilitates the decomposition of this salt by the silicic anhydride. Soluble soda glass is employed as a cement, in coating building stone in order to preserve it from decay, and in fixing colors in fresco paintings. The alkaline silicates are important constituents of glass [q.v ). COMPOUNDS OF SODIUM WITH SULPHUR AND HYDROSULPHYL. Sodic sulphide, Sodic polysidphides, and Sodic sulphhydrate.—These compounds are pre- pared like the corresponding potassium compounds, which they closely resemble. Sodic sulpharsenate, 2AsS//Nas3,150H2, is prepared like the potassium compound. It forms large, colorless rnonoclinic prisms, readily soluble iu water. Sodic sulphantimonate {Schlippe’s salt), SbS//Nas3,50H2, is obtained like the potash salt. It crystallizes in large pale yellow tetrahedra, readily soluble in boiling water. When exposed to the air, the crystals undergo superficial decomposition, becoming coated with a reddish-brown layer of antimonic sulphide. SULPHO-SALTS OF SODIUM. COMPOUNDS OF SODIUM WITH NITROGEN AND HYDROGEN. Sodic amide, NNalI2, is formed when sodium is gently warmed in a current of dry gaseous ammonia. The sodium fuses, yielding a greenish-blue liquid, which, on cool- ing, solidifies to a crystalline mass, whilst the color at the same time changes, through brown and olive-green, to a flesh tint. In presence of moisture and under the influence of heat, it behaves like potassic amide (pp. 423, 424). General Properties and Reactions of the Compounds of Sodium.—The salts of sodium are as a rule more soluble than those of potassium. The only insoluble sodium salt is the dihydric disodic pyr- antimonate (5b203H02Na02,60H2) (p. 434). Sodium compounds color the non-luminous flame an intense yellow. The color is invisible when a piece of cobalt glass or a solution of indigo is interposed between the flame and the eye. The flame spectrum of sodium consists of a double line in the yellow, coincident with D in the solar spectrum. LITHIUM, Li2? Atomic weight =7. Probable molecular weight = 14. Sp. gr = 0.59. Fuses at 180° C. (356° F.). Atomicity '. Evidence of atomicity: Lithic chloride, .... LiCl. Lithic hydrate (lithia), .... .... OLiH History.—Lithic hydrate was discovered by Arfvedsou in 1817. The metal was first isolated by Bunsen, Occurrence.—Lithium is a constituent of several rare minerals, such as lepidolite (lithia mica), petalite, spodumene, and triphyline. By the aid of spectrum analysis, lithium has been shown to be very widely dis- 436 INORGANIC CHEMISTRY, tributed: thus it occurs in minute quantities in the ashes of plants and in many mineral waters. Preparation.—Metallic lithium cannot, like potassium or sodium, be reduced from its oxygen compounds by heating with charcoal. It is obtained by the electrolysis of the fused chloride. For this purpose a battery power of five or six Grove’s cells is required. The positive pole is of hard gas coke, the evolved chlorine having no action upon this substance; for the negative pole, an iron wire is employed. A globule of molten metallic lithium soon forms on the iron wire under the surface of the fused chloride. As soon as this globule has attained the size of a pea, it is lifted out of the chloride along with the iron wire by means of a small iron spoon, a coating of lithic chloride protecting it from instantaneous oxidation, and is allowed to cool under petroleum. The globule is then detached from the wire and these operations are repeated until a sufficient quantity of the metal has been obtained. The globule must not be permitted to attain too great a size, otherwise it will detach itself from the iron wire and rise to the surface of the fused chloride, where it generally inflames. Properties.—Lithium is a silver-white metal, harder than potassium or sodium, but softer than lead. It has a sp. gr. of 0.59, and is thus the lightest solid known. It floats on petroleum. It is less oxidizable than potassium or sodium, but speedily tarnishes when exposed to the air. Heated in air to a temperature considerably above its fusing-point, it inflames, burning with an intense white light. It decomposes water, without however inflaming, even when the water is hot. The solution contains lithic hydrate, LiHo. COMPOUNDS OF LITHIUM WITH THE HALOGENS. These compounds are prepared by dissolving the hydrate or carbonate in the corresponding hydracid. Lithic chloride, LiCl, crystallizes in anhydrous octahedra, having the taste of common salt. At temperatures below 10° C. (50° F.) it crystallizes with 2 aq. It is deliquescent and readily soluble in alcohol or in a mixture of alcohol and ether, by which means it may be sepa- rated from the other chlorides of this group. It volatilizes below a red heat. Lithic iodide, LiT,3OH2, forms very deliquescent needles. Lithic fluoride, LiF, crystallizes in small opaque white granular crystals, sparingly soluble in water. COMPOUNDS OF LITHIUM WITH OXYGEN AND HYDROXYL. Lithic oxide, OLi2, is obtained as a white spongy mass, containing a certain quantity of a higher oxide, by burning lithium in dry oxygen. Lithic hydrate (Lithia), LiHo, is prepared like the hydrate of potas- sium, which it also resembles in most of its properties. It is, however, COMPOUNDS OF LITHIUM. 437 less soluble in water than potassic hydrate, and does not deliquesce when exposed to the air. Fused lithic hydrate corrodes platinum power- fully, and should therefore always be prepared in a silver capsule. OXY-SALTS OF LITHIUM. These are for the most part obtained by neutralizing the acid with lithic hydrate or carbonate. Lithic nitrate, N02Lio, crystallizes at 15° C. (59° F.) in anhydrous rhombohedra, below 10° C. (50° F.) in thin prisms of the formula 2N02Li0,50H2. It is deliques- cent and very soluble in water. roci Lithic perchlorate, -j O ,is a deliquescent salt, readily soluble in alcohol, i OLio Lithic carbonate,, COLio2, occurs in small quantities in various mineral waters. It is prepared by precipitating a solution of lithic chloride or nitrate with potassic, sodic, or aramonic carbonate. It is thus obtained as a white crystalline powder, sparingly soluble in cold water. The solution is alkaline and deposits the salt by slow evapora- tion in small prisms. At a bright red heat lithic carbonate undergoes partial decomposition, evolving carbonic anhydride. Owing to its in- solubility, this salt is frequently employed in separating lithium from potassium and sodium. Lithic sulphate, S02Lio2,0H2, forms flat, monoclinic prisms or tables, readily soluble in water, soluble also in alcohol. Potassic lithic sulphate, S306Ko4Lio2.—Hexagonal crystals. Lithic dithionate, | gQ2|Ji^,2oH2, is prepared by exactly precipitating a solution of baric dithionate with lithic sulphate and evaporating the resulting solution of lithic dithionate to crystallization. It forms large rhombic crystals, readily soluble in water and somewhat deliquescent. It is insoluble in alcohol. Lithic phosphate, 2POLio3,01T2, is precipitated, slowly in the cold, instantaneously on heating, when hydric disodic phosphate is added to a solution of a lithium salt. If the solution is rendered alkaline by the addition of sodic hydrate or carbonate, the precipitation of the lithium is complete. Lithic phosphate forms a white crystalline powder, very sparingly soluble in water ( I part of the salt requires 2500 parts of water at ordinary temperatures for solution), still less soluble in water contain- ing ammonia. When heated, it parts with its water of crystallization, but does not fuse, even at a red heat. This salt is employed in the esti- mation of lithium.—Dihydric lithic phosphate, POIIo2Lio, is formed when either the preceding salt, or lithic carbonate, is dissolved in an excess of phosphoric acid and the solution evaporated. It is thus ob- tained in large, very soluble, deliquescent crystals, with an acid reaction. General Properties and Reactions of the Compounds of Lithium.—Lithium is distinguished from the other metals of the alkali group by the sparing solubility of its normal carbonate and phosphate 438 INORGANIC CHEMISTRY. and by the solubility of lithic chloride in a mixture of alcohol and ether. Lithium compounds color the non-luminons flame carmine-red. The spectrum of lithium displays a bright line Li a in the red, and a faint line Li/9 in the yellow. At the temperature of the oxyhydrogen flame a brilliant blue line makes its appearance. RUBIDIUM, Rb2? Atomic weight = 85.3. Probable molecular weight 170,6. Bp. gr. 1.52. Fuses at 38.5° C. (101.3° F.). Atomicity '. Evidence of atomicity : Rubidic chloride, . . . . RbCl. Rubidic iodide, . . . Rbl. Rubidic hydrate, . . . . . RbHo. History.—Rubidium was discovered in 1860 by Bunsen and Kirch- lioff with the aid of spectrum analysis. Occurrence.—This rare metal is widely distributed in nature, but always in very minute quantity. It occurs along with potassium in many minerals (frequently in lepidolite), in the ashes of plants, and in some mineral springs. It was first obtained from the water of a spring at Diirkheim in Baden. Preparation.—l. Metallic rubidium may be obtained by the elec- trolysis of the fused chloride as in the preparation of lithium (p. 436). 2. A more advantageous process consists in distilling a mixture of rubidic carbonate and carbon obtained by charring rubidic tartrate, as in the corresponding method for the preparation of potassium (p. 412). Properties.—Rubidium is a lustrous white metal, with a yellowish tinge. It is soft like wax, even at —lo° C. (14° F.). It fuses at 38.5° C. (101.3° F.), and boils below a red heat, yielding a greenish-blue vapor. Exposed to the air, it instantly becomes covered with a bluish- gray film of oxide and speedily inflames spontaneously. It burns, with vivid incandescence, in chlorine and in the vapors of bromine, iodine, sulphur, and arsenic. In contact with water it behaves like potassium. COMPOUNDS OF RUBIDIUM. Rubidic chloride, RbCI, crystallizes in transparent colorless cubes, possessing a vitreous lustre. It is more soluble than potassic chloride (100 parts of water at 7° C. dissolve 83 parts), and is easily fusible and volatile. It forms double salts with other metallic chlorides. The most important of these double chlorides is rubidic platinic chloride (PtCl4,2RbCl),. which is even less soluble than the corresponding po- tassium compound, and is employed in the separation of rubidium. Rubidic bromide, Rbßr, crystallizes in lustrous cubes with subordinate octahedral faces, and is soluble in its own weight of water at ordinary temperatures. Rubidic iodide, Rbl, resembles the bromide. It dissolves in 0.7 part of water at ordinary temperatures. CAESIUM. 439 Rubidic hydrate, RblTo, resembles the potassium compound, but is a more powerful base. Rubidic nitrate, N02Rbo, forms hexagonal crystals, soluble in 2.3 parts of water at 10° C. (50° F.). Rubidic chlorate, | —This salt forms small prismatic crystals, soluble in 20- 25 parts of water at ordinary temperatures. r oci Rubidic perchlorate, O , forms small hard lustrous rhombic crystals. It is less ( ORbo soluble than the corresponding potassium salt, 1 part of the salt requiring 92 parts of water at 21° C. (70° F.) for its solution. Rubidic carbonate.—The normal salt, COE,bo2,0H2, forms indistinct crystals with a strong alkaline reaction. The water of crystallization is expelled by heating. It is readily soluble in water. Exposed to the air it deliquesces and absorbs carbonic anhydride, forming the acid salt COHolibo, which crystallizes in non-deliquescent prisms with a vitreous lustre. Rubidic sulphate.—The normal salt, S02Rbo2, crystallizes in large, hard, rhombic crystals with a vitreous lustre, more soluble in water than the potassium salt. The acid salt, S02Hoßbo, forms short rhombic prisms. Rubidic dithionate, | orms bard, hexagonal crystals, with a vitreous lustre. Rubidic borate.—A tetraborate of the formula B405Rbo2,60H2 is known. It forms small lustrous crystals belonging to the rhombic system. CiESIUM, Cs2 ? Atomic weight = 133. Probable molecular weight 266. Sp. gr. 1.88. Fuses at 26.5° C. (79.7° F.). Atomicity'. Evidence of atomicity: Cffisic chloride, .... . . . CsCl. Csesic hydrate, .... . . . CsHo. History.—This metal, which is even rarer than rubidium, was dis- covered simultaneously with the latter in the water of the Diirkheira spring by Bunsen and Kirchhoff, in 1860. Occurrence.—The rare mineral pollux, which occurs in the granite of Elba, is a silicate of aluminium, sodium, and caesium, and contains 32 per cent, of the latter metal. In minute traces caesium is found in a variety of minerals, and in many mineral springs. Preparation.—Metallic caesium cannot be obtained by the methods usually employed in the isolation of the alkali metals. Heating the oxide or carbonate with charcoal yields no result; whilst, in the elec- trolysis of the fused chloride, the reduced metal immediately acts upon the undecoraposed chloride, yielding a blue compound of unknown com- position—possibly a subchloride. If, however, fused caesic cyanide, Cs(CN), mixed with a quarter of its weight of baric cyanide, Ba(CN)2, in order to increase the fusibility, be subjected to electrolysis, pure me- tallic caesium is obtained in coherent masses. Properties.—Caesium is a lustrous white metal. At ordinary tem- 440 INORGANIC CHEMISTRY. peratures it is soft. It fuses at 26.5° C. (79.7° F.). When exposed to the air it oxidizes rapidly, and finally inflames spontaneously. Thrown on to water it behaves like potassium. Caesium is the most electro- positive of the elements. COMPOUNDS OF OJEBIUM. CiESic chloride, CsCl, crystallizes in indistinct cubes, which are very soluble and deliquescent. It fuses below redness, and is more easily volatilized than potassic chloride. When heated in moist air it is partially converted into hydrate. It forms double salts with other metallic chlorides. Ccesio antimonious chloride (SbCI3,CsCI) is obtained as a white crystalline precipitate by the addition of antimonious chloride dissolved in hydrochloric acid to a solution of caesic chloride. Ccesio platinic chloride (PtCl4,2CsCl) forms a yellow crystalline precipitate, even less soluble than the corresponding rubidium salt. Coesic hydrate, CsHo, is a caustic, crystalline substance resembling potassic hydrate. Ccesic nitrate, N02Cso, crystallizes in hexagonal prisms, and is less soluble in water than the potassium salt. Ccesic carbonate.—Both the normal and the acid carbonate resemble in almost every respect the rubidium salts. The normal carbonate is soluble in alcohol. Ccesic sulphate.—The normal salt. S02Cso2, forms prismatic crystals very soluble in water. Hydric coesic sulphate, S02IToCso, crystallizes in small rhombic prisms. General Properties and Reactions of the Compounds of Rubidium and Cesium.—The salts of rubidium and csesium cannot be distinguished from those of potassium by the ordinary chemical tests. Like these they yield precipitates with platinic chloride and tar- taric acid. Csesic platinic chloride (PtCl4,2CsCl) is more difficultly soluble in boiling water than rubidic platinic chloride (PtCl42RbCl), and this again is more difficultly soluble than the potassium compound. In this way a separation of the three metals may be effected. Caesium may also be separated from rubidium by the solubility of its normal carbonate in alcohol. The flame colorations of the caesium and rubidium compounds resemble closely that of potassium. By means of the char- acteristic spectra, however, the compounds of the three metals may be readily distinguished. The spectrum of rubidium consists of two lines, Rba and Rb/3, in the violet, and two lines, Rb«s and Rby, in the red, together with other fainter lines. The most characteristic lines in the spectrum of csesium are Cs« and Cs/3 in the blue. THE AMMONIUM SALTS. i NPT The hypothetical radical ammonium < jq-fj4 has already been referred , ( v. A 4 to (p. 235) in connection with the compounds of nitrogen. Its salts closely resemble those of the alkalies, and may therefore be appropri- ately treated of at this point. THE AMMONIUM SALTS. 441 COMPOUNDS OF AMMONIUM WITH THE HALOGENS. Ammonic chloride, NH4Cl.—This compound occurs in small quantities in the neighborhood of volcanoes, being generally formed when lava flows over fertile land. The nitrogenous vegetable mater, thus subjected to a destructive distillation, furnishes ammonia, the lat- ter combining with the hydrochloric acid which is almost always pres- ent in volcanic gases. Ammonic chloride is prepared by neutralizing the ammoniacal liquor from the gas-works—the ammonia being in this case a product of the destructive distillation of fossil vegetable matter— with hydrochloric acid, and purifying the crude ammonic chloride by crystallization and sublimation. The aqueous portion of the distillate obtained in the preparation of animal charcoal from bones is also very rich in ammonia, and serves as a source of the chloride. Ammonic chloride crystallizes from water in small indistinct octahedra or cubes, which are generally grouped in fern-shaped aggregations. When heated, it does not fuse, but sublimes, undergoing dissociation into ammonia and hydrochloric acid, which again unite as the temperature falls. When sublimed in large quantities, it forms semi-transparent, tough, fibrous masses. Dissociation also takes place when a neutral solu- tion of the salt is boiled: a small quantity of ammonia passes off with the steam, and free hydrochloric acid is found in the solution. In pres- ence of an excess of hydrochloric acid this dissociation does not occur, and solutions of ammonic chloride may be evaporated at 100° C. with- out loss, Ammonic chloride is soluble in 2J parts of water at ordinary temperatures and in its own weight of water at 100° C. Absolute alcohol does not dissolve it. Ammonic chloride forms double salts with various metallic chlorides: ammonic platinic chloride, PtCl4,2N1T4CI, crystallizes in minute octahedra, almost insoluble in water, and insol- uble in a mixture of alcohol and ether. This double salt, which closely resembles the corresponding potassium compound, is employed in the quantitative determination of ammonia. When heated, the double salt is decomposed, platinum being left behind in the finely divided con- dition in which it is known as spongy platinum. Ammonic chloride has numerous uses. It is employed in medicine, in dyeing, in solder- ing, and tinning—in which last process it serves to produce a clean metallic surface, either by reducing the oxides at a high temperature, or by converting them into fusible chlorides—in the preparation of ammonia and ammonic carbonate, as a laboratory reagent, and as a manure. Ammonic bromide, NH4Br.—This compound is prepared by the direct union of hydrobroraic acid with ammonia, or by the addition of bromine to aqueous ammonia, nitrogen being evolved in the latter reaction: 4NH3 -f 3Br = 3NH4Br + N. Ammonia. Ammonic bromide. It crystallizes in colorless cubes, readily soluble in water, less soluble in alcohol. The crystals become moist in contact with the air, and 442 INORGANIC CHEMISTRY. assume a yellow color, owing to the separation of bromine. It sub- limes without fusing. Ammonia iodide, NH4l.—This salt is prepared by the direct union of ammonia and hydriodic acid, or more conveniently by adding to a hot saturated solution of potassic iodide the equivalent quantity of ammonic sulphate, precipitating the potassic sulphate with alcohol, and evapo- rating the solution. It crystallizes in colorless cubes, readily soluble in water and in alcohol. It may be sublimed in an atmosphere free from oxygen. Exposed to the air, it assumes a yellow color, due to the liberation of iodine. Ammonic iodide is employed in photography. Ammonic fluoride, NH4F, is obtained by evaporating a solution of hydrofluoric acid supersaturated with ammonia and kept alkaline with ammonia during the evaporation, or by heating in a platinum vessel a mixture of 1 part of ammonic chloride with parts of sodic fluoride, when the ammonic fluoride sublimes and condenses in crystals on the cooled lid of the vessel. It crystallizes in colorless hexagonal prisms or laminae, deliquescent in moist air, readily soluble in water, sparingly soluble in alcohol. On evaporation, the neutral aqueous solution gives off ammonia and yields rhombic prisms of hydric ammonic fluoride, NH4F,11F. Dry ammonic fluoride absorbs gaseous ammonia, which it again parts with on heating. The dry salt decomposes silicates when heated with them. Ammonic fluoride is employed in etching glass. Ammonic silicofluoride, SiF6(NH4)2, is readily soluble in water. COMPOUND WITH HYDROXYL. Ammonic hydrate, NH4Ho.—This compound has not been iso- lated, but may be considered to exist in the aqueous solution of am- monia, which is powerfully alkaline, slightly caustic, and possesses the other properties of the solutions of the alkaline hydrates.* On evapo- ration the ammonic hydrate undergoes dissociation into ammonia and water : NH4Ho ='NH3 + OII2. (For the other properties of aqueous ammonia, see p. 232.) Ammonic oxide, 0(NVH4)2, is unknown. OXY-SALTS OF AMMONIUM. These are, as a rule, prepared by neutralizing aqueous ammonia or am- monic carbonate with the oxy-acid. Special methods will be described under the corresponding salts. Ammonic nitrate, N02(NvH40), or N02Amo, forms six-sided prisms belonging to the rhombic system. It dissolves in about half its weight of water at 18° C. (64° F.), with great absorption of heat. * Kohlransch, however, finds that, whereas the ammonic salts, when in solution, possess the same electrolytic conductivity as the corresponding potassium salts, aqueous ammonia is a bad conductor of the current, whilst a solution of potassic hydrate con- ducts the current well. From this he concludes that an aqueous solution of ammonia contains little or no ammonic hydrate. THE AMMONIUM SALTS. 443 In moist air it deliquesces, at the same time losing ammonia and be- coming acid. When heated, it is decomposed into nitrous oxide and water (p. 220). At low temperatures it absorbs gaseous ammonia with great avidity, taking up at —lo° C. (14° F.) two molecules of am- monia, and yielding a compound of the formula N(NH2)2Ho2Amo. This substance is a colorless liquid of sp. gr. 1.05, which does not solidify at lB° C. (0° F.). As the temperature rises this compound dissociates, till at 28.5° C. (83.3° F.) it parts with one molecule of am- monia, and is converted into a white crystalline mass, of the formula NO(NH2)HoAmo. This substance also suffers dissociation as the temperature rises, giving off ammonia and yielding at 80° C. (176° F.) pure arnmonic nitrate. Ammonic nitrite, NOAmo, is formed in small quantity when phosphorus undergoes slow oxidation in contact with moist air; also during the combustion of hydrogen or hydrogenous substances in air, and by the action of ozone on dilute ammonia. It may be obtained in crystals by passing simultaneously ammonia, nitric oxide, and oxygen into a dry flask. It is most easily prepared by the double decomposi- tion of argentic nitrite with ammonic chloride, or of baric nitrite with ammonic sulphate, the solution obtained by either of these methods being filtered from the insoluble precipitate and evaporated in a desic- cator over quickline. Thus obtained it forms a crystalline, very soluble mass. It decomposes slowly at ordinary temperatures into nitrogen and water (p. 212). When heated to 60-70° C. (140-158° F.), or when struck, it detonates, In concentrated aqueous solution it under- goes rapid decomposition, the process being accelerated by heat and retarded by dilution. Ammonic chlorate, | Q^mo> is prepared by neutralizing chloric acid with ammonia or ammonic carbonate, or by the double decomposition of ammonic silicofluoride with potassic chlorate, filtering from the insoluble potassic silicofluoride and evaporating over sulphuric acid. It crystallizes in colorless prisms or slender needles, readily soluble in water or alcohol. When dry the crystals turn yellow and frequently explode spontaneously with great violence. This explosive decomposition takes place at once on heating to somewhat above 100° C. The aqueous solution on boiling evolves nitrogen and chlorine. roci Ammonic perehl&rate, 1 O .—Large rhombic crystals, soluble in 5 parts of ( OAmo water. f OBr Ammonic bromate, ■< m0> forms white needles or crystalline granules. The dry salt explodes spontaneously like the chlorate. Ammonic iodate, | q \mo‘—Lustrous quadratic crystals, soluble in 38 parts of water at ordinary temperatures and in 6.9 parts of boiling water. At 150° C. (302° F.) it decomposes with a hissing noise, yielding equal volumes of oxygen and nitrogen, together with iodine and water. Ammonic carbonate : Normal ammonic carbonate, COAmo2.—This salt is deposited as a crystalline powder when a concentrated solution of the sesquicarbonate {vide infra) is saturated with gaseous ammonia, and in large tabular crystals when a hot solution of the sesquicarbonate in dilute aqueous 444 INORGANIC CHEMISTRY. ammonia is allowed to cool. It is a very unstable salt. When exposed to the air it rapidly parts with ammonia and is converted into hydric ammonic carbonate, COHoAmo. It dissociates completely at 58° C. (136° F.) into carbonic anhydride, ammonia, and water. It is soluble at ordinary temperatures in its own weight of water, but only spar- ingly soluble in concentrated ammonia.—Hydric ammonic carbonate, COHoAmo, occurs in a crystallized form in guano beds. It may be obtained from the commercial sesquicarbonate either by exposing the latter salt to the air, when it parts with ammonia, yielding the acid carbonate; or by treating the sesquicarbonate with a small quantity of water, which dissolves the normal carbonate, leaving the acid carbonate. It is also deposited when a concentrated solution of the sesquicarbonate is exposed to a low temperature, or is mixed with alcohol, or is satu- rated with carbonic anhydride. It crystallizes in hard lustrous rhombic prisms. It sublimes at 60-65° C. (140-149° F.). It dissolves in about 8 parts of water at ordinary temperatures. The solution slowly evolves carbonic anhydride and becomes ammoniacal. This decompo- sition is very rapid above 36° C. (97° F.), the liquid effervescing when warmed. It is insoluble in alcohol, but on long standing under alcohol dissolves as normal carbonate, with evolution of carbonic anhydride. Ammonic sesquicarbonate, COAmo2,2COHoAmo.—This salt is pre- pared on a large scale by heating ammonic chloride or sulphate with calcic carbonate, when the sesquicarbonate sublimes. It forms a trans- lucent, crystalline mass, which is usually coated with an opaque layer of the acid carbonate. Its composition varies, generally approximating however to the above formula. It has an ammoniacal odor, and is gradually converted by exposure to air into the acid salt. Ammonic sulphate : Ammonic sulphate, S02Amo2, is found native as mascagnine. It is prepared on a large scale by passing the ammonia from the ammoniacal liquors of the gasworks into sulphuric acid. It forms colorless rhom- bic crystals, isoraorphous with the potassium salt. It is soluble in twice its weight of cold, in its own weight of boiling, water ; insoluble in alcohol. It fuses at 140° C. (284° F.), and above 280° C. (536° F.) is decomposed into ammonia, nitrogen, water, and ammonic sulphite, the latter subliming.—Hydric ammonic sulphate, S02lloAmo, crystallizes from a solution of the normal salt in concentrated sulphuric acid in deliquescent thin rhombic crystals. It is soluble in its own weight of cold water, also in alcohol. Ammonic sulphate is employed in the manufacture of ammonia- alum ; also as a manure. Ammonic potassic sulphate, SO2 AmoKo, is obtained by evaporating a solution of mo- lecular quantities of ammonic potassic sulphates. It crystallizes in lustrous scales. Ammonic sodic sulphate, S0lAinoNao,2oll2, is prepared like the foregoing. It is also deposited in crystals when mixed solutions of sodic sulphate and ammonic cliloride, or of sodic chloride and ammonic sulphate, are evaporated. The salt is permanent in air. Ammonic sulphite, SOAmo2IOFT2, is obtained by neutralizing an aqueous solution of sulphurous anhydride with ammonia, and then adding alcohol. The salt separates in monoclinic crystals, readily soluble in water. By exposure to the air it is oxidized to THE AMMONIUM SALTS. 445 amnionic sulphate. When a solution of this salt is saturated with sulphurous anhy- dride and evaporated over sulphuric acid, it deposits crystals, not of the acid salt, but rSOAmo of ammonia pyrosulphite, -I O . This salt evolves sulphurous anhydride when (SOAmo exposed to the air, at the same time undergoing oxidation to amnionic sulphate. Ammonia dithionate, gjQ2^mo’OH2, is obtained by the double decomposition of baric dithionate with amnionic sulphate. It forms colorless capillary crystals, very soluble in water, insoluble in alcohol. Ammonia thiosulphate (Ammonia hyposulphite), is prepared by decomposing calcic thiosulphate with amnionic carbonate. It forms deliquescent, very soluble acicular crystals or rhombic plates. Ammonic phosphate : a. Ammonic phosphate, POAmo3,30H2, occurs sometimes in guano. It is formed when a concentrated solution of hydric diammonic phos- phate is mixed with ammonia, and is deposited in small prismatic or acicular crystals, which when exposed to the air part with ammonia, yielding hydric diammonic phosphate. When boiled for some time in aqueous solution, it is converted into dihydric ammonic phosphate. When strongly heated, it yields, like all the other ammonic'phosphates, metaphosphoric acid.—Hydric diammonic phosphate, POHoAmo2, is prepared by evaporating an ammoniacal solution of phosphoric acid, care being taken to keep the ammonia slightly in excess during the process. It forms large colorless monoclinic crystals, soluble in 4 parts of cold, more readily in boiling, water; insoluble in alcohol. Exposed to the air it gradually parts with ammonia.—Dihydric ammonic phos- phate, POHo2Amo, is prepared by adding phosphoric acid to ammonia till the solution is strongly acid and no longer precipitates baric chloride, and evaporating to the crystallizing point; or by boiling a solution of the monohydric salt. It crystallizes in quadratic octahedra, which are permanent in air. It is somewhat less soluble than the foregoing salt. Hydric ammonic sodic phosphate {Microcosmic salt), This salt occurs in guano and in putrid urine. It is prepared by dis- solving 6 parts of hydric disodic phosphate and 1 part of ammonic chloride in 2 parts of boiling water, and allowing the liquid to cool. It forms large colorless raonoclinic prisms, very soluble in water, yield- ing a solution which gives off a portion of its ammonia on evaporation. It fuses easily, water and ammonia being expelled, and sodic meta- phosphate left. Microcosmic salt is employed in the laboratory as a blowpipe reagent, the sodic metaphosphate, which remains on heating it, possessing the property of dissolving various metallic oxides at a high temperature to yield characteristically colored fluxes or glasses. POHoAmoNao,4OH2.— Diammonic sodic phosphate, POAmo2Nao,4OH2, separates in lustrous white pearly laminae when strong ammonia is added to a cold saturated solution of the foregoing. It evolves ammonia when exposed to the air, and is converted into hydric ammonic sodic phosphate. h. Ammonic pyrophosphate, P203Amo4, separates in small acicular laminse when alcohol is added to a solution of pyrophosphoric acid 446 INORGANIC CHEMISTRY. supersaturated with ammonia. Its solution gives off ammonia when boiled, yielding dihydric diammonic 'pyrophosphate, P203Ho2Amo2, which may be precipitated from its solution by the addition of alcohol as a syrupy mass, becoming crystalline on standing. c. Amnionic metaphosphates are also known. Ammonic borate.—The normal salt has not been prepared. Diammonic tetraborate, B40sArno2,40H2, crystallizes from a solution of boric acid in warm concentrated am- monia in quadratic crystals, which give off ammonia when exposed to the air. When this salt is dissolved in water and the solution evaporated by heat, colorless trans- parent rhombic crystals of hydric ammonic tetraborate, B405H0Am0,30H2, are deposited on cooling. COMPOUNDS OF AMMONIUM WITH SULPHUR AND HYDRO- SULPHYL. Ammonic sulphide, SAmz, is obtained in lustrous crystals by the union of 2 volumes of ammonia with 1 volume of sulphuretted hydrogen at a temperature of—18° C. (0° F.). Above this temperature it dissociates, evolving amnonia, and yielding am- monic sulphhydrate, AmHs. Ammonic sulphhydrate, AmHs, is formed by the direct union of equal volumes of ammonia and sulphuretted hydrogen at ordinary tem- peratures. It is best prepared by passing sulphuretted hydrogen into alcoholic ammonia, when the sulphhydrate separates out in a crystalline form. The aqueous solution employed as a laboratory reagent is ob- tained by saturating aqueous ammonia with sulphuretted hydrogen. Ammonic sulphhydrate forms large colorless laminae. It volatilizes readily, with dissociation into ammonia and sulphuretted hydrogen, which reunite on cooling. It becomes yellow, both in the solid state and in solution, when exposed to the air, owing to the formation of polysulphides of ammonium. The solution precipitates many metals in the form of sulphides from the solution of their salts, and dissolves sulphur to form ammonic polysulphides. Ammonic pentasulphide, ArmSs, is prepared by alternately passing ammonia and sul- phuretted hydrogen into a mixture of ammonic sulphhydrate and flowers of sulphur until the liquid solidifies on cooling. The mixture is then heated to 50° C. (122° F.) and allowed to cool with exclusion of air, when the pentasulphide is deposited in orange-yellow rhombic prisms. Water decomposes them with precipitation of plastic sulphur. Ammonic heptasulphide, Am287, is formed by the spontaneous decomposition of the foregoing compound in presence of air: 3(NH4)255 = 2(NH4)257 + NHg + NH4Hs. Ammonic Ammonic Ammonia. Ammonic pentasulphide. heptasulphide. sulphhydrate. It forms ruby-red crystals, which are not decomposed by heat below 300° C. (572° F.), but are slowly decomposed by water. General Properties and Reactions of the Ammonium Salts.—The ammonium salts are all volatile—some with decom- position, others with dissociation, in which last case the dissociated SILVER. 447 constituents recombine on cooling to form the original salt, as in the case of ammonic chloride (p. 64). Ammonium salts yield with pfatinic chloride and with tartaric acid precipitates closely resembling those ob- tained with potassium salts; ammonic platinic chloride (PtCI4,2NH4CI), however, leaves only a residue of spongy platinum on ignition. All ammonium salts, when warmed with calcic hydrate, or with concen- trated caustic potash or caustic soda, evolve gaseous ammonia, which may generally be recognized by its characteristic smell, or in case the quantity is very minute, by the white fumes of ammonic chloride which are formed when a glass rod moistened with hydrochloric acid is held over the mixture. The smallest trace of ammonia in aqueous solution may be detected by means of a solution of mercuric iodide in a mixture of potassic iodide and caustic potash (Nessler’s reagent), with which it yields a brown coloration, or, if present in larger quantity, a brown precipitate, of NHg"(Hg"Ho)HI. This reaction does not occur in presence of alkaline sulphides or cyanides. MONAD METALS. Section IV. SILVER, Ag2? Atomic weight 107.7. Probable molecular weight = 215.4. Sp. gr. 10.57. Fuses at 1040° C. (1904° F.). Atomicity '. Evidence of atomicity : Argentic chloride, AgCl. Argentic iodide, Agl. Argentic oxide, °Ag2. History.—This metal has been known from the earliest times. Occurrence.—Silver occurs native, occasionally in large masses. Native silver is rarely pure: it contains gold, copper, and other metals. In combination, silver occurs as argentic sulphide in silver glance (SAg2); as sulphantimonite in pyrargarite or dark-red silver ore (SbAgs3); as chloride in kerargyrite or horn-silver (AgCl). The bromide, iodide, telluride, antiraonide, and arsenide are rare minerals. Galena, or plumbic sulphide, the commonest form of lead ore, generally contains small quan- tities of silver. Silver also occurs in minute traces in sea-water. Extraction.—Although silver is very readily reducible from its com- pounds (the mere application of heat being generally sufficient for this purpose), yet the extraction of silver from its ores is a matter of con- siderable practical difficulty. The ores of silver are frequently mixed with earthy impurities, from which they cannot be mechanically sepa- rated, or they occur along with the ores of other metals, which are apt to undergo reduction at the same time, and thus contaminate the pro- duct. The process of extraction varies with the nature of the ore ; but the methods employed may be divided into three classes according 448 INORGANIC CHEMISTRY. as they depend upon cupellation, upon amalgamation, or upon reactions in the wet way. а. Cupellation Process.—This process is employed in separating silver from lead. The alloy of silver and lead, obtained from argen- tiferous lead ores, is fused in a reverberatory furnace, the hearth of which is composed of burnt clay. Over the molten metal, which rests upon the concave surface of this hearth or cupel, a rapid current of air is blown. The lead is thus oxidized, and the fused oxide escapes by flowing off through lateral openings in the hearth, whilst the silver remains in the cupel. At first the fused oxide flows off in large quan- tity, but towards the end of the operation it forms thin films upon the surface of the silver, exhibiting the colors of Newton’s rings. At last, as the film of oxide finally disappears, the bright surface of the silver is perceived. This phenomenon is known as the “ figuration” of the metal. The removal of the oxide is aided by skimming. When the lead is sufficiently rich in silver, it is cupelled at once; but if the silver is present in a proportion less than one-tenth of a per cent., the lead is subjected to a preliminary process, which has for its object the concentration of the silver in a relatively small portion of the lead. In this process, invented by Pattinson, the metal is fused in iron pots and allowed to cool slowly. As soon as the temperature has sufficiently fallen, crystals of pure lead are formed; these are constantly removed by means of perforated ladles, and this is continued until the lead in the pot has been reduced in quantity by about two-thirds. In this way the greater part of the silver is left behind in the pot, and by systematic recrystallization, pure and nearly desilverized lead on the one hand, and a lead very rich in silver on the other, may be obtained. The rich lead is cupelled as above described. Instead of treating the lead by Pattinson’s process, it may be fused, and zinc, in the proportion of 11.2 lbs. for every 7 ozs. of silver present per ton of lead, added. The whole is thoroughly stirred and then allowed to settle. The zinc extracts the greater part of the silver from the lead and rises to the surface, where it solidifies first, and may be removed as a solid cake. This cake is then heated to redness in a cur- rent of air, by which means the zinc is converted into zincic oxide, and may be separated from the unaltered silver by washing. Sometimes poor silver ores are roasted along with galena. The lead thus obtained contains the whole of the silver, which may then be separated by cupellation. б. Amalgamation Process.—The amalgamation process formerly em- ployed in Europe was conducted as follows: The finely-ground ore was mixed with common salt and roasted in a reverberatory furnace. By this means the silver, which was mostly present in the form of sulphide, was converted into chloride. The roasted ore was again ground very fine and then introduced, along with scrap iron and water, into casks which were made to revolve by machinery. The chloride was thus reduced to metallic silver: 2AgCI -f~ Fe FeCI3 -(- Ag2. Argentic Ferrous chloride. chloride. SILVER. 449 Mercury was then introduced into the revolving casks. The mercury combined with the silver to form a liquid amalgam, which was sepa- rated and subjected to distillation, when the mercury passed over and the silver remained in the retort. A modification of this proces is employed in Nevada. Some trouble is occasioned in this process by the tendency of the mercury to form minute globules, which, along with the silver contained in them, are lost in washing. This “flouring,” or “sickening,” as it is termed, which is due to the formation of a film of mercuric sulphide, may be prevented by the addition of about 2 per cent, of sodium to the mercury, the mercuric sulphide being thus re- duced to metallic mercury, with formation of sodic sulphide. The method of amalgamation employed in Mexico differs from the above, the scarcity of fuel in the silver-producing districts precluding the application of the roasting process. The ore is first ground very fine with water in a mill. The paste thus obtained is spread on a paved floor, and mixed with a small quantity of common salt, after which it is allowed to stand for a day. About 1 percent, of a substance known to the miners as magistral—a mixture of crude ferric and cupric sulphates obtained by roasting copper pyrites—is added, and the whole is again thoroughly mixed. Mercury is now poured in, and the mixing is renewed. All these processes of incorporation are effected by the treading of blindfolded mules. The mercury is added in successive portions, at intervals of some days, during the working of the heap, the entire quantity of mercury employed being about six times the weight of the silver contained in the ore. The time required for working a heap varies from a fortnight to two months. At the end of the time the liquid amalgam, which contains all the silver, is separated from the earthy and other impurities by washing, and, after pressing in sacks to free it from the excess of mercury, is subjected to distillation. The nature of the chemical changes which occur in the Mexican pro- cess is not thoroughly understood, but the action is supposed to take place as follows: The cupric sulphate undergoes double decomposition with the sodic chloride, yielding sodic sulphate and cupric chloride. The latter salt reacts with the argentic sulphide, converting it into argentic chloride: 2CuC12 + SAg2 = 'Ou'2CI2 + 2AgCI + S. Cupric Argentic Cuprous Argentic chloride. sulphide. chloride. chloride. The cuprous chloride thus formed dissolves in the sodic chloride pres- ent, and is thus enabled to act upon a fresh quantity of argentic sul- phide : 'Cu'2Cl2 + SAg2 = 'Cu'2S" + 2AgCI. Cuprous Argentic Cuprous Argentic chloride. sulphide. sulphide. chloride. The silver chloride held in solution by the sodic chloride is reduced by the metallic mercury, with formation of mercurous chloride: 2AgCI + 2IIg = 'Hg'2Cl2 + Ag2. Argentic Mercurous chloride. chloride. 450 INORGANIC CHEMISTRY. The whole of this mercurous chloride is lost in washing, representing a loss of mercury equal to nearly twice the weight of the silver ob- tained. c. Extraction in the Wet Way.—When argentiferous copper pyrites is roasted, the sulphides of iron, copper, and silver take up oxygen, and are converted into sulphates. By carefully regulating the temperature, a point may be reached at which the sulphates of iron and copper are decomposed, yielding insoluble oxides, whilst the more stable argentic sulphate remains unaltered, and may be obtained in solution afterwards by lixiviating the roasted mass with hot water. A small quantity of undecomposed copper salt goes into solution at the same time. The silver is precipitated from the solution by metallic copper. (Ziervogel.) Another method consists in roasting the ore with common salt, so as to convert the silver into chloride, which is then extracted with a cold dilute solution of sodic thiosulphate. From this solution the silver is precipitated as sulphide by sodic sulphide. The argentic sulphide is reduced to metal by heating to a high temperature in a current of air. (Percy-Patera.) The burnt pyrites obtained in the manufacture of sulphuric acid contains, in addition to copper, a small quantity of silver, amounting to about half an ounce to the ton. This small quantity may be profitably separated by adding to the tank-liquor obtained in the extraction of the copper (see Copper) a solution of kelp. In this way the silver, which is present in the tank-liquor in the form of chloride, and is kept in solution by the sodic chloride with which the burnt pyrites has been roasted, is precipitated as argentic iodide. A trace of gold, which is precipitated at the same time, is afterwards separated. Preparation of Pure Silver.—In order to obtain pure silver advan- tage is taken of the insolubility of the chloride. Ordinary silver is dissolved in dilute nitric acid, when gold, if present, remains undis- solved. The silver is precipitated from the filtered solution as chloride by hydrochloric acid. The washed and dried chloride is fused in a crucible with an excess of sodic carbonate. The silver collects as a regains at the bottom of the crucible. Another method is to reduce the argentic chloride by laying it on a plate of zinc under dilute hydro- chloric acid. The reduced silver is carefully washed with hydrochloric acid to free it from adhering traces of zinc, and is then dried. By this means it is obtained as a fine gray powder, devoid of metallic lustre. In this last form it is known as “molecular” silver (a misnomer, as it is very far from being in a state of molecular subdivision) and is used in organic research for acting upon organic compounds of the halogens. Properties.—Silver has a white color, with a tinge of yellow, and pos- sesses great lustre when polished. In the form in which it is obtained by the ignition of some organic silver salts, it is white like porcelain, owing to the roughness of its surface, and the consequent absence of metallic lustre. Of all the metals it is the best conductor of heat and electricity. It is a soft metal, standing between copper and gold in hardness. In malleability and ductility it is inferior only to gold ; it can be beaten into leaf 0.00025 mm. in thickness, andean be drawn into wire of which 180 metres weigh 0.1 gram. In very thin films, as in the case when it SILVER 451 is deposited upon glass from ammoniacal solutions by means of reducing agents, it transmits blue light. It possesses great tenacity; the break- ing weight for a wire 2 millimetres in diameter being 85 kilograms. Its specific gravity is 10.57. It crystallizes in regular octahedra. Na- tive silver also occurs in dendritic forms. It fuses at 1040° C. (1904° F.), and may be distilled at a white heat by means of the oxyhydrogen blowpipe, a process which was employed by Stas in purifying silver for the purpose of determining its atomic weight. When melted in contact with air, pure silver absorbs about 22 times its volume of oxygen, which it again gives up at the moment of solidification. As the metal cools, the outer crust solidifies first, and the gas evolved from the interior then escapes through this crust in sudden bursts, carrying with it small particles of molten silver. This phenomenon is known as the “spit- ting” of silver. The presence of a small quantity of copper prevents the absorption of oxygen. Pure air, oxygen, and water are without action upon silver at all temperatures, but ozone oxidizes it superficially to peroxide. Reactions.—l. Silver is blackened by sulphuretted hydrogen in pres- ence of oxygen, argentic sulphide being formed. For this reason silver articles exposed to the atmosphere become discolored. Pure sulphuret- ted hydrogen, however, is without action upon silver at ordinary tem- peratures, and the metal may even be heated with an aqueous solution of sulphuretted hydrogen to 200° C. (392° F.) without blackening. 2. Silver is acted upon by hot concentrated sulphuric acid : Ag2 + 2S02Ho2 S02Ago2 20H2 -(- S02. - Sulphuric Argentic Water. Sulphurous acid. sulphate. anhydride. 3. Dilute nitric acid readily dissolves silver 3Ag2 + 8N02Ho = 6N02Ago + 40H2 + 2'N"O. Nitric Argentic Water. Nitric acid. nitrate. . oxide. 4. At a red heat silver decomposes hydrochloric acid : Ag2 2HCI H2 + 2 AgCI. Hydrochloric Argentic acid. chloride. Strong aqueous hydriodic acid dissolves silver, even at ordinary tem- peratures, with evolution of hydrogen and formation of argentic iodide. Uses.—Pure silver is very little employed in the arts, as it is too soft to resist wrear. In order to increase its hardness and tenacity, it is alloyed with a small proportion of copper, an addition which does not affect its color, and in this form it is employed for plate, ornaments, coinage, etc. Standard silver is an alloy of silver and copper of a given composition fixed by law, and this standard varies in different countries. In England the standard contains 92.5 per cent, of silver. In France, Germany, and Austria, the standard for coinage contains 90 per cent. 452 INORGANIC CHEMISTRY. of silver, whilst there are other standards for plate and jeweller’s work. What is termed the fineness of silver is the number of parts of silver per mille which the alloy contains; thus the English standard silver has a fineness of 925. Pure silver is employed in the manufacture of various laboratory vessels; this metal, unlike glass and platinum, being capable of resist- ing the action of fused caustic alkalies. Silver is also employed in electroplating. For this purpose the ob- ject to be silvered, which must possess a conducting surface, is made the negative electrode; the positive electrode consists of a plate of silver. The electrodes are immersed in a solution of argentic cyanide in an ex- cess of potassic cyanide. The electrolytic silver is deposited as a coherent coating on the object to be silvered, and the cyanogen, liberated at the negative electrode, combines with the silver of the electrode to form argentic cyanide, which dissolves in the excess of potassic cyanide, so that the strength of the electrolytic solution remains constant. From silver solutions other than the above, the electrolytic silver is generally deposited in the form of a non-coherent powder. The silvering of glass is effected by means of a mixture of an ar»- moniacal solution of silver with milk-sugar, or some other suitable organic reducing agent. The solution is contained in a flat shallow vessel, and the glass is suspended so that the surface to be silvered, which must previously have been thoroughly cleaned, may be in contact with the surface of the liquid. A bright coherent mirror of silver is thus deposited on the glass. Reflectors for astronomical telescopes are now extensively prepared by this method. COMPOUNDS OF SILVER WITH THE HALOGENS. Argentic chloride, AgCl, occurs native, as kerargyrite, or horn- silver, in Mexico, Peru, and Chili, also in the Harz, Horn-silver crys- tallizes in forms belonging to the regular system, but more frequently occurs in wax-like, translucent masses. Its specific gravity varies from 5.3 to 5.4. Argentic chloride is obtained as a curdy precipitate by the addition of hydrochloric acid, or a soluble chloride, to the solution of a silver salt. When pure it is white; but under the influence of light it speedily assumes a violet tint, passing into black. The reason of this phenomenon, which is turned to account in photography, is not thor- oughly understood, but the change is supposed to be due to the forma- tion of a lower chloride, or to the liberation of metallic silver. The action is only superficial, and the quantity of chlorine evolved extremely small. Argentic chloride fuses at about 260° C. (500° F.) to a clear, yellow liquid, which solidifies to a translucent, horny, sectile mass. It is insoluble in water and dilute acids; slightly soluble in concentrated hydrochloric acid, and in concentrated solutions of the alkaline chlorides; readily soluble in ammonia, potassic cyanide, sodic thiosulphate, and in a concentrated solution of mercuric nitrate. On evaporation, the solu- tions in hydrochloric acid and in ammonia deposit the agentic chloride in octahedra. In contact with oxidizable metals, such as iron or zinc, COMPOUNDS OF SILVER. it is reduced, in presence of water, to metallic silver, the addition of a little acid favoring the reaction. The dry chloride absorbs gaseous ammonia to form the compound 2AgCI,3NH3, which parts with its ammonia at 37.7° C. (100° F.), and was employed by Faraday in the liquefaction of ammonia (p. 231). This compound is also obtained in large transparent rhombohedra, when a solution of argentic chloride in concentrated ammonia is allowed to stand in an imperfectly closed bottle. Argentous chloride, Ag4Cl2, is obtained by treating argentous oxide (q v.) with hydro- chloric acid. It forms a black powder, which is decomposed by ammonia into metallic silver and argentic chloride, the latter dissolving in the ammonia. Nitric acid decom- poses it in a similar manner, the silver in this case dissolving, whilst the chloride is left. Argentic bromide, Agßr, occurs native as hromargyrite in Mexico and Chili, also at Huelgoet in Britanny. It generally forms concretions, but is also found crystallized. It may be prepared by precipitating solutions of silver salts with hydrobromic acid. At ordinary tempera- tures, hydrobromicacid converts argentic chloride into argentic bromide; at 700° C. (1292° F.), on the other hand, this reaction is reversed, and the bromide is converted by hydrochloric acid into chloride. Precipi- tated argentic bromide is a faint yellow substance, soluble with difficulty in dilute ammonia, readily soluble in concentrated ammonia. The dry bromide does not absorb ammonia; but a double compound with am- monia, corresponding to that of the chloride, is deposited from the am- moniacal solution. Argentic bromide fuses below a red heat. It is employed in photography in the preparation of “ dry plates.” Argentic iodide, Agl, is of very rare occurrence. It is found as iodargyrite, in Chili, Mexico, and Spain, in the form of thin hexagonal plates which are slightly elastic. It is obtained as an amorphous yellow precipitate when potassic iodide is added to the solution of a silver salt. Concentrated hydriodic acid dissolves metallic silver with evolution of hydrogen ; from this solution lustrous larainse of the formula Agl,HI are deposited on cooling; and these, on exposure to the air, are speedily decomposed, yielding argentic iodide. When the mother liquor from these crystals is exposed to the air, or when it is left in contact with excess of metallic silver, it deposits argentic iodide in hexagonal prisms. Argentic chloride and bromide are converted by hydriodic acid with violent reaction into the iodide ; but above 700° C. (1292° F.) gaseous hydrochloric acide converts the iodide into chloride. Argentic iodide closely resembles the chloride and bromide, but differs from these in its almost perfect insolubility in concentrated ammonia, which, however, has the effect of turning it white. It is soluble in sodie thiosulphate, though not so readily as the chloride. It also dissolves in a concen- trated solution of potassic iodide, the hot solution depositing on cooling acicular crystals of the formula Agl,HI; from this solution the iodide is precipitated by dilution with water. It fuses at a dull red heat, yielding a yellow liquid which becomes darker colored at a higher tem- perature, and on cooling solidifies to a yellow mass with a sp. gr. of 6.687. The sp. gr. of the precipitated iodide is 5.807, that of the crys- tallized variety 5.47-5.54. Fizeau has made the remarkable observa- 454 INORGANIC CHEMISTRY. tion that between the temperatures of —lo° and -|-70o C. (14° and 158° F.) argentic iodide contracts on heating and expands on cooling. Pure argentic iodide is not acted upon by light, but in presence of sub- stances which are capable of combining with the liberated iodine it is slowly blackened. A slight admixture of argentic nitrate produces this effect. By exposure to light, however, even for a very short time, argentic iodide passes into a peculiar active condition, in which it pos- sesses the property of immediately precipitating upon its surface black, finely-divided metallic silver from solutions of silver salts in presence of some reducing agent, such as pyrogallic acid. Upon this property the application of argentic iodide in photography depends, and the process of thus blackening the iodide is that of “ developing” the pho- tographic image. Dry argentic iodide absorbs gaseous ammonia, form- ing a white compound, 2AgI,NH3, which, when exposed to the air, parts with ammonia, and is reconverted into yellow argentic iodide. Argentic fluoride, AgF, is prepared by dissolving argentic oxide or argentic carbonate in hydrofluoric acid, and evaporating the solution. Argentic fluoride crystallizes either in colorless quadratic pyramids with 1 aq., or in prisms with 2 aq. It is deliquescent, and soluble in half its weight of water. It is not readily obtained in an anhydrous state When the compound AgF,OH2 is dried in vacuo, it undergoes partial decomposition, and a brownish-yellow earthy mass is formed, which, when heated with exclusion of air, may be fused, and on cooling solidifies to a black horny sectile mass. Unlike the chlorides of many of the metals, which in the fused state may be subjected to electrolysis, fused argentic fluoride conducts the electric current without undergoing decomposition. When heated in moist air it is reduced to the metallic state. The dry fluoride absorbs 844 times its volume of gaseous am- monia; at higher temperatures ammonia reduces it to metallic silver. COMPOUNDS OF SILVER WITH OXYGEN. The following three oxides of silver are known : Argentous oxide (argentous quadrantoxide), . OAg4, Argentic oxide, • OAg2. Argentic peroxide, / OAg \ 0Ag‘ Argentous oxide, OAg4, is obtained by heating argentic citrate in a current of hydro- gen to 100° C.; on adding potassic hydrate to the solution of the bronze-colored mass thus obtained, argentous oxide is precipitated. It forms a black powder. Hydrochloric and hydrobromic acid convert it into argentous chloride and bromide. Oxy-acids de- compose it, yielding an argentic salt and metallic silver. On heating, it breaks up into metallic silver and oxygen. Argentic oxide, °Ag2, is prepared by precipitating nitrate of silver with potassic hydrate or baryta-water, taking care to avoid the forma- tion of carbonate; or by boiling freshly precipitated argentic chloride with a concentrated solution of potassic hydrate. When precipitated in the cold, it forms a dark-brown powder, which becomes black and 455 COMPOUNDS OF SILVER. anhydrous on drying at 60° or 70° C. (140-158° F.). The recently precipitated and still moist brown oxide is in some respects more active in its combining properties than the dried black oxide; thus it absorbs carbonic anhydride from the air and substitutes hydroxyl for chlorine in the chlorides of organic ammonium bases; it therefore probably con- sists of the hydrate AgHo, which has not, however, been prepared in a state of purity. One part of argentic oxide dissolves in about 3000 parts of water, the solution possessing a marked alkaline reaction. The sp. gr. of the dry oxide is 7.25. In the dry state it acts as a powerful oxidizing agent, inflaming various oxidizable substances, such as sulphur, amorphous phosphorus, and the sulphides of arsenic and antimony, when triturated along with them. At a temperature of 250° C. (482° F.) it is decomposed into silver and oxygen, whilst in a current of hydrogen it undergoes reduction to metallic silver at 100° C. Argentic oxide is the salifiable oxide of silver: OAg2 + S02Ho2 = S02A.go2 + OII2. Argentic Sulphuric Argentic Water, oxide. acid. sulphate. Strong ammonia converts argentic oxide into fulminating silver (q.v.). Argentic peroxide, —This compound is formed by the action of ozone on finely divided silver. When a concentrated solution of argentic nitrate is submitted to electrolysis, argentic peroxide is deposited on the positive electrode. In like manner, in the electrolysis of acidu- lated water, if a silver plate be employed as positive electrode, the nascent oxygen combines with the silver, and the plate becomes coated with argentic peroxide. It forms minute black lustrous octahedra, which are frequently attached to each other. It is decomposed a little above 100° C. into oxygen and argentic oxide. Chlorine rapidly con- verts it at ordinary temperatures into argentic chloride and oxygen. Hydroxyl and argentic peroxide mutually reduce each other, oxygen being evolved from both substances : 2 {gif + 2{gg = 2Ag2 + 20HS + 30, Argentic Hydroxyl. Water, peroxide. Argentic peroxide possesses more powerful oxidizing properties than argentic oxide: when triturated with antimonious sulphide, the mixture deflagrates; sulphuretted hydrogen inflames in contact with the peroxide, the latter being converted into argentic sulphide; in aqueous ammonia the peroxide dissolves with evolution of nitrogen ; when warmed in hydrogen it is reduced to metallic silver with a slight explosion. It seems to possess the properties of a weak base, forming salts which are stable only in solution with an excess of acid. Thus concentrated sul- phuric acid dissolves it, forming a green liquid; but, on diluting with water, oxygen is evolved, and the solution contains argentic sulphate. With strong nitric acid it yields a brownish-red solution, which on dilu- 456 INORGANIC CHEMISTRY. tion with water deposits the unchanged peroxide, the latter then redis- solving in the dilute acid with evolution of oxygen and formation of argentic nitrate. OXY-SALTS OF SILVER. Argentic nitrate, N02Ago.—This salt is prepared by dissolving silver in dilute nitric acid, evaporating the solution, and allowing it to crystallize. It is thus obtained in colorless rhombic tabular crystals of sp. gr. 4.3, which fuse at 198° C. (388° F.), and solidify on cooling to a fibrous crystalline mass. Argentic nitrate is soluble in half its weight of water at ordinary temperatures, less soluble in nitric acid ; soluble in four parts of boiling alcohol. The aqueous solution has a neutral reac- tion. Argentic nitrate has a disagreeable metallic taste, and is very poisonous. Applied to the flesh of animals, it acts as a powerful caustic, destroying the vitality of the part; the fused salt, cast into sticks, in which form it is known as lunar caustic, is employed in surgery for this purpose. The pure salt is not altered by exposure to light; but in con- tact with organic substances, light speedily blackens it. The hot, con- centrated solution dissolves argentic chloride slightly, argentic bromide more readily, and still more readily argentic iodide and cyandide. From these solutions the following compounds are deposited in needles on cooling: N02Ago, AgCl; N02Ago,Agßr; 2N02Ago,AgI; N02Ago,2Ag(CN). These compounds are all decomposed by water with precipitation of the chloride, bromide, etc. Solid argentic nitrate absorbs gaseous ammonia, yielding a compound N02Ago,3NH3.* A concentrated solution of argentic nitrate, when saturated with ammonia, deposits rhombic crystals of the formula N02Ago,2NH3,f Argentic nitrate is extensively em- ployed in photography. It also forms the basis of most of the indelible inks used for marking linen. Argentic nitrite, NOAgo, is precipitated when concentrated solutions of potassic nitrite and argentic nitrate are mixed. It crystallizes in colorless or yellow prisms, which are sparingly soluble in cold, more readily soluble in warm water. At a temperature be- tween 90° and 140° C. (162-284° F.) it is decomposed into metallic silver, nitric oxide, and argentic nitrate : 4NOAgo 2N02Ago + 2/N//0 + Ag2. Argentic Argentic Nitric nitrite. nitrate. oxide. Argentic chlorate, | q, is obtained by dissolving argentic oxide in chloric acid. It is more readily prepared by passing chlorine into water in which argentic oxide is suspended; a mixture of chloride and hypochlorite (cf. p. 181) is thus formed, the latter decomposing in the dark at 60° C. (140° F.) into chloride and chlorate; 30IAgo = 2AgCI + {g“0. Argentic Argentic Argentic hypochlorite. chloride. chlorate. * N(NH2)2Ho2fNvAgH3O). f Nv]SrF2)2Ho2Ago or NO(NH2)Ho(NvAgH3O). COMPOUNDS OF SILVER. 457 The liquid is filtered from the chloride and evaporated. Argentic chlorate crystallizes in white opaque quadratic prisms, soluble in 10 parts of cold water. It fuses at 230° C. (446° F.), and decomposes at 270° C. (518° F.) into oxygen and argentic chloride, a trace of chlorine being evolved at the same time. When rapidly heated it defla- grates. A mixture of argentic chlorate with sulphur detonates with great violence on friction. Argentic bromate, | q^0, and argentic iodate, / oAgo’ are Stained as sparingly soluble precipitates by the addition of solutions of the corresponding potash salts to a solution of argentic nitrate. Argentic periodate.—When argentic nitrate is added to a neutral or slightly acid solution of an alkaline periodate, a dark-brown precipitate of the formula 031Ag0,- ?OAg2 is obtained, which when heated to 200° C. (392° F.) is decomposed into argentic iodide, metallic silver, and oxygen. This salt dissolves in nitric acid, and deposits on evaporation orange-colored octahedra of argentic metaperiodate, O3IAgO, which is de- composed by water into free periodic acid and an insoluble yellow salt of the formula 20sIAgo,OAg2,30H2* Argentic carbonate, COAgo2, is precipitated when potassic or sodic carbonate is added to a solution of argentic nitrate. It forms a pale-yellow amorphous powder, insoluble in water. When exposed to light, or when warmed, it blackens. At a temperature of 100° C. it evolves carbonic anhydride, and is converted into argentic oxide. Argentic sulphate, S02Ago2, is prepared by dissolving silver in hot concentrated sulphuric acid, or by precipitating a concentrated solu- tion of argentic nitrate with sulphuric acid. It forms small lustrous crystals belonging to the rhombic system, of sp. gr. 5.4. It is soluble in about 200 parts of cold and in 68.35 parts of boiling water; more readily soluble in dilute sulphuric or nitric acid. At a dark red heat it fuses without decomposition ; at a higher temperature it breaks up into metallic silver, oxygen, and sulphurous and sulphuric anhydrides. The solid salt absorbs two molecules of gaseous ammonia, forming the compound S02Ago2,2NH3 = S02(NyAgH30)2. A solution of the salt in warm aqueous ammonia deposits on cooling quadratic crystals of the compound S02Ago2,4NH3 = S(NH2)2Ho 2(NTAgH3O2).—Hydric argentic sulphate, S02HoAgo, crystallizes in pale yellow prisms from a solution of the normal salt in less than three parts of sulphuric acid. If more sulphuric acid be employed, double compounds of the acid salt with sulphuric acid are obtained. Argentic sulphite, SOAgo2, is prepared by dissolving argentic oxide in sulphurous acid, or by precipitating argentic nitrate with an alkaline sulphite or with sulphurous acid, avoiding an excess of the precipitant. It crystallizes in white shining needles, or forms a curdy precipitate, only slightly soluble in water. When exposed to light, it blackens. At a temperature of 100° C. it is decomposed into sulphurous anhydride, argentic sulphide, and metallic silver: 2SOAgo2 = S02Ago2 + S02 + Ag2. Argentic Argentic Sulphurous sulphite. sulphate. anhydride. Argentic dithionate, / OH2, is prepared by dissolving argentic carbonate in the aqueous acid. It crystallizes in rhombic prisms. Argentic thiosulphate (Argentic hyposidphite), S02AgoAgs.—When a dilute solution * On the formulation of these compounds on the basis of heptaic iodine, e.g., lOAgo5, and lOHo3Ago2, see p. 305. 458 INORGANIC CHEMISTRY. of argentic nitrate is added to an excess of a solution of sodic thiosulphate, a gray precipitate is formed, consisting of a mixture of argentic sulphide with argentic thio- sulphate. The thoroughly washed precipitate is treated with ammonia which extracts the thiosulphate. On carefully neutralizing the ammoniacal solution with nitric acid the argentic thiosulphate is reprecipitated as a white powder, sparingly soluble in water. It must be quickly dried by pressure, as in the moist state it readily decom- poses into argentic sulphide and sulphuric acid : S02AgoAgs + OH2 = SAg2 + S02Hq2. Argentic Water. Argentic Sulphuric thiosulphate. sulphide. acid. Sodic argentic thiosulphate, S0.2NaoAgs,2OTT,, is obtained by gradually adding, with constant stirring, a solution of sodic thiosulphate to a solution of argentic nitrate till the precipitate no longer redissolves. On adding alcohol to the filtrate, the double salt separates in lustrous laminae. Argentic Phosphate : a. Argentic orthophosphate, POAgos, is precipitated when argentic nitrate is added to a solution of any normal or monohydric alkaline phosphate, nitric acid being liberated in the latter case. It forms a yellow amorphous precipitate, insoluble in water, readily soluble in nitric acid and in ammonia. It becomes dark-colored when exposed to light. When heated it assumes a deep orange-red color, and fuses at a strong red heat without decomposition.—Hydric diargentic orthophos- phate, POHoAgo2, is deposited as a white crystalline powder when ether is added to a solution of the normal salt in excess of phosphoric acid. b. Argentic pyrophosphate, P203Ago4, is obtained as a white precipi- tate when argentic nitrate is added to solutions of either normal or acid pyrophosphates of the alkali metals. It is insoluble in water, readily soluble in nitric acid or ammonia. It fuses without decompo- sition below redness, yielding a dark brown liquid which solidifies on cooling to a radio-crystalline mass. Under the influence of light it turns red. c. Argentic metaphosphate, P02Ago.—The various modifications of metaphosphoric acid yield corresponding silver salts. Thus, if argentic nitrate be added to a solution of the vitreous sodic metaphosphate, an amorphous white precipitate of the silver salt is obtained ; whereas crystal lizable sodic trimetaphosphate yields, when so treated, well- formed crystals of argentic trimetaphosphate, P3O6Ago3,0H2. Argentic arsenate, AsOAgo3, is obtained as a reddish-brown amorphous precipitate when an alkaline arsenate is added to the solution of a silver salt. The same salt may be obtained as a dark-red crystalline powder by precipitating a boiling solution of argentic nitrate with a concentrated solution of arsenic acid. It is insoluble in water, readily soluble in nitric acid and in ammonia. When heated it fuses, yielding a reddish-brown glass on cooling. Argentic arsenite, AsAgo3, is prepared by cautiously adding ammonia to a mixed solution of argentic nitrate and arsenious acid as long as a precipitate is produced. 11 forms a yellow precipitate, readily soluble in nitric acid and in ammonia. On heating, it decomposes into arsenious anhydride, argentic arsenate, and metallic silver: 5AsAgo3 = 3AsOAgo3 + As20s + 3Ag2. Argentic Argentic Arsenious arsenite. arsenate. anhydride. COMPOUNDS OF SILVER. 459 By boiling with sodic hydrate it is decomposed into arsenic anhydride, which dissolves with formation of sodic arsenate, and metallic silver, the latter being mixed with ar- gentous oxide (OAg4). COMPOUNDS OF SILVER WITH SULPHUR. Argentic sulphide, SAg2.—This compound occurs native as silver glance or argentite in blackish-gray regular crystals with a metallic lustre. It has a sp. gr. of from 7.196 to 7.356. Artificial crystals of argentite are obtained when silver is heated in a current of sulphuretted hydrogen, and the same substance may be prepared as a crystalline mass by fusing together silver and sulphur. A black amorphous pre- cipitate of argentic sulphide is formed when sulphuretted hydrogen is passed into solutions of silver salts. Argentic sulphide is insoluble in water, soluble with decomposition in strong nitric acid, insoluble in am- monia. When heated in air, avoiding too high a temperature, it is oxi- dized to argentic sulphate. Cupric chloride in presence of sodic chloride converts it into argentic chloride (see Mexican Amalgamation Process). SULPHO-SALTS OF SILVER. Argentic sulphar senile, AsAgs3, occurs native as proustite or light red silver ore, in red translucent rhombohedral crystals. It generally contains more or less antimony, which is present in isomorphous replacement of a portion of the arsenic. Argentic sulphantimonite, SbAgsg, occurs as pyrargyrite or dark red silver ore, in rhom- bohedral crystals, isomorphous with the preceding. It varies in color from dark red to grayish-black, is opaque, and possesses metallic lustre. COMPOUNDS OF SILVER WITH NITROGEN AND PHOSPHORUS. Fulminating silver.—This compound is formed when freshly precipitated argentic oxide is dissolved in strong ammonia, and the solution is evaporated with the aid of a gentle warmth. It forms black crystals, which when dry explode violently on the slightest touch, and even when moist may be made to explode by shaking the liquid in which they are immersed. Owing to the dangerous character of this compound its composition has not been ascertained with certainty. It is possibly argentic amide, NAgH2. Argentic phosphide is formed when phosphorus is added to molten silver, or when argentic phosphate is fused with charcoal. It is thus obtained as a dark gray mass, which, when strongly heated, parts with a portion of its phosphorus. This compound has not been obtained of constant composition. General Properties and Reactions of the Compounds of Silver.—The salts of silver with colorless acids are colorless. The soluble salts are neutral to test-paper, have an acrid metallic taste, and act as violent irritant poisons. From solutions of silver salts caustic alkalies precipitate brown argentic oxide. Ammonia also precipitates the oxide, which is soluble however in an excess of the precipitant. Sulphuretted hydrogen gives a black precipitate of argentic sulphide, in- soluble in aramonic sulphide, soluble in hot nitric acid. The hydracids precipitate the corresponding haloid compounds of silver (p. 452). Hydrocyanic acid and potassic cyanide give a curdy precipitate of ar- gentic cyanide (AgCy) soluble in excess of potassic cyanide. Argentic cyanide is decomposed on ignition, leaving a residue of metallic silver. Copper, zinc, iron, and other oxidizable metals, further, sulphurous acid 460 INORGANIC CHEMISTRY. and ferrous sulphate, precipitate metallic silver from the solutions of its salts.. Insoluble compounds of silver, when heated with sodic carbonate on charcoal before the blowpipe, are reduced to metallic silver. The silver compounds give no flame spectrum; but the spark spectrum ex- hibits two characteristic bright lines in the green. CHAPTER XXXIII. DYAD ELEMENTS. Section 11. BARIUM, Ba. Atomic weight 137. Probable molecular weight = 137. Sp. gr. be- tween 4.0 and 5.0. Atomicity ", Evidence of atomicity: Baric chloride, Ba"Clo. Baric hydrate, Ba"Ho2. Baric oxide, Ba;/0. History.—Metallic barium was first prepared by Davy in 1808. Occurrence.—Barium is never found native. It occurs abundantly as sulphate in the mineral heavy-spar and as carbonate in witherite. In many calcium minerals it sometimes replaces a portion of the calcium, with which it is isoraorphous. Traces of it are found in various mineral waters and in sea-water. Preparation.—Barium is not reduced from its oxide, hydrate, or carbonate, by heating with charcoal. It may be obtained by the fol- lowing methods : 1. By the electrolysis of the fused chloride (see Preparation of Lith- ium, p. 436). The barium is thus obtained in the form of a metallic powder. 2. By electrolyzing moistened baric hydrate, carbonate, nitrate, or chloride* the negative electrode being formed of mercury. A liquid amalgam of barium is thus obtained, which may be freed from the ex- cess of mercury by pressing through a cloth. The solid amalgam which remains is only slowly oxidized by exposure to the air. On subjecting it to distillation mercury passes over and metallic barium remains in the retort as a porous mass. 3. By acting with sodium amalgam upon a hot concentrated solu- tion of baric chloride, barium amalgam is obtained, which is further treated as above. 4. Barium amalgam is also obtained by passing the vapor of potas- sium or sodium over baric oxide or chloride strongly heated in an iron tube, and extracting the mass with mercury. Properties.—Barium is a pale yellow metal. Its fusing-point appears to be higher than that of cast-iron. It is rapidly oxidized by expo- COMPOUNDS OF BABIUM. 461 sure to the air, and decomposes water at ordinary temperatures like sodium : Ba + 20H2 == H2 + BaHo2. Water. Baric hydrate. COMPOUNDS OF BARIUM WITH THE HALOGENS. Baric chloride, BaCl2,01T2, may be prepared either from the native carbonate or from the native sulphate. The carbonate is dissolved in hydrochloric acid, and the liquid is digested with an excess of the car- bonate in order to precipitate iron and other foreign metals that are present. The addition of a small quantity of baric hydrate facilitates this precipitation. The filtered liquid is acidified with hydrochloric acid and evaporated. In order to prepare baric chloride from the native sulphate, this mineral is ground to a fine powder and then strongly heated with calcic chloride, limestone, and coal. The follow- ing reactions occur: S02Bao" + 4C = BaS" + 4CO. Baric Baric Carbonic sulphate. * sulphide. oxide. BaS" + CaCl2 = BaCl2 + CaS". Baric Calcic Baric Calcic sulphide. chloride. chloride. sulphide. The calcic sulphide unites with the calcic oxide present to form an insoluble calcic oxysulphide, which remains behind when the baric chloride is extracted with water.—Baric chloride crystallizes in colorless lustrous rhombic tables, with 2 aq., permanent in air. The sp. gr. of the crystallized salt is 3.05. It has an unpleasant bitter taste, and, like all the soluble salts of barium, is very poisonous. The anhydrous salt is soluble in 3 times its weight of water at 10° C. (50° F.), and in about 1-| times its weight of water at 100° C. It is almost insoluble in concentrated hydrochloric and nitric acids; in the dilute acids it is soluble, but less freely than in water. Absolute alcohol does not dis- solve it. When heated above 100° C., the crystallized salt parts with its water of crystallization, and is converted into a white powder fusible at a red heat. When fused in air a small quantity of the salt is con- verted into baric oxide with evolution of chlorine; when heated in a current of steam hydrochloric acid is given off below the fusing-point of the salt, and baric hydrate is formed.—Baric chloride is chiefly used in the preparation of the pigment permanent tvhite, which consists of artificial baric sulphate. Baric bromide, Baßr2,20H2, is prepared by dissolving baric carbonate in hydro- bromic acid. The following method is the most convenient; 12.5 parts of bromine and 1 part of amorphous phosphorus are brought together under water. As soon as the color of the bromine has disappeared the liquid, which now contains hydrobromic and phosphoric acid, is neutralized with baric carbonate, rendered alkaline with baryta Water, filtered from the insoluble baric phosphate, and evaporated to the point of crys- 462 INORGANIC CHEMISTRY. tallization. Baric bromide closely resembles the chloride, but is soluble in absolute alcohol. Banc iodide, BaT2,20H2, is prepared like he bromide, substituting iodine for bro- mine. It forms large, colorless, rhombic crystals, which are very deliquescent, and are soluble in alcohol. When exposed to the air it assumes a reddish tint, owing to the liberation of iodine. It may be heated in a close'd vessel without decomposition, but when heated in air the whole of the iodine is expelled, and baric oxide is formed. Baric fluoride, BaF2, is obtained by neutralizing hydrofluoric acid with baric car- bonate or hydrate, or by precipitating a concentrated solution of baric nitrate with potassic or sodic fluoride. It forms a white granular crystalline powder, sparingly soluble in water, readily soluble in nitric, hydrochloric, and hydrofluoric acids. Baric sihcofluoride, SißaF6, is precipitated as a white crystalline powder, when hy- drofluosilicic acid is added to the solution of a barium salt. It is almost insoluble in water, requiring 3500 parts of cold, and 1200 parts of boiling water for its solution; totally insoluble in alcohol. COMPOUNDS OF BARIUM WITH OXYGEN. Baric oxide [baryta), BaO. Ba=o. 01 Baric peroxide, | | . Baric oxide, BaO.—This is the oxide which is formed by the com- bustion of the metal in air. It may be prepared by heating the nitrate, gently at first, in order to avoid frothing, and afterwards to bright redness. The frothing may also be prevented by mixing the nitrate with its own weight of baric sulphate, the presence of the insoluble sulphate-in the product not being objectionable for many purposes to which the baric oxide may be put, for instance in the preparation of baric hydrate. The carbonate may also be converted into baric oxide by heating to a very high temperature, but the whole of the carbonic anhydride can be expelled only with difficulty; hqwever, by mixing the carbonate with carbon, or with some substance which yields carbon when heated, such as tar or resin, the conversion into baric oxide is greatly facilitated, carbonic oxide being evolved, thus : COBao" + C = BaO + 2CO. Baric Baric Carbonic carbonate. oxide. oxide. Much of the baryta employed in sugar refining (p. 464) was prepared by this method. Baric oxide is a grayish-white, porous, friable mass, of sp. gr. 4.73. It is fusible in the flame of the oxyhydrogen blowpipe. It slakes with water, forming baric hydrate, the combination taking place with such energy that, if an excess of water is avoided, the mass becomes incandescent. Baric peroxide, BRq j>, is formed when baric oxide is heated to low redness in oxygen or air. Baric hydrate is also converted into the per- oxide under these circumstances, but less readily, inasmuch as it fuses below the temperature at which the absorption of oxygen occurs. The product obtained by these means is not pure, a portion of the baric COMPOUNDS OF BARIUM. 463 oxide or hydrate escaping conversion. It is also contaminated with iron, silica, and other matters derived from the vessels in which it has been prepared. In order to obtain the substance in a state of purity, the finely-powdered crude product is added in small portions at a time to an excess of dilute hydrochloric acid, any considerable rise of temper- ature being avoided. The crude peroxide dissolves, with formation of baric chloride and hydroxyl (cf. p. 175). To the solution, after filter- ing from insoluble matters, baryta water is carefully added until the silica and ferric oxide, along with a small quantity of hydrated baric peroxide regenerated by the action of the hydroxyl upon the baric hydrate, are precipitated: {oh + BaHo2 = Ba°| + 20H2. Hydroxyl. Baric hydrate. Baric peroxide. Water. This liquid is again filtered, and then supersaturated with baryta. In this way the whole of the remaining hydroxyl regenerates baric perox- ide, which is precipitated in minute prisms or laminae of the formula BRq j* ,80H2. In the moist condition this aquate may be preserved for any length of time in closed vessels, and forms a convenient source of hydroxyl. By drying at 130° C., or at ordinary temperatures in vacuo, it is converted into anhydrous baric peroxide.—Baric peroxide forms a white impalpable powder, insoluble in water, but forming with it the aquate BRq | ,80H2. It fuses at a bright red heat, and is de- composed into oxygen and baric oxide. Heated with steam it evolves oxygen at the same temperature at which the peroxide is formed, and is converted into baric hydrate. Dilute acids dissolve it with formation of a barium salt and hydroxyl; with concentrated sulphuric acid it forms baric sulphate, whilst oxygen mixed with traces of ozone and hydroxyl is evolved. When heated in a current of sulphurous anhy- dride it becomes incandescent, and is converted into baric sulphate: S02 + = So{°Ba)". Sulphurous Baric Baric anhydride. peroxide. sulphate. COMPOUND OF BARIUM WITH HYDROXYL. Baric hydrate, BaHo2.—This compound is formed, with great evolution of heat, by the direct union of baric oxide with water. A hot concentrated solution of equivalent quantities of baric nitrate and sodic hydrate deposits, on cooling, crystals of baric hydrate. Potassic hydrate may be substituted for sodic hydrate in this reaction; but ammonia does not precipitate baric hydrate from solutions of barium salts. On a large scale, baric hydrate is prepared as follows : By heating 464 INORGANIC CHEMISTRY. powdered heavy-spar with carbon a crude baric sulphide is obtained. Moist carbonic anhydride is passed over the heated sulphide, convert- ing it into carbonate : BaS" -f 0O2 + OH2 = OOBao" + SH2. Baric Carbonic Water. Baric Sulphuretted sulphide. anhydride. carbonate. hydrogen. Superheated steam is then passed over the heated carbonate, which parts with carbonic anhydride and forms baric hydrate: COBao" + OH2 = BaHo2 + C02. Baric Water. Baric Carbonic carbonate. hydrate. anhydride. —Baric hydrate crystallizes from water in large four-sided prisms or plates, of the formula BaHo2,801T2, which are soluble in 20 parts of water at ordinary temperatures, and in 3 parts at 100° C. The solution of the hydrate, generally known as baryta water, is much used in chemical analysis, particularly in the determination of carbonic anhy- dride, which it rapidly absorbs, with formation of insoluble baric car- bonate. The crystals of the hydrate are efflorescent, and when exposed in vacuo over sulphuric acid, give off the greater part of their water of crystallization, leaving a white powder of the formula BaHo2,0H2. When heated, the whole of the water of crystallization is expelled, and the hydrate fuses at a red heat, solidifying on cooling to a crystalline mass. It cannot be converted into baric oxide by the action of heat alone. Heated in a current of air, it is converted into baric peroxide with elimination of water : BaHo2 O = j- -f- 0H2. Baric hydrate. Baric peroxide. Water. Baric hydrate was extensively employed in sugar-refining for sepa- rating crystallizable sugar from molasses. It forms with cane sugar an insoluble compound of the formula Cl 2H22OüBaO, which when sus- pended in water and treated with carbonic anhydride is decomposed, yielding insoluble baric carbonate and sugar, the latter dissolving. Strontic hydrate, which, unlike the barium compound, is not poisonous, has of late been substituted for baric hydrate in the sugar industry. OXY-SALTS OF BARIUM. NO Baric nitrate, No!Bao" —This salt is prepared by dissolving the carbonate or the crude sulphide (p. 461) in dilute nitric acid. It crystallizes in colorless, lustrous, regular octahedra, of sp. gr. 3.2. It is soluble in 12 parts of cold, in 3 parts of boiling water ; almost insoluble in concentrated nitric acid; insoluble in absolute alcohol. It fuses at COMPOUNDS OF BARIUM. 465 597° C. (1107° F.). Heated to redness it decomposes, giving off oxygen, nitrogen, and nitric peroxide, whilst a residue of pure baric oxide remains. It is largely employed in pyrotechny for the preparation of green fire. NO Baric nitrite, NOBao"OH2 .—Baric nitrate is moderately heated so as to convert it into nitrite; carbonic anhydride is then passed into the solution of the fused salt to precipitate any baryta that may have been formed ; an excess of alcohol is added to precipitate unaltered nitrate, and the filtered solution is evaporated to the crystallizing point. It is most readily prepared pure by adding baric chloride to a boiling solution of argentic nitrite, and filtering from the argentic chloride. It forms colorless prisms, very soluble in water. | OCI Baric chlorate, Jr q Bao//,OH2, is formed when chlorine is passed into a hot solu- i OCI tion of baric hydrate, but its separation from the chloride which is formed at the same time is a matter of difficulty. It is best prepared by neutralizing a solution of chloric acid with baric carbonate and evaporating to the crystallizing point. It crys- tallizes in colorless monoclinic prisms, with 1 aq., soluble in 4 parts of cold, in less than 1 part of boiling water. f OCI I o Baric perchlorate, ■; q Bao//,40H2, is prepared by neutralizing perchloric acid with 1 0 i °CI baric hydrate or carbonate. It crystallizes in long deliquescent prisms, readily solu- ble in water and in alcohol. Baric carbonate, COBao".—This salt occurs abundantly in nature as witherite. It may be prepared by pouring a solution of baric chlo- ride or nitrate into an excess of a solution of amnionic carbonate, and washing the precipitate with hot water. The native carbonate forms lus- trous crystals belonging to the rhombic system, of sp. gr. 4.29-4.35; that prepared by precipitation is a dense white powder. It is insoluble in pure water; slightly soluble in water containing carbonic anhydride, probably with formation of an unstable acid carbonate. It fuses at a white heat, giving off carbonic anhydride very slowly; but it is more readily decomposed by heat in presence of carbon, or when steam is passed over it (pp. 462 and 463). The artificial carbonate is employed in chemical analysis. Witherite is used in the preparation of the other salts of barium and as a rat poison. Baric sulphate, S02Bao", occurs in large quantities as heavy-spar, sometimes forming distinct veins. It is frequently found in large rhombic crystals. The sp. gr. of the mineral varies between 4.3 and 4.72. By precipitating solutions of barium salts with dilute sulphuric acid, baric sulphate is obtained as a white impalpable powder ofsp.gr. 4.53. It is insoluble in water, slightly soluble in dilute acids. When freshly precipitated it is readily soluble in concentrated sulphuric acid at 100° C., the solution depositing, on cooling, lustrous prisms of dihy- ( S02Ho dric baric sulphate, < Bao" . If the acid solution is exposed to the (SO2Ho air it absorbs moisture, and deposits silky needles of a salt having the 466 INORGANIC CHEMISTRY rso2Ho Bao" ,20H2. Both these salts are decomposed by water, (SO2Ho yielding sulphuric acid and the neutral salt. Artificial baric sulphate is used as a pigment, under the name of permanent white or blanc fixe. The finely ground mineral is also employed for this purpose, but is too crystalline and transparent, and hence lacks “ body.” (80,1 Baric pyrosulphate, -j O Bao//.—Precipitated baric sulphate dissolves in fuming i S02 j sulphuric acid, and the solution, on heating to 150° C. (302° F.), deposits lustrous gran- ular crystals of this salt. It decomposes at a dull read heat, without previously fusing. Baric sulphite, SOBao//, is obtained as a white crystalline precipitate by the addition of an alkaline sulphite to the solution of a barium salt. It crystallizes from its solu- tion in warm aqueous sulphurous acid in six-sided prisms. When heated in air it is converted into sulphate; in closed vessels it yields, when heated, a mixture of sulphate and sulphide. 4SOBao" = SSCtoßao" + BaS. Baric sulphite. Baric sulphate. Baric sulphide. r so Baric dithionate, f gQ2Bao//,20H2.—Preparation, see p. 278. This salt crystallizes in large, lustrous, monoclinic crystals, soluble in 4 parts of water at 18° C. (64° F.), in 1.1 part at 100° C. The solution may be boiled without undergoing decomposition ; but when boiled with hydrochloric acid, it evolves sulphurous anhydride, and baric sulphate is precipitated. In like manner, when the dry salt is ignited it breaks up into sulphurous anhydride and baric sulphate: | gQ2Bao// = SOaBao" + SO2. Baric Baric Sulphurous dithionate. sulphate. anhydride. Baric dithionate is employed in the preparation of the other dithionates. Baric thiosulphate, SO^^//Ba^//,0H2, is obtained as a sparingly soluble crystal- line precipitate when sodic thiosulphate is added to a solution of baric chloride. POBao" Baric orthophosphate, poßao"^lo^^2’ is prepared by add- ing hydric disodic orthophosphate to a solution of baric chloride rendered strongly alkaline with ammonia. It forms a white precipitate, insolu- ble in water, soluble in dilute nitric and hydrochloric acids.—Hydric baric orthophosphate, POHoBao", is precipitated when hydric disodic orthophosphate is added to neutral solutions of barium salts. It is a white crystalline powder, slightly soluble in water, readily soluble in dilute acids.—Tetrahydric baric orthophosphate, is ob- tained by evaporating a solution of the monacid salt in phosphoric acid. It forms colorless crystals, apparently triclinic, with an acid reaction. It is soluble without decomposition in a small quantity of water; excess of water precipitates 'the monacid salt, whilst free phos- phoric acid remains in solution. COMPOUNDS OF BARIUM. 467 COMPOUNDS OF BARIUM WITH SULPHUR. Baric sulphide, BaS", is obtained by passing sulphuretted hy- drogen over the heated oxide. It is prepared on a large scale by heat- ing heavy-spar with carbon. The materials must be thoroughly in- corporated ; otherwise, owing to their infusibility, the action will be only partial. The finely ground heavy-spar is mixed with powdered bituminous coal; the latter fuses, yielding by its decomposition a carbon which permeates the entire mass of the sulphate, and insures its com- plete reduction. The sulphide obtained by this method is always con- taminated with the excess of carbon, and is only used for the prepara- tion of the various salts of barium (see p. 464). Baric sulphide forms a white mass which, when exposed to the air, absorbs water, oxygen, and carbonic anhydride, and is gradually converted into a mixture of sulphate and carbonate. Water dissolves it, but the solution contains a mixture of hydrate and sulphhydrate : 2BaS" -f 20H2 = BaHo2 + BaHs2. Baric Water. Baric Baric sulphide. hydrate. sulphhydrate. The so-called Bolognian phosphorus is a sulphide of barium prepared by heating 5 parts of precipitated baric sulphate with 1 part of carbon. It must be sealed up while still hot in glass tubes. After exposure to sunlight, or to any other light rich in chemically active rays, it displays in the dark a brilliant orange-colored light, and retains this phospho- rescent property, though with gradually diminishing intensity, for some time. The luminosity may be renewed indefinitely often by fresh ex- posure to light. The sulphides of calcium and strontium are also phos- phorescent, and emit a green, blue, violet, or red light, according to the mode of preparation. These various sulphides are at present manufac- tured under the name of luminous 'paints, and are employed for coating clock-faces, match-boxes, and other objects which it is desired to render luminous in the dark. It is necessary, in order that these paints may preserve their efficiency unimpaired, that they should be protected from the moisture of the air. This is effected by a transparent coating of glass or varnish. Baric tetrasulphide, is obtained in pale-red rhombic prisms by boiling a solution of baric sulphhydrate with sulphur and allowing the solution to cool. It is readily soluble in water. Various other polysulphides of barium have been prepared. They are unstable compounds, which in contact, with water are decomposed with formation of the tetra- sulphide. Baric sulphhydrate, BaHsa, is formed along with baric hydrate by the action of water on baric sulphide (supra). It may be prepared pure by saturating a solution of baric hydrate with sulphuretted hydrogen. It forms colorless very soluble crystals, contain- lng water of crystallization. When heated with exclusion of air, it parts with the water of crystallization, and at a higher temperature evolves sulphuretted hydrogen, whilst baric sulphide remains. COMPOUND OF BARIUM WITH HYDEOSULPHYL. 468 INORGANIC CHEMISTRY. General Properties and Reactions of the Compounds of Barium.—The salts of barium with colorless acids are colorless. The soluble salts have a bitter taste and are poisonous. Baric chloride and baric nitrate are both insoluble in absolute alcohol. Sulphuric acid and soluble sulphates produce in solutions of barium salts a white precipi- tate of baric sulphate insoluble in dilute acids. Alkaline carbonates precipitate baric carbonate. Hydrofluosilicic acid gives a colorless crys- talline precipitate of baric silicofluoride. Potassic chromate and potassic dichromate precipitate yellow baric chromate, insoluble in acetic acid. Barium salts color the non-luminous flame yellowish-green. Of the numerous lines in the complex spectrum, the two green lines, Ba« and Ba/5, are the brighest. STRONTIUM, Sr. Atomic weight 87.5. Probable molecular ioeight= 87.5. Sp.gr. 2.5. Fuses at a red heat. Atomicity ". Evidence of atomicity: Strontic chloride, . . . Strontic hydrate, . . . . . . Sr"Cl2. . . . Sr"Ho2. Strontic oxide, .... . . . Sr"0. History.—Hope showed in 1792 that the mineral strontianite con- tained a new earth. The metal was first isolated by Davy in 1808. Occurrence.—Strontium occurs as carbonate in strontianite, and as sulphate in celestlne. Traces of it are present as carbonate in many kinds of limestone, marble, and chalk. It also occurs in minute quan- tities as chloride and sulphate in brine-springs, mineral waters, and sea-water. Preparation.—Strontium is most readily prepared by the electrolysis of the fused chloride. By this means it is obtained in coherent pieces sometimes weighing half a gram.—By heating a saturated solution of strontic chloride with sodium-amalgam, an amalgam of strontium is formed, from which the mercury may be expelled by heating. Properties.—Strontium is a yellow malleable metal. It undergoes rapid oxidation on exposure to the air, burns with a brilliant light when heated, and decomposes water at ordinary temperatures. COMPOUNDS OF STRONTIUM WITH THE HALOGENS. Strontic chloride, SrCl.2,60H2, is prepared like the barium salt (p. .461). It crystallizes in deliquescent hexagonal needles or prisms of sp. gr. 1.603, readily soluble in water, soluble also in alcohol. When heated, it parts with its water of crystallization, leaving the anhydrous salt in the form of a white powder, which fuses at a higher tempera- ture. The anhydrous chloride absorbs dry ammonia, forming the com- pound SrCl2,8NH3. Strontic bromide, SrBro.GOEL, is prepared like the barium salt (p. 461). It resem- bles strontic chloride in its properties. COMPOUNDS OF STRONTIUM. 469 Strontic iodide, SrI2,70ET2, prepared like the barium salt (p. 462). It crystallizes in six-sided plates, and is very soluble. When heated in air it parts with iodine, and is converted into strontic oxide. Strontic fluoride, SrF2, is prepared like the barium salt (p. 462), which it also resem- bles in its properties. Strontic silicofluoride, SiSrF6,20H2, is obtained by neutralizing hydrofluosilicic acid with strontic carbonate, and evaporating to the crystallizing point. It forms rnonoclinic crystals, readily soluble in water. COMPOUNDS OF STRONTIUM WITH OXYGEN AND HYDROXYL. Strontic oxide, Strontia, . . SrO. Sr=0. Strontic peroxide, . • Sro}‘ Sr ,SOH2, is deposited in yellow crystals from the solution obtained by boiling milk of lime with sulphur and filtering hot. COMPOUND OF CALCIUM WITH PHOSPHORUS. Calcic phosphide, /P//2Ca2(?).—This compound has not been prepared pure. It is formed by the direct combination of metallic calcium and phosphorus, when the two substances are heated together under petroleum. It may be obtained mixed with calcic pyrophosphate by passing the vapor of phosphorus over lime heated to redness : 14 CaO + 14P = 5/P//2Ca2 + 2P203Cao//2. Calcic Calcic Calcic oxide. phosphide. pyrophosphate. 484 INORGANIC CHEMISTRY. The mixture thus 'obtained, which forms a reddish-brown mass, is employed in the preparation of liquid phosphoretted hydrogen (p. 343). It also contains tricalcic di- phosphide, P2Ca3. General Properties and Reactions of the Compounds of Calcium.—The calcium salts, as a rule, closely resemble in their properties those of barium and strontium. Those formed with color- less acids are colorless. Calcic nitrate and calcic chloride are both soluble in absolute alcohol. From solutions of calcium salts alkaline carbonates precipitate calcic carbonate. The sulphate of calcium is more soluble than that of strontium ; in dilute solutions of calcium salts sulphuric acid and soluble sulphates produce a precipitate only on addition of alcohol. Amnionic oxalate precipitates white calcic oxalate, soluble in hydrochloric and in nitric acid, insoluble in acetic acid. Calcium compounds color the non-luminous flame yellowish-red. The flame spectrum is complex; the two most characteristic lines are Caa in the orange, and Ca/5 in the green. On Potable Water and on the Impurities occurring in Natural Waters. In describing the properties of water (p. 173), it was mentioned that natural waters always contain impurities; and as some of the most important of these are compounds of two of the metals belonging to the section under consideration, it will be convenient to return here to the subject in order to complete the chemical history of water. Pure water never occurs in nature; as soon as it quits the vaporous condition, and assumes the form of clouds and rain, it becomes more or less contaminated by atmospheric impurities. When it reaches the earth, it flows over surfaces, or percolates through strata, more or less soluble, and thus acquires further impurities in addition to, or some- times in the place of, those which it had previously contracted from the atmosphere. It thus becomes, in some cases more, in others less, suit- able for domestic use. The nature of the changes which water suffers from such influences must obviously depend, to a great extent, upon the character of the geological formations over or through which it passes. If the formation be hard and insoluble, then little saline or other matter is taken up. Thus the River Loka, in Sweden, contains only 0.07 part of solid matter in 100,000 parts of water. Loch Kat- rine contains 3.2 parts per 100,000, Ullswater Lake 3.9 parts, and the Dee at Aberdeen 5.7 parts per 100,000 parts of water. As a rule, however, water meets with more soluble matter than this, and the pro- portion generally varies from 7 to 50 parts in 100,000 parts of water. Thus the Thames and Lea contain about 30 parts, and the water of deep wells sunk into the chalk about 40 parts, per 100,000. An excessive amount of these foreign matters renders the water unpalatable, and constitutes it a mineral or abnormal water. Such ac- cumulations of soluble saline matter take place in the ocean, which contains from 3140 to 4000 parts per 100,000, and in lakes without outlet. Thus the Dead Sea, which is 1312 feet below the level of the Mediterranean, and is fed by the Jordan and six other streams (con- taining on the average 104 parts of soluble solid matter per 100,000) POTABLE WATERS. 485 contains 22,857 parts of solid matter per 100,000. And the Elton lake in Russia contains 27,143 parts per 100,000, although upwards of 200,000 tons of salt are annually extracted from it. We propose here, however, to confine attention chiefly to drinking or potable water—a subject which is, year by year, acquiring an in- creased sanitary importance. Numerous researches, made by both physiologists and chemists, have led investigators to the conclusion that several, at least, of those dis- eases, which are propagated in the manner of epidemics, diffuse them- selves by living germs or spores, which, finding a suitable nidus in the bodies of animals, there multiply and produce that specific disturbance of the normal vital functions which characterizes a disease of the zymotic class. It is indeed in consequence of the extensive prevalence of this view respecting the mode of propagation of such diseases that the term zymotic (from Cu/Aw, I ferment) has come to be almost uni- versally employed to designate them. Long continued observations and carefully compiled statistical records have conclusively demonstrated that drinking-water is the chief medium through which zymotic diseases, especially cholera and typhoid fever, are propagated. In these latter diseases the infectious or zymotic matter is contained in the discharges from the intestinal canal of the patient. Many of our arrangements for disposing of these secretions have the effect of diffusing them through water, and the drinking of such polluted water not unfrequently conveys the infection to whole communities. Shortly stated, it is absolutely necessary for the propa- gation of cholera and typhoid fever, that the excrements of persons suffering from these diseases should be swallowed by other persons. That such an unspeakably disgusting mode of infection is not only possible, but imminent over a very large proportion of the inhabitants of Great Britain, is conclusively proved by the numerous analyses of the water used by them for drinking. So far from the horrible prac- tice just indicated being exceptional, it is the rule. It is a widely spread custom, both in towns and villages, to drink either the water of rivers into which the excrements of man are discharged, or the water from shallow wells which are largely fed by soakage from middens, sewers, or cesspools. Thus many millions of the population are daily exposed to the risk of infection from typhoidal discharges, and periodi- cally to that from cholera dejections. It would obviously be of the very highest importance to mankind, if the presence of cholera or typhoid poison in water could be demon- strated by chemical or microscopical analysis. This is, however, at present impossible. It is only by their action on human beings that their presence can be proved. But, chemical analysis can show us the presence, in water, of excremental matter, or of the characteristic products of its decomposition, although it cannot distinguish between normal and infected excrement. From this point of view, therefore, the analyical examination of water assumes an importance second to no other application of chem- istry. It would be out of place, however, in this work to describe the mode of performing these analyical operations, and we shall therefore 486 INORGANIC CHEMISTRY. confine ourselves to an enumeration of the data obtained in water analysis and to the interpretation of these data. Water Analysis.—The exhaustive chemical examination of a sample of water is one of the most tedious and troublesome operations known to chemists. It requires weeks, sometimes even months, for its com- pletion. This arises partly from the great multiplicity of separate sub- stances which may be present in the water, both in solution and in suspension, partly from the very minute proportion in which these substances sometimes exist, and partly on account of the difficulties attending their exact determination, when they are diffused through vast volumes of water. Such an exhaustive examination includes: 1. The extraction and separate volumetric measurement of the dis- solved gases. 2. The separate determination of the weight of each constituent of the saline matters in solution. 3. The determination of the two chief elements of the organic matters in solution. 4. The separation of the suspended matters, if any, and the determi- nation of their total weight when dry. 5. The separation and determination of each mineral constituent of the suspended matters. 6. The separation and determination, as far as possible, of each or- ganic constituent of the suspended matters. Fortunately, many of the more tedious and laborious of these opera- tions may be omitted, if the object of the analysis be only to ascertain the suitability or otherwise of the water for domestic or manufacturing purposes. Thus, the extraction and volumetric measurement of the gases may be safely dispensed with; since, in the present state of our knowledge, the gaseous constituents of water throw but little light upon its character. The existence of dissolved atmospheric gases in water doubtless adds to its platability ; recently boiled water, for instance, has a notoriously flat and vapid taste, but the solution of these gases by water is so rapid as almost to preclude the possibility of lack of aeration in natural waters. This is seen from the following comparison of the proportional volumes of atmospheric gases expelled on boiling 100 cubic centimetres of rain-water, Welsh and Cumberland upland surface water, Loch Katrine water as delivered in Thames water as de- livered in London, and water drawn from deep wells in the chalk, re- spectively : Volume and Composition of the Gases dissolved in 100 Cubic Centi- metres of Various Waters. Rain water. Cumber- land mountain water. LOch Katrine water. Thames water. Deep chalk well water. Nitrogen, Oxygen, Carbonic anhydride, . 1.308 c.c. 0.637 “ 0.128 “ 1.424 c.c. 0.726 “ 0.281 “ 1.731 c.c. 0.704 “ 0.113 “ 1.325 c.c. 0.588 “ 4,021 “ 1.944 c.c. 0.028 “ 5.520 “ 2.073 “ 2.431 “ 2.548 “ 5.934 “ 7.492 “ POTABLE WATERS. 487 A comparison of the numbers in the foregoing table shows that the total volume of dissolved atmospheric gases differs but little, even in waters from the most widely different sources. It was at one time sup- posed that the proportion of oxygen in these gases was an important item in the history of the water, and a deficiency of this gas was be- lieved to indicate the presence of putrescent organic matters; but the subsequent discovery that deep well waters (in which putrescent or- ganic matter is certainly not present) contained little or no dissolved oxygen, deprived this analytical fact of much of its importance. The large proportion of carbonic anhydride which is present in Thames wrater and in deep chalk well water scarcely adds to the effective aeration of these waters, because nearly the whole of this car- bonic anhydride is in chemical combination with lime, and not in the condition of dissolved gas. The separate determination of the weight of each constituent of the saline matters in solution is also rarely required. These constituents have, with very few exceptions, no appreciable influence upon the whole- someness of the water; hence, in the great majority of cases, it is not necessary to determine the weight of each. Certain of them, however— ammonia, nitrates, nitrites, and chlorides—are very useful in tracing the previous history of the water, and the separate determination of these must, therefore, on no account be omitted. Moreover, if the presence of lead, arsenic, or barium be suspected, these poisonous metals must be carefully sought for, and, if found, their respective quantities determined. The degree of hardness ought also to be ascertained in all cases. The separation and determination of each mineral constituent of the suspended matters may be dispensed with, unless poisonous substances occur amongst them. The separate determination of each organic constituent of the sus- pended matter is of comparatively little use in the present state of our knowledge, because it is impossible to distinguish, amongst the sus- pended matters in water, those which are injurious from those which are harmless. The really injurious organic suspended matters are probably not merely organic but organized matters, entozoic ova, or zymotic germs, capable of reproduction in the human body with the simultaneous development of disease. Investigations of this class belong rather to microscopical than to chemical analysis, but even mi- croscopic research is not yet competent to reveal any facts of direct im- portance in connection with such organized suspended matters. The microscope has rarely if ever discovered, even in the most pol- luted drinking water, any germ or organism which is known to be deleterious to human health; but by showing the presence of living organisms in water, it proves, either that the water has not been so effi- ciently filtered as to remove these organisms, or that it has subsequently become polluted by them; and thus it is indirectly demonstrated that the water has not been treated, preserved, or stored under such condi- tions as would preclude the access of deleterious germs or organisms. A microscopic examination of the suspended matters in potable waters thus becomes indirectly of considerable importance. 488 INORGANIC CHEMISTRY. The analytical determinations, deemed sufficiently important to warrant the expenditure upon them of the necessary time and labor, are the following; those which are of primary importance being printed in bold type: In Solution. 1. Total solid matters. 2. Organic carbon, or carbon contained in the organic matter actually present. 3. Organic nitrogen, or nitrogen contained in the organic mat- ter actually present. 4. Ammonia. 5. Nitrogen as nitrates and nitrites. 6. Total combined nitrogen. 7. Estimation of the previous sewage or animal contamination. 8. Chlorine. 9. Temporary, permanent, and total hardness. In Suspension. 10. Mineral matters in suspension. 11. Organic matters in suspension. We have now to explain the object and significance of each of these determinations. 1. Total Solid Matters in Solution, or Total Solid Impurities.—When water is evaporated to dryness, there is left behind a solid residue con- taining the mineral and organic matters with which the water had be- come contaminated since its condensation from the atmosphere. Leav- ing out of consideration the quality of the ingredients contained in potable waters, the proportion of solid residue left on evaporation affords an approximate, though rough, indication of the comparative purity of such waters. On the one hand it may be safely concluded, that waters leaving very large residues on evaporation are unfit for domestic use, whilst on the other, those containing very small residues are, on this account alone, well adapted for such purposes, and but very rarely con- tain amongst their constituents any which are seriously objectionable. Not only do waters leaving small residues on evaporation generally possess a superiority for domestic purposes, but they are also much more valuable than less pure waters for a large number of manufacturing purposes. Thus, for the feeding of steam boilers, their use precludes the formation of incrustations, which not only seriously interfere with the transmission of heat from the fuel to the water, but are probably a frequent cause of disastrous explosions. 2. Organic Carbon.—From a sanitary point of view, the most im- portant constituent of the -total solids is organic matter, and various processes have from time to time been devised for the quantitative de- termination of this matter or of some of its constituents. The problem is surrounded with unusual difficulties, and hitherto no method, worthy of any degree of confidence, has been discovered by which the weight of organic matter dissolved in water can be even approximately deter- POTABLE WATERS. 489 mined. Even of several analytical processes which do not pretend to the estimation of the total weight, and aim at the quantitative deter- mination of only some of the elements of the organic matter, there is only one which yields trustworthy results. This process is both trouble- some and tedious, and requires considerable manipulative skill; but it is the only method which throws any light whatever upon the actual pollution of water by organic matter. It consists in transforming by combustion in close vessels the carbon and nitrogen of the organic mat- ter into carbonic anhydride and free nitrogen, and then measuring the respective volumes of these gases. By a simple calculation, the weights of carbon and nitrogen contained in the original organic matter present in the water can be arrived at, from these volumetric determinations, with great precision. The weight of organic carbon, or carbon contained in the organic matter found in different samples of water, indicates the amount of organic matter with which the water is contaminated, but it does not indicate the source, animal or vegetable, whence that organic matter was derived. Cceteris paribus, the smaller the proportion of organic carbon, the better the quality of the water. Even if the source of the organic matter be altogether vegetal, experience has shown that a proportion of organic carbon larger than 0.2 part in 100,000 parts of water is undesirable, because it renders the water slightly bitter and unpalatable. A larger proportion of organic carbon, if it be contained in animal matter, does not interfere with the palatability of the water, but it exposes the consumer to the risk of infection, and potable water which contains organic matter, even only partially derived from animal sources, should not yield much more than 0.1 part of organic carbon from 100,000 parts of water. 3. Organic Nitrogen.—The character of the organic matter con- tained in potable water, that is to say, its animal or vegetable origin, may in most cases be judged of by the relative proportions in which the two elements, carbon and nitrogen, occur in the organic matters. Hence the necessity for determining the amount of organic nitrogen in waters used for domestic purposes. This determination, taken in connection with that of organic carbon, frequently affords information of great value as to whether the organic matter is of animal or vegetable origin ; and this information acquires additional importance and trustworthi- ness when it is subsequently tested by a chemical investigation of the previous history of the water as revealed by the proportions of the chief products derived from sewage and animal matters, viz., ammonia, nitrates, nitrites, and chlorine. The smaller the absolute quantity of organic nitrogen, and the less the proportionate amount as compared with organic carbon, the better is the quality of the water as regards present or actual pollution, and the less likely is the water to contain any organic matters of animal origin. In connection with this part of the analytical investigation, however, it must be borne in mind that vegetable organic matter is far from being destitute of nitrogen.. Peat, for instance, which is a form of vegetable matter least likely to contain nitrogen, yields to water organic substances in solution containing much nitrogen. Doubtless, different samples of peat vary in the nitrogenous character of the soluble vegetable matter which they contain ; but, on 490 INOEGANIC CHEMISTEY. the average, the proportion of nitrogen to carbon may be taken to be N : C = 1 : 12, and it is found that such peaty matters dissolved in water may, after prolonged exposure to oxidizing influences, lose carbon so much more rapidly than nitrogen, as materially to increase the pro- portion of the latter element to the former. The following table shows the proportion of nitrogen to carbon in waters containing organic matter of peaty origin : Proportion of carbon to 1 part Unoxidized peaty matter contained in upland surface water, 11.92 Peatv matter contained in upland surface water after exposure to atmospheric oxidation in natural lakes or artificial reservoirs, Peaty matter contained in spring water, .... 3.21 'Thus the proportion of carbon to nitrogen in the peaty organic mat- ter of water decreases rapidly as oxidation progresses. After storage for weeks or months in lakes it is reduced to one-half its original amount; but after the water containing the peaty matter has been sub- jected to the powerful oxidizing influences which accompany filtration through porous strata, it reappears as spring water with a greatly aug- mented proportion of organic nitrogen, although the absolute quantity has greatly diminished. In other words, large quantities of both carbon and nitrogen have been oxidized and converted into mineral matter, but the carbon has undergone this transformation more rapidly than the nitrogen. This concentration of nitrogen during oxidation assimilates oxidized vegetable to unoxidized animal organic matter in chemical composition, so far, at least, as the proportion between the chief elements, nitrogen and carbon, is concerned. There is still, however, a considerable differ- ence in this respect between these two kinds of organic matter; but even this disappears when the water containing animal organic matter is sub- jected to oxidizing influences; for whilst vegetable organic matter suffers a concentration of nitrogen during oxidation, animal organic matter exhibits, as a rule, a concentration of carbon, and a diminution in the proportion of nitrogen under the same influence. Thus the proportions of nitrogen to carbon in soluble vegetable and animal organic matters vary in opposite directions during oxidation ; a fact which renders more difficult the decision as to whether the organic matter present in any given sample of water is of animal or vegetable origin. This difficulty can, however, be greatly diminished or entirely overcome by an appeal to the previous history of the water as revealed partly by a knowledge of its source, and of the kind of contamination to which it has been exposed, and partly through the information af- forded by chemical analysis. In the first place, if the water is known by an inspection of its source to have been polluted by animal matters, and to have been subjected, after such pollution, only to the slight oxida- tion effected in rivers or streams, a portion at least of the organic matter which it contains must have been derived from animal matter. For POTABLE WATERS. 491 there is no river in Great Britain long enough to completely oxidize or destroy the soluble animal organic matter present in polluted water. In the second place, if the water is found, on analysis, to contain consid- erable quantities of one or more of the mineral compounds—ammonia, nitrates, and nitrites—into which animal organic matter is resolved during its decomposition or oxidation, the inference may be drawn that the soluble organic matter of such water is derived from animal sources. But this inference must only be provisional; it must stand or fall by an investigation into the source of the water; for although the presence of the products of the decomposition of animal matter indubitably con- victs the water of previous pollution, yet it is obviously possible, from the facts and considerations which have just been adduced, that the whole of the original organic matter may have been oxidized and con- verted into innocuous mineral compounds during the prolonged filtra- tion of the water through a great thickness of porous strata, and that the water so purified may afterwards have become contaminated with vegetable matter only. In other words, water polluted by animal mat- ters may become pure spring water, retaining only the innocuous evi- dence of its former pollution, and may then become polluted by the soluble matter of peat. Such water would be suspicious owing to the evidence of its previous pollution, which it still bears about with it, and it could only be cleared from this suspicion on proof of efficient purification after its pollution by animal matter. To render the water safe for domestic use the animal pollution must have occurred before it became spring water. It is upon this part of the investigation of potable water that the next four determinations have a very important bearing. 4. Ammonia.—This mineral nitrogenous compound is rarely absent from potable waters, which derive it, sometimes from the atmosphere, but more usually from decomposing animal matters. Bain water fall- ing in London sometimes contains as much as 0.21 part of ammonia in 100,000 parts of water, but this is exceptional, and the proportion rarely exceeds one-third of that amount. The average quantity present in 71 samples of rain water collected at Rothamsted, near St. Albans, was 0.049 part in 100,000 parts of water. In river water the proportion rarely exceeds 0.01 part, in unpolluted well water it is usually less, whilst in spring water it is either absent altogether or present in only very minute proportion. On the other hand, it often abounds in the water of much polluted shallow wells. The proportion of ammonia in the London shallow well waters sometimes rises as high as 2.7 5 parts in 100,000 parts of water. In contact with animal matter and under the operation of oxidizing influences, ammonia is very rapidly converted into nitrites and nitrates, and its presence therefore in considerable pro- portion in shallow well waters indicates their very recent contamination with animal matters. Its occurrence in water from deep wells, however, does not permit of the same inference being drawn, because we find that in such water the decomposition of nitrates not unfrequently gives rise to ammonia. This is particularly the case in very deep wells, and in those which are sunk into the Chalk beneath the London Clay. The ammonia which occurs under such circumstances is obviously still more 492 INORGANIC CHEMISTRY. remote from the animal matter whence it originated, than the nitrates from which it was immediately derived, and which were themselves generated by the oxidation of animal matter. The chief significance attaching to the determination of ammonia in potable water lies in the circumstance that this compound is derived almost exclusively from the decomposition of animal matter. It is obvious, however, from the consideration just mentioned, that all infer- ences to be drawn from its presence must be controlled by a study of the physical and chemical history of the water. 5. Nitrogen as Nitrates and Nitrites.—In the presence of oxygen, the nitrogen of animal matters is transformed, in great part, into nitric acid and nitrous acid; and these, by combining with the basic substances always present in polluted water, are in their turn transformed into nitrates and nitrites. This transformation takes place most rapidly and completely when the polluted water soaks through aerated soil. Thus 97 per cent, of the combined nitrogen of London sewage is converted into nitrates during its slow percolation through a stratum of gravelly soil only 5 feet thick. Whilst the oxidation of animal matters in solution in water yields abundance of nitrates and nitrites, vegetable matters furnish under like circumstances none, or mere traces, of these compounds. Upland waters, which have been in contact only with mineral matters or with the vege- table matter of uncultivated soil, contain, if any, mere traces of nitrogen in the form of nitrates and nitrites ; but as soon as the water comes into contact with cultivated land, or is polluted by the drainage from farm- yards or human habitations, nitrates in abundance make their appear- ance. The presence of these salts in sufficient quantity is, therefore, trustworthy evidence of the previous pollution of the water with animal matters. It must be borne in mind, however, that nitric and nitrous acids are present, though in but minute quantity, in the atmosphere, and that rain washes them out of the air through which it falls. In 71 samples of rain water collected at Rothamsted the proportion of nitrogen as nitrates and nitrites varied from 0 to 0.044 part in 100,000 parts of water. Even the highest proportion, which occurred only once, is a very small one, and one that is never met with in unpolluted upland waters. 6. Total combined Nitrogen.—The element nitrogen may exist in water in four forms; viz.; firstly as a constituent of organic matter, secondly as a constituent of ammonia, thirdly as a compound of nitrates and nitrites, and fourthly as a constituent of dissolved atmospheric air. In the last case, the nitrogen is in the free or elementary condition ; and as it neither pollutes the water nor throws any light upon its previous pollution, it may be left out of consideration. In all the other three forms, the nitrogen is combined with other elements, constituting either polluting matter or the resultant of previously existing polluting mat- ter. With a slight deduction for the minute amount of this element which is met with in combination in rain water, the determination of total combined nitrogen sums up, as it were, the evidence of the past and present pollution of each water by nitrogeneous organic matter of either animal or vegetable origin. The evidence is unfortunately de- POTABLE WATERS. 493 fective, especially in spring and summer, because some of the compounds containing nitrogen constitute an important part of the food of both animal and vegetable organisms. Combined nitrogen also suffers dimi- nution whenever the organic matter in the water enters into putrefac- tion or undergoes oxidation in the absence of atmospheric oxygen and in the presence of nitrates and nitrites. The latter salts supply, under these circumstances, the oxygen required to transform the carbon and hydrogen of the organic matter into carbonic anhydride and water, whilst their nitrogen is converted only to a slight extent into ammonia, the rest being set free and consequently ceasing to exist as combined nitrogen. It is thus that the water of very deep wells frequently retains few or no traces of the nitrates and nitrites which it previously held in solution, whilst a comparatively small proportion of ammonia is found in their place. The artesian wells of London afford striking instances of this destruction of nitrates and consequently of combined nitrogen. 7. Previous Sewage or Animal Contamination.—lt has been established by very numerous chemical analyses, that animal matters dissolved in water, such as those contained in sewage, the contents of privies and cesspools, or farmyard manure, undergo oxidation in lakes, rivers, and streams very slowly, but in the pores of an open soil very rapidly. When this oxidation is complete, they are resolved into min- eral compounds;—their carbon is converted into carbonic anhydride, and their hydrogen into water, products which can no longer be iden- tified in the aerated waters of a river or spring; but their nitrogen is transformed partly into ammonia, chiefly however into nitrous and nitric acids, which, combining with the bases present in nearly all water that has been in contact with the earth, form nitrates and nitrites, and frequently remain dissolved in the water for a long time;—there con- stituting a record of the sewage or other analogous contamination, to which it has been subjected since its last descent to the earth as rain. It is convenient to have a concrete expression for the amount of pre- vious animal contamination revealed by this record of the past history of water. Such an expression is obtained by taking as a standard of comparison the amount of total combined nitrogen contained in solution in 100,000 parts of average London sewage. Although a considerable proportion of this nitrogen is found at the sewer outfall in the condition of ammonia, it is well known that in the perfectly fresh sewage the nitrogen of this ammonia was present as a constituent of animal organic matter. The earlier analyses of London sewage made by Hofmann and Witt, give the number 8.363 as the amount of total combined nitrogen contained in 100,000 parts of average London sewage. More recent analyses show that 100,000 parts of average London sewage now con- tain only 7.06 parts of total combined nitrogen. This difference is doubtless owing to the more abundant supply of water to the metropolis at the later period. For simplicity, however, a round number (10) is assumed as the amount of total combined nitrogen in solution in 100,000 parts of average London sewage. In estimating, in terms of this standard, the previous animal contam- ination of water, from the proportion of nitrogen, in the form of ammonia and of nitrates and nitrites, which it holds in solution, it is 494 INORGANIC CHEMISTRY. necessary to bear in mind that rain water itself contains these sub- stances, although in minute quantities. The average composition of samples of rain water collected at Rothamsted gives the amount of nitrogen in these forms as 0.032 in 100,000 parts of water. After this number (0.032) has been substracted from the amount of nitrogen, in the forms of nitrates, nitrites, and ammonia, found in 100.000 parts of a potable water, the remainder, if any, represents the nitrogen derived from oxidized animal matters with which the water has been in contact. Thus a sample of water which contains, in the forms of nitrates, nitrites, and ammonia, 0.326 part of nitrogen in 100.000 parts, has obtained 0.326 nitrogen from animal matters. Now this last amount of combined nitrogen is assumed to be contained in 2940 parts of average London sewage, and hence such a sample is said to exhibit 2940 parts of pre- vious sewage or animal contamination in 100,000 parts; or in other words, 100,000 lbs. of the water contain the mineral residue of an amount of animal organic matter equal to that found in 2940 lbs. of average London sewage. It must not be forgotten, however, that the absence of nitrogen in these forms is not absolutely conclusive evidence of immunity from this pollution. There are several agencies at work by which this testimony, as to the amount of animal matter previously in the water, may be weakened or altogether destroyed. Thus we look in vain for the full evidence of previous animal pollution in the effluent water from fields irrigated with sewage; because the growing plants have removed a considerable proportion of ammonia, nitrates, and nitrites, from the liquid as it flows amongst their rootlets. In like manner the aquatic vegetation of rivers, lakes, and reservoirs, slowly removes these com- pounds from the water, and to that extent destroys the evidence of anterior animal contamination. Nitrates and nitrites are also rapidly destroyed when the organic matter in the water containing them enters into putrefaction, a condition which often occurs in streams or reser- voirs containing much polluting organic matter. The same not unfre- quently takes place in water-bearing strata far removed from the surface, although the water in this case may contain but a comparatively small amount of organic matter ; the latter, however, cut off from a supply of atmospheric oxygen—as in the Chalk beneath the London Clay for instance—accomplishes its oxidation at the expense of the nitrates or nitrites, and thus destroys them. Owing to this cause, the evidence of previous animal contamination is not met with in the water of some deep wells in which it might otherwise be expected to occur. The previous animal contamination of water, as deduced from chem- ical analysis, must therefore always be regarded as a minimum quantity; it does not denote the comparative freedom of different samples of water from anterior pollution; but whenever analysis shows this ex- cess of nitrogen in the shape of nitrates, nitrites, and ammonia, the water stands convicted of previous contamination at least to the extent so indicated. The importance of the history of water as regards its anterior pollu- tion with organic matters of animal origin, does not arise from the POTABLE WATERS. 495 presence of the inorganic residues (nitrates, nitrites, and ammonia) of the original polluting matters, for these are in themselves innocuous, but from the risk lest some portion (not detectable by chemical or microscopical analysis) of the noxious constituents of the original animal matters should have escaped that decomposition, which has resolved the remainder into innocuous mineral compounds. This evidence of previous contamination implies, however, much more risk when it occurs in water from rivers and shallow wells, than when it is met with in the waters of deep wells or of deep-seated springs. In the case of river water, there is great probability that the morbific matter, sometimes present in animal excreta, will be carried rapidly down the stream, escape decomposition, and produce disease in those persons who drink the water; for the organic matter of sewage undergoes decomposition very slowly when it is present in running water. In the case of shal- low well water, the decomposition and oxidation of the organic matter are also very liable to be incomplete during the rapid passage of pol- luted surface water into shallow wells. In the case of deep well and spring water, however, if the proportion of previous contamination do not exceed 10,000 parts in 100,000 parts of water, this risk is very inconsiderable, and may be regarded as nil if the direct access of water from the upper strata be rigidly excluded ; because the prolonged filtra- tion to which such water has been subjected in passing downward through so great a thickness of soil or rock, and the rapid oxidation of the organic matters contained in water, when the latter percolates through a porous and aerated soil, afford a considerable guarantee that all noxious constituents have been removed. It has been already stated that chemical analysis cannot discover the noxious ingredient or ingredients in water polluted by infected sewage or animal excreta; and as it cannot thus distinguish between infected and non-infected sewage, the only perfectly safe course is to avoid altogether the use, for domestic purposes, of water which has been pol- luted with excrementitious matters. Nevertheless, as it is very difficult in some localities to obtain water which has not been more or less polluted by excrementitious matters, it is desirable to classify such previously contaminated drinking waters into Reasonably safe water. Suspicious or doubtful water. Dangerous water. Reasonably Safe Water.—Water, although it exhibits previous sew- age or animal contamination, may be regarded as reasonably safe when it is derived either from deep wells (say 100 feet deep), or from deep- seated springs; provided that all contaminated surface water has been rigidly excluded from the well or spring, and that the proportion of previous contamination does not exceed 10,000 parts in 100,000 parts of water. Suspicious or doubtful water is, first, river or flowing water which exhibits any proportion, however small, of previous sewage or animal contamination; and, secondly, well or spring water containing from 496 INORGANIC CHEMISTRY. 10,000 to 20,000 parts of previous contamination in 100,000 parts of water. Dangerous water is, first, river or flowing water which exhibits more than 20,000 parts of previous animal contamination in 100,000; secondly, river or flowing water containing less than 20,000 parts of previous contamination in 100,000 parts, but which is known, from an actual inspection of the river or stream, to receive sewage, either dis- charged into it directly or mingling with it as surface drainage; thirdly, as the risk attending the use of all previously contaminated water increases in direct proportion to the amount of such contamina- tion, well or deep-seated spring water exhibiting more than 20,000 parts of previous contamination in 100,000 must be regarded as dan- gerous. River or running water, containing less than 10,000 parts of previous animal contamination, should only be provisionally placed in the class of suspicious waters, pending an inspection of the banks of the river and tributaries ; which inspection will obviously transfer it either to the class of reasonably safe waters, if the previous contamination be derived exclusively from spring water, or to the class of dangerous w7aters, if any part of the previous contamination be traced to the di- rect admission of sewage or excrementitious matters. 8. Chlorine.—The chlorine found in potable waters is always com- bined with other elements, and chiefly with sodium in the form of sodic chloride or common salt. A knowledge of the proportion of chlorine in water often throws important light upon the history of the water as regards its previous contamination with the liquid, as distinguished from the solid excrements of animals. Human urine contains about 500 parts of chlorine or 824 parts of common salt in 100,000 parts, whilst upland surface water free from previous or present pollution rarely contains more than 1 part of chlorine or 1.648 parts of common salt in the same weight; and it is present in but comparatively minute proportion in the solid excrements of animals. It is scarcely necessary to state that the determination becomes valueless, for the purpose of indicating previous sewage contamination, in the neighborhood of the sea and of natural deposits of salt. The normal proportion of chlorine, as common salt, existing in British waters which have never been pol- luted by excrementitious matters is, as just stated, about 1 part in 100.000 parts of water; but it varies considerably in different parts of the country. Thus at the Land’s End with a strong wind from the S.W. even rain water contains as much as 21.8 parts of chlorine in 100.000 parts, while the Gelder Burn at Balmoral contained on March 9th, 1872, only 0.35 part in 100,000 parts. Unpolluted rivers and lakes in inland countries contain still less. Thus the Rhine at Schaff- hausen contains only 0.2 part, and the lakes of Zug and Zurich 0.27 and 0.17 part respectively in 100,000 parts of water. The proportion of chlorine in rain water varies in like manner, and the variation is also here doubtless due to the varying distance from the sea at which the rain falls. Thus whilst rain water at the Land’s End was found to contain 21.8 parts, the average proportion of rain falling in the centre of India was only 0.03 part. POTABLE WATERS. 497 9. Hardness.—Some of the mineral substances which occur in solu- tion in potable waters communicate to the latter the quality of hardness. Hard water decomposes soap, and cannot be efficiently used for washing. The chief hardening ingredients met with in potable waters are the salts of lime and magnesia. In the decomposition of soap, these salts form curdy and insoluble compounds containing the fatty acids of the soap, and the lime and magnesia of the salts. So long as this decompo- sition goes on, the soap is useless as a detergent, and it is only after all the lime and magnesia salts have been decomposed at the expense of the soap, that the latter begins to exert a useful effect; as soon as this is the case, however, the slightest further addition of soap produces a lather when the water is agitated, but this lather is again destroyed by the addition of a further quantity of the hard water. Thus the addi- tion of hard water to a solution of soap, or the converse of this opera- tion, causes the production of the insoluble curdy matter above men- tioned. These facts render intelligible the process of washing the skin with soap and hard water: The skin is first wetted with the water and then soap is applied; the latter soon decomposes all the hardening salts contained in the small quantity of water with which the skin is covered, and there is then formed a strong solution of soap which penetrates into the pores. This is the process which goes on whilst a lather is being produced in personal ablution ; and now the lather, and the impurities which it has imbibed, require to be removed from the skin,—an opera- tion which can be performed in one of two ways, viz., either by wiping the lather off with a towel, or by rinsing it away with water. In the former case, the pores of the skin are left filled with soap solution; in the latter they become clogged with the greasy, curdy matter which re- sults from the action of the hard water upon the solution which had previously gained possession of the pores of the cuticle. As the latter process of removing the lather is the one universally adopted, the ope- ration of washing with soap and hard water is analogous to that used by the dyer and calico printer when he fixes his pigments in calico, woollen, or silk tissues. The pores of the skin are filled with insoluble, greasy, and curdy salts of the fatty acids contained in the soap, and it is only because the insoluble pigment produced is white, or nearly so, that such a repulsive operation is tolerated. To those, however, who have been accustomed to wash in soft water, the abnormal condition of the skin thus induced is for a long time extremely unpleasant. Of the hardening salts present in potable water, carbonate of lime is the one most universally met with ; and to obtain a numerical expres- sion for this quality of hardness, a sample containing 1 lb. of carbonate of lime or its equivalent of other hardening salts in 100,000 lbs. is said to have one degree of hardness. Each degree of hardness indicates the destruction and waste of 12 lbs. of the best hard soap by 100,000 lbs. or 10,000 gallons of the water, when used for washing. Hard water frequently becomes softer after it has been boiled for some time. When this is the case, a portion at least of the original hardening effect is due to the acid carbonates of lime and magnesia. These salts are decomposed in boiling water into free carbonic anhy- dride, which escapes, and the carbonates of lime and magnesia. The 498 INORGANIC CHEMISTRY. latter, being nearly insoluble in water, cease to exert more than a very slight hardening effect. As the hardness resulting from the carbonates of lime and magnesia is thus removable by boiling the water, it is designated temporary hardness, whilst the hardening effect which is due chiefly to the sulphates of lime and magnesia, and cannot be got rid of by boiling, is termed permanent hardness. The total hardness of a water is therefore commonly made up partly of temporary and partly of permanent hardness. Hard water not only acts injuriously when it is used for washing; but, when it is employed for the generation of steam, it forms trouble- some and dangerous incrustations in the boiler. A constant supply of hot water has become almost a necessity in every household, but great difficulties are thrown in the way of its attainment by the* supply of hard water to towns, owing to the formation of thick calcareous crusts in the heating apparatus. Waters which have much temporary hard- ness are most objectionable in this respect, and the evil is so great where the heating is effected in a coil of pipe, as practically to prevent the use of this most convenient mode of heating water. The hardness of rain water varies from 0° to 10°, The latter degree of hardness is, however, only attained near the seashore and in rough weather. At Rothamsted, in seventy-one samples, it never exceeded 1.7° and averaged only 0.49°. The hardness of water which has once touched the earth depends obviously upon the character of the gather- ing ground or water-bearing stratum over or through which it passes, and also upon the length of time during which it has been in contact with the earth. Calcareous and magnesian soils or strata cause the water passing over or through them to be hard. If the calcareous or magnesian matter contain carbonate of lime or carbonate of magnesia, a portion at least of the hardness will be temporary. If, on the other hand, gypsum (sulphate of lime) be the calcareous material, the hard- ness will be permanent. Unpolluted water collected from Igneous rocks, either as surface drainage or springs, is the softest. Its hardness varies from 0.4° to 5.9°, and averages 2.4°. Next to this in softness, must be ranged the unpolluted waters from Metamorphic, Cambrian, Silurian, and Devonian rocks, the Millstone Grit, London Clay, and Bagshot Beds, which range from 0.4° to 32.5°, and average 5,6°. The Lower Greensand also yields very soft water (about 4° of hard- ness) when the water does not previously percolate through calcareous strata, but this is so rarely the case as to prevent any reliance from being placed upon the softness of Greensand water. The hardness of unpolluted Greensand water sometimes ranges as high as 44°. Amongst the slightly calcareous strata, the New Red Sandstone generally yields water of medium hardness; a large proportion of the hardness is, however, frequently permanent. In fifty-one samples of unpolluted New Red Sandstone water, the temporary hardness ranged from 0° to 19.8°, and averaged 7.7° ; whilst the total hardness varied from 5.7° to 35.7°, and averaged 17.9°. Of true calcareous strata, the Mountain Limestone yields water of least total hardness, whilst the permanent hardness is in general only a small proportion of the total. The analysis of nineteen samples of un- POTABLE WATERS. polluted limestone water showed a total hardness varying from 9.8° to 27.9°, and averaging 15.7°. The permanent hardness ranged from 3.3° to 12.9°, and averaged 7.1°. The Dolomite or Magnesian Limestone generally imparts to water great hardness, of which a large proportion, and sometimes nearly the whole, is permanent. This stratum occupies, however, a comparatively small area in this country, and the water is consequently but little used for domestic purposes. In five samples the total hardness varied from 14.7° to 67.3°, and averaged 41.2°; whilst the permanent hardness varied from 8.3° to 40.8°, averaging 24.8° ; and the temporary hard- ness from o.B° to 26.5°, averaging 16.4°. The Lias yields water of variable, but nearly always great, hardness. The permanent hardness of water from this geological formation is also almost invariably high. In ten samples, the total hardness ranged from 10.3° to 50°, and averaged 29° ; the permanent hardness varied from 1.7° to 17.1°, averaging 8.2° ; and the temporary hardness from 8.6° to 35.3°, averaging 20.9°. The Oolite and Chalk strata yield water of great, but chiefly tempo- rary, hardness. In forty-two samples of unpolluted Oolitic water, the total hardness ranged from 4.2° to 35.2°, and averaged 22.4°; the permanent hardness varied from 3.5° to 13.5°, averaging 6.1° ; whilst the temporary hardness was from 0° to 25.7°, and on the average 16.3°. In ninety-five samples of unpolluted water from the Chalk, the total hardness ranged from 12.4° to 50°, and averaged 26.1°; the perma- nent hardness ranged from 2.7° to 13.8°, averaging 6.1°; whilst the temporary hardness varied from 6.8° to 38.6°, and averaged 20.2°. The Chalk beneath the London Clay yields water which is usually much softer than that obtained from Chalk which is not covered by an impervious stratum. In fourteen samples of water from this source, the total hardness ranged from 0.9° to 48.5°, the average being 18.9°; the permanent hardness varied from 0.9° to 25.4°, but this extreme number and the extreme of total hardness occurred only in the water from a deep well at Harrow-on-the-Hill. Omitting this well, the ex- treme total hardness was 28.2° and the extreme permanent hardness 9.7°; whilst, omitting the Harrow sample, the temporary hardness varied from 0° to 21.2, and averaged 7.1°. The Coal Measures yield water of very variable hardness, owing to the variety in chemical composition presented by these rocks. The surface waters are generally very soft, but those derived from springs and deep wells are not unfrequently very hard. In sixty samples, the total hardness varied from 2.3° to 75°, and averaged 14.7°; the per- manent hardness ranged from 1.2° to 48.5°, and averaged 9.6° ; whilst the temporary hardness varied from 0° to 28.2°. AYater obtained from any stratum permeable to the foul liquids of sewers, middens, and cess-pits is always hard, and generally exhibits a large proportion of permanent hardness. The food of man and beast contains considerable quantities of lime, nearly the whole of which, in the adult, is discharged in the liquid and solid excrements. In 268 samples of shallow well water polluted by excrementitious matters to such an extent as to exhibit evidence of 10,000 parts and upwards of 500 INORGANIC CHEMISTRY. previous sewage or animal contamination, the total hardness ranged from 9.8° to 191°, and averaged 50.7° ; the permanent hardness varied from 3.8° to 164.3°, and averaged 31.7°; whilst the temporary hard- ness ranged from 0° to 49.2°, and averaged 19°. 10. Mineral Matters in Suspension.—The mineral matters in suspen- sion in potable water are almost invariably of an innocuous character, but they diminish or altogether destroy the transparency and brilliancy of the water, and impart a repulsive appearance, which often leads to the rejection of a wholesome water for a bright and sparkling though dangerous one. Slow filtration through sand is almost invariably effec- tive for the removal of visible suspended matters, but the washings of clay soils are very difficult to render bright by sand filtration ; and in all cases filtered water, if turbid previous to filtration, may always be shown, by suitable optical means, to be full of minute suspended parti- cles, although to unassisted vision it is perfectly clear and transparent. 11. Organic Matters in Suspension.—The organic matters in sus- pension in potable water possess not only all the objectionable qualities of similar matters of mineral origin, but in addition they are sometimes actively injurious, and they always promote the development of crowds of animalcules. Their presence in drinking water is therefore much more objectionable than is the occurrence of mineral matters in suspen- sion. Like the suspended mineral matters, the finely divided organic matters in suspension cannot be entirely removed by sand filtration. The Sixth Report of the Rivers Pollution Commission gives the result of the chemical examination of 1272 samples of potable water collected under the most widely different conditions, and comprehend- ing 81 samples of rain water, 372 samples of surface water, 419 sam- ples of shallow well water, 180 samples of deep well water, and 220 samples of spring water. This extended investigation of waters which have drained from the surface of, or percolated through the most im- portant geological formations of, Great Britain affords, the Commis- sioners say, a broad basis hitherto unattainable upon which to found conclusions as to the relative merits of potable waters from these various sources. The results of this research are quite conclusive as to the sources from which the best water for domestic purposes is to be ob- tained. They show that rain water contains the smallest proportion of total solid impurity, but by no means the smallest proportion of that most objectionable of impurities, organic matter. The rain drops con- centrate within themselves the organic dust and dirt diffused through vast volumes of atmospheric air, and everywhere visible when a ray of sunlight illuminates them. Rain water, collected from the roofs of houses at a distance from towns, carefully stored and filtered, may be made into a fairly good and wholesome potable water; but when it is col- lected from the surface of uncultivated land, allowed to subside in lakes or reservoirs, or filtered through sand, it becomes of good quality for domestic, and still more so for manufacturing purposes. Numerous large towns, both in England and Scotland, are supplied with water of this description. Non-calcareous strata are generally selected as gathering ground, and then the water is soft and well adapted both for washing and for almost all manufacturing operations. It is nearly POTABLE WATERS, always wholesome, but sometimes suffers in palatability by containing an excessive quantity of peaty matter in solution. This evil may be materially abated by the use of sand filters. Seeing that rapid filtration through a few feet of sand can materially improve the quality of surface water, by removing some of the organic impurity which it contains in solution, we are prepared to find a much greater improvement when the water is drawn from deep wells or springs, to which it could only gain access by slow natural percolation through a great thickness of porous rock or earth. Under such cir- cumstances, the powerful oxidizing influences of a porous and aerated soil are brought to bear upon the organic matter dissolved in the water. It is not, therefore, surprising to find that surface water should be almost, or eyen quite, exhaustively purified from such matter, by the natural intermittent filtration which transforms it into spring or deep well water. Mere exposure to the air, however, even if accompanied by violent agitation, is comparatively powerless for the removal of pollut- ing organic matter from water. Surface water, draining from cultivated land, is always more or less polluted with the organic matter of manure. Such water, of course, contributes very largely to rivers and streams which have already de- scended from their mountain or upland sources. Even when not con- taminated bv the actual admission into it of the sewage of towns and * O villages, it is not of suitable quality for domestic purposes, but when it is further polluted by excremental drainage, its use for drinking and cooking becomes fraught with great risk to health. Still more dan- gerous to health is the water drawn from shallow wells, no matter upon what geological formation they may be sunk, when they are situated, as is usually the case, near privies, drains, or cesspools. Many severe outbreaks of epidemic disease have been traced to the use of such water in villages and towns, and there is strong reason to believe that sporadic attacks of typhoid fever often occur in isolated country houses from the same cause. In respect of wholesomeness, palatability, and general fitness for drinking and cooking, waters may be classified in the following order of excellence: ( 1. Spring water. 1 -\r i , , , Wholesome. < 2. Deep well water. J Pa atl e* ( 3. Upland surface water. 1 Moderately pala- te 4, Stored rain water. j table. Suspicious. -< 5. Surface water from cultivated ( land. {6. River water to which sewage j- Palatable. gains access. 7. Shallow well water. Preference should always be given to spring and deep well water for purely domestic purposes, over even upland surface water—not only on account of the much greater intrinsic chemical purity and palatability of these waters, but also because their physical qualities render them peculiarly valuable for domestic supply. They are almost invariably 502 INORGANIC CHEMISTRY. clear, colorless, transparent, and brilliant—qualities which add greatly to their acceptability as beverages—whilst their uniformity of temperature throughout the year renders them cool and refreshing in summer and prevents them from freezing readily in winter. Such waters are of in- estimable value to communities, and their conservation and utilization are worthy of the greatest efforts of those who have the public health under their charge. The foregoing remarks have reference exclusively to the use of water for drinking and cooking—applications of paramount importance from a sanitary point of view; but a large proportion of the water supplied for domestic purposes is used for washing, whilst in many towns con- siderable volumes are used in manufactories. For all these latter pur- poses it is of the utmost importance that the water should be soft—a quality that is not always associated with wholesomeness and palata- bility. Classified according to softness, the waters from the various sources fall into the following order: 1. Rain water. 2. Upland surface water. 3. Surface water from cultivated land. 4. Polluted river water. 5. Spring water. 6. Deep well water. 7. Shallow well water. The interests of the laundress and of the manufacturer are thus evi- dently opposed to those of the householder, inasmuch as they lead to a preference for moderately palatable or even unwholesome water over that which is very palatable and wholesome. Most of the hard waters from springs and deep wells can, however, be easily and cheaply ren- dered soft, and the interests of the householder and manufacturer thus made identical. In Clark’s process of softening water with lime, the sanitary authorities of towns have at their disposal a method of render- ing hard water from springs or deep wells available for washing and manufacturing purposes, without diminishing either its palatability or its wholesorneness. The influence of geological formation upon the palatability and wholesomeness of water is very considerable. In the case of surface water this influence is to a great extent masked, or indeed often alto- gether annulled, by superficial deposits of vegetable matters, such as peat, upon the rocks; and thus, except in respect of hardness and saline constituents, unpolluted surface waters from the most widely different geological formations differ but little in the proportions of organic mat- ter which they contain, and consequently in their palatability and whole- someness. But when the water percolates or soaks through great thick- nesses of rock, its quality, when it subsequently appears as spring or deep well water, depends greatly upon the nature of the material through which it has passed. When the formation contains much soluble saline matter, the water becomes loaded with mineral impurities, as is fre- quently the case when it percolates through certain of the Carboniferous POTABLE WATERS. 503 rocks, the Lias, and the Saliferous Marls. When the rock is much fissured, or permeated by caverns or passages, like the Mountain Lime- stone, for instance, the effluent water differs but little from surface drain- age, and retains most of the organic impurities with which it was origi- nally charged. But when the rock is uniformly porous, like the Chalk, Oolite, Greensand, or New Red Sandstone, the organic matter, at first present in the water, is gradually oxidized and transformed into innocu- ous mineral compounds. In effecting this most desirable transforma- tion, and thus rendering the water sparkling, colorless, palatable, and wholesome, the following water-bearing strata are the most efficient: 3. Chalk. 2. Oolite. 3. Greensand. 4. Hastings Sand. 5. New Red and Conglomerate Sandstone. This is seen from the following table, in which the average composi- tion of unpolluted water from various sources is contrasted: INORGANIC CHEMISTRY. Average Composition of Unpolluted Potable Waters. Results of Analysis expressed in parts per 100,000. Description. Dissolved Matters. Number of samples analyzed. Total Solid Matters. Organic Carbon. Organic Nitrogen. Am- monia. Nitrogen as Nitrates and Nitrites. Total Combined Nitrogen. Previous Sewage or Animal Contami- nation. Chlorine. Hardness. Temporary. Permanent. Total. WATWR,, 2.95 0.070 0.015 0.029 0.003 0.042 42 0.22 0.3 39 Upland Surface Waters. From Non-Calcareous Strata. From Igneous Eocks, .... 5.15 0.278 0.033 0.001 0.002 0.035 O 1.13 0.1 2.0 2.1 18 From Metarnorphic, Cambrian, Silurian, and Devonian Eocks 5.12 0.293 0.024 0.002 0.000 0.031 3 0.92 0.3 2.5 25 81 From Yoredale and Millstone Grits and the Coal Measures, 8.75 0.373 0.037 0.003 0.010 0.050 6 1.05 0.4 4.3 4.7 47 From Lower London Tertiaries and Bagshot Beds, .... 8.40 0.379 0.048 0.004 0,007 0.058 O 2.06 0.3 3.5 3.8 3 POTABLE WATERS. 505 From Calcareous Strata. From Calcareous portions of Si- lurian and Devonian Rocks, 13.71 0.302 0.026 0.000 0.021 0.047 77 1.20 1.2 7.4 8.6 3 From Mountain Limestone, From Calcareous portions of the 17.07 0.370 0.047 0.001 0.011 0.059 26 1.24 5.7 7.0 12.7 7 Coal Measures, From the Lias, New Red Sand- stone, Conglomerate Sand- stone, and Magnesian Lime- 22.79 0.346 0.037 0.003 0.016 0.056 33 1.52 4.0 8.3 12.3 26 stone, From the Oolites, . . . » . Deep Were Waters. 18.80 0.286 0.042 o;oo2 0.010 0.054 4 1.49 7.6 6.5 14.1 9 17.46 0.326 0.025 0.004 0.042 0.070 130 1.55 6.6 5.8 12.4 1 In Devonian Rocks and Mill- stone Grit, 32.68 0.068 0.012 0.005 0.294 0.310 2,671 2.70 8.8 8.6 17.4 7 In the Coal Measures, . . . 83.10 0.119 0.034 0.044 0.207 0.278 2,243 18.05 15.1 20.6 35.7 9 In Magnesian Limestone, . . 61.14 0.076 0.030 0.000 1.426 1.456 13,937 4.31 16.9 26.9 43.8 3 In New Red Sandstone, . . . 30.63 0.036 0.014 0.003 0.717 0.734 6,895 2 94 7.4 10.5 17.9 28 In the Lias, 70.98 0.146 0.027 0.001 0.389 0.417 3,730 4.42 21.9 8 2 30.1 2 In the Oolites, In the Hastings Sand, Green- 33.60 0.037 0.010 0.022 0.625 . 0.654 6,118 2.69 13.8 6.8 20.6 5 sands, and Weald Clav, . . 45.20 0.068 0.014 0.016 0.196 0.223 1,864 5.38 16.8 10.5 27.3 20 In the Chalk In the Chalk below London 36.88 0.050 0.017 0.001 0.610 0.628 5,801 2.76 21.2 6.5 27.7 66 Clav, 78.08 0.093 0.028 0.048 0.068 0.135 797 15.02 9.7 8.7 18.4 13 In Thanet Sand and Drift, . . 53.84 0.113 0.020 0.072 0.116 0.202 1,517 6.32 14.4 7.6 22.0 4 Spring Waters. From Granite and Gneiss Rocks 5.94 0.042 0.008 0.001 0.106 0.115 846 1.69 0.4 2.6 3.0 8 From Silurian Rocks, . . ■. From Devonian Rocks and Old 12.33 0.051 0.014 0.001 0.178 0.192 1,587 1.84 1.5 5.3 6.8 15 Red Sandstone, 25.06 0.054 0.012 0.001 0.764 0.777 7,339 3.85 4.8 7.2 12.0 22 506 INORGANIC CHEMISTRY, Dissolved Matters. S Description. Nitrogen as Nitrates and Nitrites. Previous Hardness. a . CC o £ Total Solid Matters. Organic (Jar bon. Organic JSi itrogen. Am- monia. Total Combined Nitrogen. Sewage or Animal Contami- nation. Chlorine. Temporary. Permanent. Total. ?-< a v £ & as a s A From Mountain Limestone, 32.06 0.087 0.010 0.001 0.224 0.235 2,008 4.63 10.9 8.9 19.8 15 From Yoredale and Millstone Grits and the Coal Measures, 21.91 0.050 0.014 0.001 0.393 0.408 3,704 1.85 5.2 7.9 13.1 22 From Magnesian Limestone, . 66.52 0.058 0.038 0.002 1.686 1.726 16,560 3.40 24.9 34.8 59 7 1 From New Eed Sandstone, . . 28.69 0.065 0.017 0.001 0.330 0.349 3,047 2.19 8.1 10.7 188 15 From the Lias, 36.41 0.073 0.019 0.001 0.467 0.487 4,406 2.48 21.3 8.8 30.1 7 From the Oolites, From the Hastings Sand and Greensands, 30.33 0.043 0.011 0.001 0.402 0.414 3,730 1.55 18.2 6.2 244 35 30.05 0.053 0.012 0.000 0.326 0.338 2,941 2.98 13.6 6.6 20.2 19 From the Chalk, 29.84 0.044 0.010 0.001 0.382 0.392 3,511 2.45 18.1 5.5 23.6 30 From Fluvio-marine, Drift, and Gravel, 61.32 0.086 0.019 0.001 0.354 0.374 3,264 2.76 18.0 19.6 37.6 10 MAGNESIUM. 507 MAGNESIUM, Mg. Atomic weight = 24.4. Probable molecular weight = 24.4. ftp. gr. 1.743. Fuses at a red heat. Volatilizes at a red heat. Atomicity". Evidence of atomicity: Magnesic chloride, .... Mg"Cl,. Magnesic oxide, .... Mg"0. Magnesic hydrate, . . . . • . History.—Magnesic sulphate was described and its medicinal prop- erties pointed out by Grew at the close of the seventeenth century. The metal was first isolated by Davy. Occurrence.—The compounds of magnesium are widely distributed in nature. It occurs as carbonate in magnesite, COMgo" ; as dihydric magnesic sulphate in kieserite, SOHoJVIgo", and Epsom salts, SOHo2Mg0",60H2; as silicate in enstatite, SiOMgo", in ophite or noble serpentine, Si2OMgo"3, in talc, SigOgMgo"*, and other minerals. In combination with other bases, as double salts, it occurs in enormous quantities as dolomite, a carbonate of isomorphous calcium and mag- nesium, mCOCao",nCOMgo" ;* as kainite, S02Ko(^Q|Mg^,3oH2; as carnallite, MgCl 2,KC1,60H2; and in a great number of silicates. The sulphate and chloride are also found in saline springs and in sea-water. It occurs in small quantities in the animal and vegetable kingdoms: thus, in the bones of animals and in the seeds of plants. Preparation.—Magnesium may be obtained by the electrolysis of the fused chloride, but is more conveniently prepared by the action of sodium on the chloride. A mixture of 6 parts of fused magnesic chloride, 1 part of powdered fluorspar, 1 part of a mixture of sodic and potassic chloride in equal molecular proportions, and 1 part of sodium in small pieces, is thrown into a red-hot crucible, which is quickly closed. As soon as the reaction is over the crucible is removed from the fire and allowed to cool to below redness, after which the contents are stirred with a pipe-stem, in order to cause the globules of magnesium to unite. When quite cold, the solidified slag is broken up, and the magnesium removed. Magnesium is now manufactured on a large scale. Properties.—Magnesium is a silver-white lustrous metal, of sp. gr. 1.743. The pure metal preserves its lustre in dry air, but becomes covered with a film of oxide when exposed to the action of moisture. At a higher temperature it maybe pressed into the form of wire or ribbon, an operation which must be performed with exclusion of air. It fuses at a red heat, and may be distilled in a current of hydrogen. Magnesium wire or ribbon may be ignited at the flame of a candle, and burns with an intensely brilliant white light very rich in chemically active rays, a property which has led to its use in photography. Pure * See p. 65. 508 INORGANIC CHEMISTRY. magnesium does not decompose water even at 100° C. (212° F.). Dilute acids dissolve it with violent evolution of hydrogen. Unlike zinc it does not evolve hydrogen when heated with solutions of caustic alkalies. This is due to the fact that the magnesic hydrate, which would be formed, is not soluble in the alkali. Magnesium gives off hydrogen when heated with solutions of ammonia salts, the magnesium dissolving in the form of a double salt of magnesium and ammonium. Uses.—Except for laboratory purposes, magnesium is employed ex- clusively in the production of the magnesium light. Besides its appli- cation in photography already referred to, the magnesium light has been used in signalling. The light has been seen at sea at a distance of 28 miles. COMPOUNDS OF MAGNESIUM WITH THE HALOGENS. Magnesic chloride, MgCl2.—This compound occurs in sea-water and in salt deposits. It is formed when the metal, the oxide, or the carbonate, is dissolved in hydrochloric acid. On concentrating the solu- tion, the chloride is deposited in monoclinic crystals of the formula MgCl2,60[T2, which when heated give off their water of crystallization, but at the same time are partially resolved into magnesic oxide and hydrochloric acid. In order to obtain the anhydrous salt in a state of purity, 12 parts of the commercial oxide are dissolved in hydrochloric acid ; the solution is shaken with an excess of oxide, in order to pre- cipitate alumina and iron, and, after filtering, evaporated to dryness with 27 parts of ammonic chloride. The resulting magnesic ammonic chloride is carefully heated to expel the water of crystallization, and is afterwards ignited in a platinum crucible, until fumes of ammonic chloride cease to be given off, and the whole has fused to a clear liquid. The anhydrous chloride solidifies on cooling to a colorless laminated crystal- line mass with a lustrous fracture. It deliquesces when exposed to moist air, dissolves in water with evolution,of heat, and is also readily soluble in alcohol. It volatilizes at a bright red heat. Magnesic chloride is employed in dressing cotton goods.—Magnesic chloride combines with magnesic oxide to form oxychlorides of varying composition. If strongly ignited magnesia be made into a paste with a concentrated solution of magnesic chloride, the mixture solidifies in the course of a few hours to a solid mass, sufficiently hard to be polished. Magnesic polemic chloride, MgCl2,KC1,60H2, occurs native as carnallite in large de- posits at Stassfurt, and is frequently deposited from the last mother-liquors of sea-water and brine-springs. It forms colorless rhombic prisms, which deliquesce on exposure to'the air. On heating, the water of crystallization is expelled without decomposition of the salt, and the anhydrous salt fuses at a red heat. Anhydrous carnallite may be employed in the preparation of magnesium by means of sodium. Magnesic ammonic chloride, MgCl2,N 11,01,601f, is deposited in small rhombic crystals from mixed solutions of magnesic and ammonic chlorides. It is soluble in 6 parts of water. Magnesic calcic chloride, 2MgCl2,CaCl2,120H2, occurs native in deliquescent masses as taehydrite, at Stassfurt. Magnesic bromide, Mgßr2, occurs in sea-water and in saline springs. A solution of magnesia in hydrohromic acid deposits needle-shaped crystals of the formula COMPOUNDS OF MAGNESIUM. 509 Mgßr2,60H.2, which when heated behave like the aquate of magnesic chloride. Mag- netic bromide forms double salts with the alkaline bromides. Magnesic. iodide, Mgl2, occurs in sea-water and in saline springs, and may be pre- pared by dissolving magnesia in hydriodic acid. It forms deliquescent crystals which readily decompose when heated. Magnesic fluoride, MgF,,, occurs native as selldite in colorless quadratic crystals. It is obtained as a white insoluble powder by digesting magnesia with hydrofluoric acid. By fusion with common salt this powder is converted into crystals having the same form as sellai'te. Magnesic sodie fluoride, MgF.2,NaF.—This salt is obtained in insoluble, cubical crys- tals by fusing magnesic chloride with a large excess of sodic fluoride and cooling slowly. It is also formed by digesting magnesia with a solution of sodic fluoride. COMPOUNDS OF MAGNESIUM WITH OXYGEN AND HYDROXYL. Magnesic oxide. Magnesia, . MgO. Mg=C). Magnesic hydrate, .... Mg’ H 02.o2. ll—O—Mg—O—H. Magnesic oxide {Magnesia), MgO, occurs native as peridase, a rare mineral found at Monte Sornma, near Naples. The natural com- pound forms regular octahedra, generally of a greenish color, due to the presence of ferrous oxide. It is formed when magnesium burns in the air. It is usually prepared by prolonged ignition of the carbonate, and is thus obtained as a bulky white powder known as magnesia usta, or calcined magnesia. It is insoluble in water. It possesses a sp. gr. of 3.07, but when very strongly ignited, itssp. gr. is increased to 3.61, the substance becoming at the same time crystalline. By heating magnesia in a current of gaseous hydrochloric acid, it is obtained in crystals identical with those of peridase. It fuses in the oxyhydrogen flame. Magnesia is employed in medicine. Magnesic hydrate, Mgllcq, occurs native as brucite in colorless laminated masses. By the addition of sodic or potassic hydrate to solutions of magnesia salts, a gelatinous precipitate is obtained, which, after drying at 100° C. (212° F.), consists of pure magnesic hydrate. It forms a white powder, almost insoluble in water, in solutions of sodic and potassic hydrate, and in aqueous ammonia; readily soluble in solutions of ammonia salts. It absorbs carbonic anhydride from the air. At a low red heat it is decomposed into magnesia and water. The magnesia formed at this low temperature has the property of again taking up water, with evolution of heat, to form the hydrate. OXY-SALTS OF MAGNESIUM. ' STO Magnesic nitrate, jj.Q2Mgo//, GOT!,, forms deliquescentnaonoclinic prisms, soluble in half their weight of cold water, soluble also in alcohol. The water of crystallization cannot be completely expelled without partial decomposition of the salt. Magnesic caebonate, COMgo", occurs native as magnesite, some- times in rhombohedral crystals isomorphous with those of calcite, more frequently massive. The native carbonate generally contains iron and manganese. By precipitating a hot solution of a magnesia salt with potassic or sodic carbonate, and boiling the precipitate with 510 INORGANIC CHEMISTRY. water as long as any acid carbonate is dissolved; a basic magnesic r CHo2(OMgHo) Mgo carbonate of the formula C{Ho6Mgo''2(OMgHo)2 = -< CHo2 Mgo v CHo2(OMgHo) is obtained. This compound also occurs native as hydromagnesite in acicular monoclinic crystals or amorphous masses. By precipitating a magnesia salt with a large excess of sodic carbonate, and boiling with the solution until the precipitate becomes crystalline, a carbonate is ( CHo2(OMgHo) obtained having the formula C2Ho4Mgo"(OMgHo)9== < Mgo" (CHo2(OMgHo) The pharmaceutical preparation known as magnesia alba is a mixture of various complex carbonates of magnesia, obtained by precipitating soluble magnesia salts with sodic carbonate, and varies in composition according to the mode of preparation. It forms a very light, bulky white powder. When magnesia alba is suspended in water and the liquid saturated with carbonic anhydride, the powder dissolves with formation of an acid carbonate. On allowing the solution to stand, carbonic anhydride gradually escapes, and a salt of the formula COMgo" ,30H2 separates in fine needles, which when exposed to the air part with their water of crystallization and become opaque. At a very low temperature crystals of a salt having the formula COMgo",5O H2 are deposited. When the solution of the acid carbonate is evaporated to dryness, anhydrous magnesic carbonate remains as a fine powder, which under the microscope exhibits rhombic forms corresponding to those of arragonite. But if the solution be heated under pressure to 300° C. (572° F.), at the same time allowing the carbonic anhydride to escape gradually, the anhydrous carbonate is obtained in minute rhom- bohedra, identical with those of native magnesite. Magnesic carbonate is, therefore, isodimorphous with calcic carbonate. When the salt COMgo",3OH2 is boiled with water it gives off carbonic anhydride, and is converted into a basic salt, whilst when heated in the dry state to 300° C. (572° F.), it is entirely decomposed into carbonic anhydride and magnesia. Native magnesite is not altered by boiling with water, and does not evolve carbonic anhydride at 300° C. It is also only slowly attacked by acids in the cold. Magnesic dipotassic carbonate, ColoMS°"40H” is formed when magnesia alba is digested with a solution of hydric potassic carbonate for some time at a temperature of 60-70° C. It forms small rhombic prisms, which are decomposed by water. Magnesic diammonic carbonate, qq y Mgo//,40 H3, separates in colorless rhombic crystals, when a solution of magnesia salt is added to a large excess of a mixed solu- tion of ammonic carbonate and free ammonia. It is almost insoluble in water. Magnesic calcic carbonate.—This compound, which, as the mineral dolomite, forms entire mountain ranges, is not a true double salt, but an isomorphous mixture of mag- nesic and calcic carbonates in varying proportions. As bitter-spar it occurs crystallized in rhombohedra. It is employed in the preparation of magnesia alba. Magnesic sulphate, S02Mgo".—A dihydric magnesic sulphate, SOHo2Mgo", occurs in layers in the salt-beds at Stassfurt as the COMPOUNDS OF MAGNESIUM. 511 mineral hieserite. It generally forms granular masses, and is almost insoluble in water, but when allowed to remain long in contact with water gradually dissolves with formation of the salt SOHo2Mg0",60H2. The latter compound occurs native as epsomite or Epsom salt, both solid as an efflorescence of fibrous crystals, and in solution in many mineral ■waters. Magnesic sulphate is deposited from hot concentrated solu- tions in large transparent rhombic prisms of the above formula SOHo2Mg0",60H2, isomorphous with the corresponding aquates of zincic and nickelous sulphates; but a salt having the same composition is sometimes deposited from cold supersaturated solutions in monoclinic forms isomorphous with those of ferrous sulphate, SOHo2Feo//,60H2, with which magnesic sulphate also crystallizes in varying proportions. Above 70° C. (158° F.) it separates from its solutions in monoclinic crystals of the formula SOHo2Mgo//,50H2; at 0° C, (32° F.) a salt having the composition SOHo2Mgo'/,110H2 is deposited. Epsom salt is soluble in four-fifths of its weight of water, still more soluble in water at 100° C. (212° F.), insoluble in alcohol. It has an unpleasant bitter taste. When heated, it fuses in its water of crystallization, which is given off below 150° C. (302° F.), leaving the salt SOHo2Mgo"; this in turn, when heated above 200° C. (392° F.), parts with the ele- ments of water, and is converted into the anhydrous sulphate S02Mgo", which fuses at a red heat without decomposition. The acid salt, ,dihy- dric magnesic disulphate, gQ2| |||Mgo//, crystallizes in six-sided tables from a solution of the anhydrous normal salt in concentrated sulphuric acid. It is instantly decomposed by water. Large quantities of Epsom salt were formerly prepared from dolomite by treating the mineral with sulphuric acid and then separating the soluble Epsom salt from the in- soluble calcic sulphate; but at the present day nearly all the Epsom salt is obtained from the kieserite of Stassfurt. The crude kieserite from the upper salt layer, or Abraumsalz, is placed in sieves suspended in water. Sodic and magnesic chloride dissolve, the kieserite disintegrates and falls through the meshes of the sieve in a fine powder, whilst earthy impurities are retained by the sieve. The powdered kieserite is then pressed, wdiile wet, into wrooden moulds, where it speedily solidifies to a hard mass, owing to the combination of the water with a portion of the kieserite to form Epsom salt, which binds the powder together. The mass is then powdered, and is either brought into the market as kieserite, or is converted first into Epsom salt. Kieserite is employed as a manure, and in the preparation of potassic and sodic sulphate. Epsom salt is used as a purgative. It is also employed in dressing cotton goods and in aniline dyeing.—Magnesic sulphate forms double salts with the alkaline sulphates. Magnesic dipotassic disulphate, so!koms°"’60H» and magnesic diammonic disulphate, ,, £/)rr S02AmoMS° >bU±d2’ are deposited, from mixed solutions of magnesic sulphate with potassic or with ammonic sulphate, in monoclinic crystals. The potassium salt occurs native at Stassfurt as the mineral schonite. 512 INORGANIC CHEMISTRY. • POMgc/ • Magnesic wthophosphate, pQ\[go//Mgo//, occurs in bones and the seeds of plants. It is obtained as a white pulverulent precipitate when a solution of trisodic orthophos- phate is added to a solution of a magnesia salt. It is almost insoluble in water, but dissolves readily in dilute acids. A double phosphate and fluoride of magnesium having the formula P( )Mgo// 'j occurs in raonoclinic crystals as the mineral wagnerite.—Hydric magnesic orthophosphate, POHoMgo//,70H2, is deposited in hex- agonal needles when dilute solutions of magnesic sulphate and hydric disodic phos- phate are mixed. When concentrated solutions are employed, the salt is obtained as an amorphous precipitate which becomes crystalline on standing. It is sparingly soluble in water, and is decomposed by boiling into the normal salt which is deposited and free phosphoric acid which remains in solution.—Tetrahydric magnesic diortho- phosphate has not been prepared. Magnesic potassic orthophosphate, POKoMgo//,60H2, and magnesic sodic orthophos- phate, PONaoMgo//,90H2, are obtained in minute crystals by adding to solutions of potassic or sodic dihydric orthophosphate the requisite quantity of magnesia. Both salts are decomposed by washing with water. Magnesic ammonic orthophosphate, POAmoMgo//,60H2, separates from putrid urine, and is frequently a constituent of urinary calculi; it occurs also in guano in rhombic crystals as guanite or struvite. It separates as a crystalline powder when hydric disodic phosphate is added to a mixed solution of a magnesia salt with an ammonia salt and free ammonia. In dilute solutions the precipitate is not formed till after some time; it then attaches itself in small crystals to the sides of the vessel, particularly to parts which have been rubbed with a glass rod in stirring the liquid. It is almost totally insoluble in water, especially in water containing ammonia. When ignited, it is con- verted into magnesic pyrophosphate: 2POAmoMgo// = P203Mgo//2 + 2NHS + OH2. Magnesic amnionic Magnesic Ammonia. Water, orthophosphate. pyrophosphate. Magnesic ammonic phosphate is employed in the estimation both of magnesia and of phosphoric acid. Magnesic arsenate, //, and hydric magnesic arsenate, AsOHoMgO//,70H2; are prepared like the corresponding phosphates, and form white precipitates, almost insoluble in water, readily soluble in acids. Tetrahydric magnesic diarsenate is soluble in water, but uncrystallizable. Magnesic ammonic arsenate, AsOAmoMgo//,60TT2, is prepared like the correspond- ing phosphate, which it resembles in almost every particular. When dried at 100° C. (212° F.), it parts with of its water of crystallization, yielding the salt— (AsOAmoMgo//)2,0H2, or As2OHo2Amo2Mgo//2. The rest of the water cannot be expelled without partial decomposition of the salt, a portion of the ammonia being driven off and a portion of tbe arsenic acid undergoing reduction to arsenious acid. This water is therefore probably to be regarded as water of constitution, as represented in the second of the above formulae. Magnesic ammonic arsenate is employed in the estimation of arsenic acid. Magnesic borates.—When magnesia and boric anhydride are fused together at a very high temperature, and the fused mass is allowed to cool slowly, nacreous crystals of trimagnesic diorthoborate, gj^°°//Mgo//, are formed. The same salt with' 9 aq. is obtained by precipitating a solution of a magnesia salt with borax. No precipitate is formed in the cold, but on boiling the solution the salt, BM|o"M*0"9OH>- separates as an amorphous white powder, which dissolves again on cooling. A double dioctoborate and chloride of the formula occurs native, in large crystals belonging to the regular system, as boracite and massive as stassfurtite. The COMPOUNDS OF MAGNESIUM. 513 same compound may be obtained artificially in the crystallized form by fusing mag- nesia orthoborate with boric anhydride, magnesic chloride, and sodic chloride, allow- ing the mass to cool slowly, and treating with dilute hydrochloric acid, when the crystals of boracite remain undissolved. Magnesic silicates.—A number of magnesic silicates occur in nature as minerals. Peridote is a dimagnesic silicate (orthosilicate) of the for- mula SiMgo"2. It occurs in rhombic crystals, generally green-colored, owing to the presence of iron, or in granular masses. Enstatite is mono- magnesic silicate (metasilicate) SiOMgo", It forms monoclinic crystals, which generally contain iron. The following natural magnesic silicates are also known : Ophite or noble serpentine. Trimagnesic disilicate, Si2OMgo"3. Meerschaum. Tetrahydric dirnagnesic trisilicate, . Si302Ho4Mgo"2. Steatite. Trimagnesic tetrasilicate, Si405Mgo"3. Talc. Tetramagnesic pentasilicate, Si506Mgo"4. Numerous natural compound silicates of magnesium with other metals are also known. COMPOUNDS OF MAGNESIUM WITH SULPHUR AND WITH HYDE OS ULPHYL. Magnesic. sulphide, MgS//.—Magnesium is not acted upon by sulphur at the boiling- point of the latter; but when the metal is heated to redness in the vapor of sulphur, magnesic sulphide is formed. It may also be prepared by passing the vapor of car- bonic disulphide over red-hot magnesia. It forms a gray or brown, hard, brittle slag. Water decomposes it, yielding a mixture of magnesic hydrate and sulphhydrate. When an excess of sodic sulphide is added to the solution of a magnesium salt, the precipitate which is formed consists not of magnesic sulphide, but of magnesic hydrate. Magnesic sulphhydrate, MgJls.,, has not been prepared pure. It may be obtained in solution by passing sulphuretted hydrogen into water in which magnesia is suspended. On evaporating the solution, sulphuretted hydrogen is given off and magnesia re- mains. COMPOUNDS OF MAGNESIUM WITH NITROGEN AND WITH BORON. Magnesic nitride, N2Mg3, is prepared by heating magnesium in nitrogen or gaseous ammonia. The product is an amorphous greenish-yellow mass, which in contact with water, or even in moist air, is decomposed with formation of ammonia and magnesia : N2Mgs + SOH2 = 2NH3 + 3MgO. Magnesic nitride. Water. Ammonia. Magnesic oxide. Magnesic boride, B2Mgs, is formed when magnesium is heated with amorphous boron in a closed crucible. It can be obtained, mixed with magnesia, by heating boric an- hydride with magnesium. In contact with hydrochloric acid, it evolves boric hydride, 8113, mixed, however, with a large excess of hydrogen. COMPOUND OF MAGNESIUM WITH SILICON. Magnesic silicide, SiMg2.—For the method of preparing this compound, see Silicic hydride, p. 311. General Properties and Reactions of the Compounds of Magnesium.—The salts of magnesium with colorless acids are colorless. 514 INORGANIC CHEMISTRY. The soluble salts have a bitter taste. The hydrates of the alkalies and of baryta precipitate from solutions of magnesium salts gelatinous mag- nesia hydrate, insoluble in an excess of the precipitant. When salts of ammonia are present in sufficient quantity, no precipitation occurs with the above reagents in the cold, owing to the formation of double salts of ammonium and magnesium, which are not decomposed at ordinary temperatures. For the same reasons the salts of magnesium are only imperfectly precipitated by ammonia. Sodic carbonate precipitates a basic carbonate; ammonium salts prevent the precipitation. Ammo- nia phosphate gives a white crystalline precipitate of magnesia am- nionic phosphate, POAmoMgo//,60H2, very sparingly soluble in water, insoluble in aqueous ammonia. Magnesium compounds im- part no coloration to the non-luminous flame. The spark spectrum of magnesium displays characteristic lines in the green, coincident with lines of the solar spectrum. ZINC, Zn. Atomic iveight = 65.3. Molecular weight = 65.3. Molecular and atomic volume I I I. 1 litre of zinc vapor iveighs 32.65 criths. Sp. gr. 6.8f07.2. Fuses at 420° C. (788° F,). Boils at 1040° C.(1904°F.). Atomicity". Evidence of atomicity: Zincic chloride, Zn"Cl2. Zincic oxide, Zn"0. Zincic hydrate, Zn"Ho2. History.—The ores of zinc were employed by the ancients in the preparation of brass, which they obtained by melting copper with these ores; but zinc was not recognized as a distinct metal till the sixteenth century. Occurrence.—Zinc is asserted to have been found native near Mel- bourne, in Australia. It occurs as oxide (ZnO) in red zinc ; as sul- phide (ZnS") in the mineral zinc-blende; as carbonate (COZno'') in calamine, or zinc-spar ; as silicate (SiZno"2,OH2) in siliceous calamine, or zinc-glass; and as double oxides of the general formula /R//202Ro" in franklinite (/Fe//202Zn0//) and gahnite or zinc-spinelle (/Al//2G2Zno//). Extraction.—Zinc is obtained from the carbonate, less frequently from the sulphide. Siliceous calamine, red zinc and franklinite are also worked. The first operation in the process of extracting the zinc con- sists in roasting the ore in order to convert it into oxide. In the case of the carbonate this is effected simply by expulsion of carbonic anhy- dride; the sulphide is oxidized by the oxygen of the air with evolution of sulphurous anhydride. In roasting the sulphide it is necessary to avoid the formation of zincic sulphate, as this salt would, in the subse- quent reducing process, be reconverted into sulphide and thus lost. The roasted ore is then mixed with half its weight of powdered coal, and distilled from fire-clay tubes or from muffles placed in a furnace. At first a finely divided powder known as zinc-dust, and consisting of ZINC, 515 a mixture of zinc with zincic oxide, frequently also accompanied by cadmiujßi, passes over. Afterwards the liquid metal distils over and is collected in iron receivers, from which it is removed from time to time during the distillation and cast into plates. Commercial zinc is seldom pure. It generally contains lead, iron, carbon, and sometimes arsenic and cadmium. It may be obtained almost pure by redistillation from clay retorts, the first portions of the distillate, which contain arsenic and cadmium, being rejected, and the operation being interrupted before all the zinc has passed over. The iron, lead, and other less volatile impurities remain in the retort. In order to prepare perfectly pure zinc, the crude metal is dissolved in sulphuric acid, and sulphuretted hydrogen is passed through the acid solution of zincic sulphate. In this way lead, cadmium, and arsenic are precipitated as sulphides. The filtered solution is boiled to expel sulphuretted hydrogen, and the zinc is precipitated as carbonate by the addition of sodic carbonate. The zincic carbonate is converted into oxide by ignition, and the oxide is reduced by distillation from a porce- lain retort with pure charcoal prepared from sugar. Any iron which may have been contained in the purified carbonate remains in the retort. Properties.—Zinc is a white lustrous metal, with a slightly bluish tinge. It has a crystalline, somewhat laminar fracture, and may be obtained in crystals by fusing the metal, allowing it to partially solidify, and then pouring off the still liquid portion. It generally crystallizes in flat hexagonal pyramids, but occasionally exhibits forms belonging to the regular system, especially when it contains traces of copper. At ordinary temperatures it is brittle; between 100° C. (212° F.), and 150° C. (302° F.), it is so malleable and ductile that it may be rolled into plates and drawn into wire; at 205° C, (401° F.) it again becomes so brittle that it may be powdered in a mortar. It may be distilled at a bright red heat. In dry air it preserves its lustre at ordinary tem- peratures ; in moist air it becomes covered with a thin coating of basic carbonate, which preserves it from further action. Reactions.—l. When heated in air, zinc inflames, emitting a brilliant bluish light, and giving off clouds of zincic oxide. The combustion of zinc is best shown by pressing thin zinc turnings into the form of a cyl- inder ; this, when ignited at a flame, readily burns. 2. Pure zinc is very slowly attacked by dilute sulphuric and hydro- chloric acids, but the addition of a few drops of platinic chloride to the liquid causes the zinc to dissolve rapidly, with evolution of hydrogen, the finely divided platinum, which is deposited on the zinc, forming with the latter a voltaic couple. For the same reason commercial zinc, which always contains traces of electronegative metals, is rapidly dis- solved by dilute acids. In cold nitric acid the metal dissolves without evolution of gas, the nascent hydrogen being employed in reducinganother portion of the acid to ammonia; in hot nitric acid it dissolves with evo- lution of nitric oxide, nitrous oxide, and free nitrogen, whilst ammonia is also formed. When zinc is acted upon by hot dilute sulphuric acid, or by concentrated sulphuric acid even in the cold, sulphuretted hydro- 516 INORGANIC CHEMISTRY. gen, formed by the redaction of a portion of the acid, is mixed with the hydrogen which is given off: 5S02Ho2 + 4Zn = 4S02Zno" + SH2 + 40H2. Sulphuric Zincic Sulphuretted Water, acid. sulphate. hydrogen. 3. Zinc also dissolves in warm solutions of potassic, sodic, and ammonia hydrate, with evolution of hydrogen and formation of a double oxide: 20KH + Zn = ZnKo2 + H2. Potassic Dipotassic hydrate. zincic oxide. 4. It slowly decomposes aqueous vapor at 100° C. (212° F.): 20H2 + Zn = ZnHo2 + H2. Water. Zincic hydrate. Uses.—Zinc, in the form of sheets, is employed for roofing and other purposes in which lightness and the power of resisting the action of the weather are required. In order to preserve iron from rust, the metal is sometimes coated with zinc, in which condition it is known as galvanized iron. Zinc is used in the preparation of plates for voltaic batteries. The finely divided powder obtained in the distillation of zinc, and known as zinc-dust, is frequently employed as a reducing agent in organic chemistry, many oxygenated organic substances, which are unacted upon by all other reducing agents, parting with their oxy- gen when distilled with zinc-dust. The use of zinc in the desilveriza- tion of lead has already been described (p. 448). COMPOUNDS OF ZINC WITH THE HALOGENS. Zincic chloride, ZnCl2. Molecular weight = 136.3. Molecular volume I I I. Zinc foil inflames spontaneously at ordinary temperatures in chlorine gas and burns, forming zincic chloride. The chloride may also be ob- tained by dissolving zinc in hydrochloric acid, evaporating the solu- tion, and distilling the residue; or by distilling anhydrous zincic sul- phate with sodic or calcic chloride. Zincic chloride is a white very deliquescent mass. At ordinary temperatures it is soft like wax; it fuses somewhat above 100° C, (212° F.); at a higher temperature it sublimes in white needles, and may be distilled without decomposition. It is very soluble both in water and in alcohol. The concentrated so- lution is powerfully caustic: it destroys vegetable fibre, and therefore cannot be filtered through paper. When a little hydrochloric acid is COMPOUNDS OF ZINC. 517 added to a syrupy solution of zincic chloride, the liquid deposits deli- quescent octahedra of the monaquate, ZnCl2,0H2. The solution of zincic chloride cannot be evaporated without decomposition : hydro- chloric acid is given off, and an oxychloride of zinc remains. Oxychlo- rides of varying composition, consisting of mixtures of ZnHoCl and ZnHo2, are also obtained by heating the concentrated solution of zincic chloride with zincic oxide, and then adding water, when the oxychlo- rides are precipitated. In the same way, by boiling the solution of the chloride with the requisite quantity of oxide, a plastic mass is obtained which, like the mixture of magnesic chloride and magnesia (p. 508) speedily becomes quite hard.—Owing to its great affinity for water, zincic chloride frequently abstracts the elements of water from organic substances, thus producing new compounds, a property of which appli- cation is made in organic research. It is also used as a caustic in med- icine, for which purpose it is cast into sticks.—Zincic chloride forms crystalline, deliquescent double salts with the chlorides of the alkalies: for example, dipotassic zincic chloride, ZnCl2,2KCJ; disodic zincic chlo- ride, ZnCl2,2NaCI. Zincic bromide,' Znßr2, is prepared like the chloride. It crystallizes in very deli- quescent prisms, is readily fusible, and may be sublimed in white needles. Zincic iodide, Znb.—Zinc filings and iodine, when heated together, unite to form the iodide. Zincic iodide is readily fusible, and sublimes in colorless needles. From a concentrated aqueous solution it crystallizes in deliquescent regular octahedra. The concentrated solution takes up oxygen from the air, with liberation of iodine. In like manner, when zincic iodide is heated in air, iodine is given off, and zincic oxide is produced. Zincic iodide combines with the alkaline iodides to form double salts. Zincic fluoride, ZnF2, is obtained by dissolving zincic oxide in aqueous hydrofluoric acid. On evaporation, the solution deposits small, shining, rhombic octahedra of the formula ZnF2,4OH2, sparingly soluble in water. Zincic fluoride forms crystalline double salts with potassic and other fluorides. The potassium salt has the formula ZnF2,2KF. Zincic silicofluoride, SiZnF6,60t12, forms very soluble hexagonal crystals. COMPOUNDS OF ZINC WITH OXYGEN AND HYDROXYL. Zincic oxide, .... ZnO. Zn=o. Zincic hydrate, . . . ZnPIo2. H—O—Zn—O—H. Zincic oxide, ZnO, occurs native, sometimes in hexagonal crystals, more frequently in granular masses, as red zinc ore, the color being due to an admixture of manganese. It is formed when zinc is- burnt in air (p. 515). On a large scale it is prepared by distilling zinc from earth- enware retorts, allowing the zinc vapor to burn as it issues from the retort, and passing the products of combustion through chambers in which the oxide collects. It may also be prepared by igniting the basic carbonate obtained by precipitating the solution of a zinc salt with an al- kaline carbonate. The zincic oxide prepared by combustion is a white floc- culent substance, and was known to the alchemists as lana philosophica; that obtained by the ignition of the carbonate is an amorphous powder. The artificial oxide may be obtained in the hexagonal forms of the 518 INORGANIC CHEMISTRY. natural variety by igniting it strongly in a current of oxygen. Crystals of zincic oxide are also sometimes found in the cooler parts of the muffles of the zinc furnaces. Zincic oxide has a sp. gr. of 5.6. It is insoluble in water, readily soluble in acids. When heated it assumes a yellow color, changing to white again on cooling. When heated in the oxy- hydrogen flame it does not fuse, but emits a brilliant light, and on cooling continues to phosphoresce for some time in the dark. Zincic oxide is employed as a very permanent white pigment under the name of zinc white. As the sulphide of zinc is also white, zinc white does not change color when exposed to sulphurous exhalations, possessing in this respect a marked superiority over white lead. Zincic hydrate, ZnHo2, is precipitated as a white amorphous powder by the addition of sodic orpotassic hydrate, or ammonia, to the solution of a zinc salt. The precipitate is insoluble in water, but soluble in an excess of the precipitant. It may be obtained in a crystalline form by immersing a sheet of zinc, round which a copper wire has been wound, in a solution of the hydrate in ammonia ; rhombic prisms of the hy- drate are formed upon the surface of the zinc. A saturated solution of the hydrate in caustic potash deposits on standing regular octahedra of the formula ZnHo2,0H2. When heated, zincic hydrate is readily decomposed into zincic oxide and water. NO Zincic nitrate, N0:Z"0",6OH2) separates from a concentrated solution of the oxide in nitric acid in deliquescent, colorless, four-sided prisms. It is readily soluble in water and in alcohol. At 36° C. (96.8° F.) it fuses in its water of crystallization, and, when heated to 100° C. (212° F.), parts with water and nitric acid, yielding a basic salt. OXY-SALTS OF ZINC. Zincic carhonaie, COZno//, occurs native in translucent rhombohedra as calamine. The native carbonate is rarely pure, a portion of the zinc being generally replaced by calcium, iron, and other metals isomorphous with zinc. Zincic carbonate is precipitated when hydric potassic car- bonate is added to the solution of a zinc salt. Normal potassic and sodic carbonates precipitate basic zincic carbonates of variable compo- sition. The basic precipitate is insoluble in water and in solutions of potassic and sodic carbonate, but soluble in ammonic carbonate. Zincic sulphate (White vitriol), S02Zno/;, is prepared on a large scale by roasting the native sulphide and extracting the mass with water, but is most readily obtained pure by dissolving zinc in sulphuric acid. At ordinary temperatures it crystallizes in large transparent rhombic prisms of the formula isomorphous with Epsom salt (p. 511), soluble in two-thirds of their weight of water at ordinary temperatures, in one-sixth of their weight of boiling water ; insoluble in alcohol. The crystals effloresce slowly in air, and, when heated to 100° C. (212° F.), orexposed in vacuo over sulphuric acid, part with 6 aq., leaving the salt SOHr^Zno7', which is converted at a temper- ature of 240° C. (464° F.) into anhydrous zincic sulphate (So2Zno") and water. At temperatures above 40° C. (104° F.) solutions of zincic COMPOUNDS OF ZINC. 519 sulphate deposit monoclinic crystals having the formula SOHo2Zno//,- 50H2, also isomorphous with the corresponding magnesium salt. When the anhydrous salt is heated to a high temperature it gives off sulphur- ous anhydride and oxygen, yielding a basic salt, a hot saturated solu- tion of which deposits on cooling lustrous laminae of the formula SO(OZnHo)4. The same compound may be obtained by boiling a solution of zincic sulphate with zincic oxide. At a white heat the anhydrous sulphate is converted into zincic oxide. Zincic sulphate forms double sulphates with the sulphates of the alkalies, zincic dipo- tassic disulphate, gQ2|^|)>Zno'G 0II2, and zincic diammonic disulphate, gQ2^™^Zno",6oH2, which are isomorphous with and closely resemble the corresponding magnesium compounds. Mixed solutions of zincic and raagnesic sulphates deposit crystals containing the two salts in variable proportions.—Zincic sulphate is employed in medicine and in calico printing. Zincic orthophosphate.—The normal or irizincic salt, POZno//Zno//’40IT2’ *s f° when hydric disodic phosphate is added to a solution of a zinc salt. It is a white precipitate, which, when deposited from cold solutions, is gelatinous, but becomes crystalline on standing or on heating.—-The acid phosphates have not been prepared. Zincic silicate.—A dizincic silicate, SiZno//2, occurs native in hexagonal prisms as willemite. It may be obtained artificially in the crystallized form by passing silicic fluoride over zincic oxide heated almost to whiteness, or by the action of zincic fluoride on silicic anhydride.—The same compound with 1 aq., SiZno//2,OH2—perhaps to be regarded as SiO(OZnHo)2—occurs in rhombic crystals as the mineral zinc glass or siliceous calamine. COMPOUNDS OF ZINC WITH SULPHUR. Zincic sulphide, ZnS", occurs native as zinc blende, either crystal- lized in forms belonging to the regular system, or massive. The color of the mineral varies from a pale yellow, in the purer specimens, to a brown or black in the massive variety, due to the presence of iron and other impurities. Zincic sulphide is occasionally found in hexagonal prisms as the mineral icurtzite. It is obtained as a white amorphous precipitate when sulphuretted hydrogen is passed through a solution of zincic acetate. From neutral solutions of zinc salts with mineral acids the zinc is only partially precipitated by sulphuretted hydrogen, and in acid solutions no precipitate is produced. All zinc salts, however, are completely precipitated by the addition of alkaline sulphides or sulph- hydrates to their solutions. The precipitated zincic sulphide is insolu- ble in water and in acetic acid, but readily soluble in mineral acids with evolution of sulphuretted hydrogen, Zincic sulphide is difficultly fusi- ble. When the amorphous sulphide is heated to a very high tempera- ture in a current of sulphuretted hydrogen, or sulphurous anhydride, it sublimes in colorless hexagonal crystals identical with those of wurtzite. Trizincie dipotassic tetrasulphide, SIZn3K2. K—S—Zn—S—Zn—S—Zn—S—K. This compound is obtained by fusing together 1 part of zincic sulphide, 24 parts of potassic carbonate, and 24 parts of sulphur, at a red heat for ten minutes. On extract- 520 INORGANIC CHEMISTRY. ing the cooled mass with water, the double sulphide remains in the form of colorless transparent laminae, which may be boiled with water without decomposition.—The cor- responding sodium compound S4Zn3Na2, may be obtained in a similar manner, and forms a pale flesh-colored crystalline powder. Zineic pentasulphide, S5Zn, is obtained as a white precipitate by the addition of po- tassic pentasulphide to the neutral solution of a zinc salt. It assumes a pale yellow color on drying, and, when heated with exclusion of air, gives off sulphur, and is con- verted into the monosulphide. COMPOUNDS OF ZINC WITH THE PENTAD ELEMENTS. Zincic nitride, N2Zn3.—When zinc ethyl (see Organic Chemistry) Is acted upon by gaseous ammonia, ethylic hydride is evolved, and zinc diamine is formed: Zn(C2HS)2 + 2NHS - Zn(NH2)a + 2{g*H*. Zinc ethyl. Ammonia. Zinc Ethylic diamine. hydride. The zinc diamine thus obtained is a white amorphous powder, which is decomposed by water with formation of ammonia and zincic hydrate; Zn NHj)2 + 20H2 = ZnHo3 + 2NH3. Zinc Water. Zincie Ammonia, diamine. hydrate. When zinc diamine is heated to low redness in absence of air, ammonia is evolved, and zincic nitride remains as a green powder: 3Zn(NH2)2 = N2Zn3 + 4NIT3. Zinc Zincic Ammonia, diamine. nitride. In contact with water zincic nitride is decomposed with great evolution of heat, yield ing ammonia and zincic oxide. Zincic phosphide, P2Zn3, is prepared by heating finely divided zinc in the vapor of phosphorus. An impure compound is obtained by heating a mixture of phosphoric anhydride, zincic oxide, and charcoal. Zincic phosphide forms a steel-gray metallic mass, which dissolves in hydrochloric acid with evolution of phosphoretted hydrogen. Zincic arsenide, As2Zn3, is formed with incandescence when zinc and arsenic are heated together in the proportions required by the formula. It is a gray, brittle me- tallic mass, which, when acted upon by dilute hydrochloric acid, evolves pure arseni- uretted hydrogen (p. 367). Zincic antimonide, Sb2Zn3, is obtained as a white crystalline metallic mass by fusing together 57 parts of antimony and 43 parts of zinc. By allowing the fused compound partially to solidify, and pouring off the still liquid portion, it may be obtained in well- formed hexagonal prisms. When treated with hydrochloric acid, it evolves a mixture of hydrogen and antimoniuretted hydrogen (p. 380).—A dizincic diantimonide of the formula /Sb//2Zn2, crystallizing in rhombic octahedra, is prepared by fusing 68.5 parts of antimony with 31.6 parts of zinc. General Properties and Reactions of the Compounds of Zinc.—The salts of zinc are colorless when the constituent acid is so. They have an astringent metallic taste, and are poisonous. From their solutions caustic alkalies and ammonia precipitate white zincic hydrate, soluble in excess of the precipitant. Alkaline carbonates precipitate a basic carbonate, soluble in amnionic carbonate, but not in excess of po- tassic or sodic carbonate. Baric carbonate does not precipitate solutions of zinc salts. Sulphuretted hydrogen gives no precipitate in acid solu- tions, except in the case of salts of organic acids in solutions acidulated COMPOUNDS OF BERYLLIUM. 521 with these acids; ammonic sulphide precipitates white hydrated zincic sulphide. Hotassicferrocyanide gives a white precipitate of zincic fer- rocyanide. Heated on charcoal in the reducing flame of the blowpipe, zinc compounds yield a characteristic incrustation of zincic oxide, yel- low while hot, white when cold. If this incrustation be moistened with cobaltous nitrate and again heated, it assumes a fine green color (Rin- mann’s green). The salts of zinc do not color the non-lurainous flame. The spark spectrum of zinc shows characteristic lines in the red and in the blue. BERYLLIUM, Be. (Sometimes termed Glucinum, symbol G.) Atomic weight =9. Probable molecular weight = 9. Sp.gr. 2.1. Fuses at a red heat. Atomicity ". Evidence of atomicity: Beryllic chloride, . . . Be"Cl9. Beryl lie oxide, . , . Beryllic hydrate, . . . . . Be"0. . . . Be"Ho2, History.—Beryllic oxide was prepared by Vauquelin in 1798. Wohler first isolated the metal in 1828. Occurrence.—Beryllium occurs in combination in a few rare minerals. Beryl, a native double silicate of beryllium and aluminium of the for- mula Si606(/Al///206)vlßeo,/3, is the most abundant source of the beryl- lium compounds. This mineral crystallizes in hexagonal prisms, gen- erally opaque, and of a greenish tint. The precious stone emerald is a transparent beryl of a brilliant green color; bluish-green specimens, when transparent, are known as aquamarine, and are also employed as gems. The mineral phenacite is a silicate of beryllium having the for- mula Sißeo'^. ibeparation.—Metallic beryllium is prepared by passing the vapor of beryllic chloride along with a current of hydrogen over heated so- dium, and afterwards fusing the metal thus obtained in a crucible under sodic chloride. Properties.—Beryllium is a lustrous silver-white malleable metal of sp. gr. 2.1. It fuses below the melting point of silver. When fused in air it becomes covered with a thin coating of oxide, which checks further oxidation ; but when heated in a finely divided state it inflames, burning with a very brilliant light. It does not decompose water, even at 100° C. (212° F.). Dilute hydrochloric acid dissolves it readily in the cold, with evolution of hydrogen, but dilute sulphuric acid does not attack it till heated, whilst nitric acid, even when hot and concentrated, acts upon it only very slowly. It is not attacked by ammonia, but dissolves readily in caustic potash with evolution of hydrogen. COMPOUNDS OF BERYLLIUM WITH THE HALOGENS. Beryllic chloride, BeCk.—Molecular weight = 80. Molecular vol- ume I I I.—The anhydrous chloride is obtained in lustrous, colorless, needle-shaped crystals by passing chlorine over a heated mixture of 522 INORGANIC CHEMISTRY. beryllic oxide and charcoal. It is readily fusible and volatile. The crystals deliquesce rapidly when exposed to air, and, when thrown into water, dissolve with a hissing sound, evolving heat. The aqueous solution, which may also be obtained by dissolving the oxide in hydro- chloric acid, deposits, by spontaneous evaporation over sulphuric acid, colorless crystals of the formula BeCl2,40H2, from which the water of crystallization cannot be expelled without decomposition of the salt. Beryllic bromide, Beßr2, and Beryllic iodide, Bel 2, are both obtained in the form of colorless needles by the direct union of their elements. Beryllic fluoride, BelY—The anhydrous salt is not known. The solution of beryllic hydrate in hydrofluoric acid deposits on evaporation an amorphous mass, which when further heated gives off water and hydrofluoric acid, being partially converted into oxide. It forms double fluorides with the fluorides of the alkali metals. COMPOUNDS OF BERYLLIUM WITH OXYGEN AND HYDROXYL. Beryllic oxide, Beryllia, . BeO. Be=o. Beryllic hydrate, .... BeHo2. H—O—Be—O—H. Beryllic oxide or Beryllia, BeO.—This oxide is prepared from the mineral beryl, a beryllic aluminic silicate (p. 521). The finely powdered mineral is fused with three parts of anhydrous potassic carbonate, and the cooled mass is treated with concentrated sul- phuric acid, the excess of acid being expelled by heating. On extracting with water, the sulphates of beryllium, aluminium, and potassium dissolve, whilst the silica remains and may be filtered off'. . The' solution is evaporated until a crust begins to form on the surface. On standing, the greater portion of the alumina crystal- S02Ko—I Of) I lizes out as potash alum, Grk2 the beryllia S02Ko—J remaining in solution. A fresh crop of alum crystals may be obtained by the further exaporation of the mother liquor from the first crop. The filtered liquid from the second crop of crystals is then poured into an excess of a warm solution of amnionic carbonate, and the whole is allowed to remain for some days in a stoppered bottle, agitating from time to time. The precipitate, consisting of alumina and ferric oxide, is filtered off, and the beryllia is precipitated from the solution, either as basic carbonate by protracted boiling, or as hydrate by acidulating with hydrochloric acid and afterwards rendering alkaline with am- monia. By ignition the carbonate or hydrate is converted into oxide. Thus prepared beryllia forms a white bulky amorphous powder of sp. gr. 3.08, resembling magnesia in appearance. It is insoluble in water, and, after being strongly ignited, does not dissolve in dilute acids. Like magnesia, it becomes crystalline by exposure to a very intense heat. Beryllic hydrate, BeHo2, is obtained as a gelatinous precipitate when ammonia is added to a solution of a beryllium salt. After drying at 100° C. it forms a bulky white COMPOUNDS OF BERYLLIUM. 523 powder, which at a higher temperature is converted into the oxide. It is insoluble in water, soluble in solutions of caustic potash, caustic soda, and ammonic carbonate, but insoluble in ammonia. If the solution in caustic potash be diluted and boiled, the beryl lie hydrate is reprecipitated. From the solution in ammonic carbonate a precipi- tate of a basic beryllic carbonate separates on boiling. Beryllic hydrate dissolves in a boiling solution of ammonic chloride with formation of beryllic chloride and with lib- eration of ammonia. OXY-BALTB OF BERYLLIUM. Beryllic nitrate, |“’Reo//,30H2, forms deliquescent crystals, readily soluble in alcohol. At a temperature of 250° C. it is completely converted into oxide. Beryllic carbonate.—The precipitate produced in solutions of beryllium salts by alka- line carbonates is a basic beryllic carbonate of the formula ClIo(OBeHo)3,30H2. This salt dissolves in water containing carbonic anhydride, and the solution, when evapo- rated over sulphuric acid in an atmosphere of carbonic anhydride, deposits crystals of the normal carbonate. COBeo//,40F12- These, on exposure to the air, spontaneously part with carbonic anhydride and are re-converted into the basic salt. Beryllic sulphate, SOIIo2Beo",30IT2, crystallizes from aqueous solu- tions in quadratic octahedra, which are soluble in their own weight of water at ordinary temperatures, and effloresce on exposure to the air. The water of crystallization is expelled at 110° C., leaving the salt SOHo2Beo". This salt is stable at 150° C., but at a higher temperature the water of constitution is expelled and the anhydrous salt SCXBeo" re- mains. At a red heat the anhydrous salt is converted into beryllia. From solutions containing free sulphuric acid, beryllic sulphate crystallizes in large efflorescent monoclinic prisms of the formula isomorphous with those of Epsom salt.* Mixed solutions of beryllic and magnesic sulphates deposit crystals containing the two metals in variable proportions. Beryllic orthophosphate.—A hydric beryllic phosphate, POHoBeo//,SOH2, is obtained as a white amorphous precipitate when hydric disodic phosphate is added to the solution of a beryllium salt. When the sodic phosphate is added to a solution containing beryllic nitrate and ammonic chloride, the triple salt, disodic diammonic beryllic phos- phate, P-AfNaoaf NII/))2Beo//,70!l2, is precipitated as a white crystalline powder. Beryllic silicate, Sißeo//2, occurs native in hexagonal crystals as phenacite. Beryllic sulphide, BeS", is formed as a gray infusible mass when beryllium is heated in sulphur vapor. Alkaline sulphides precipitate only beryllic hydrate from solutions of beryllium salts. COMPOUND OF BERYLLIUM WITH SULPHUR. General Properties and Reactions of the Compounds of Beryllium.—The salts of beryllium with colorless acids are colorless; they have a sweet, slightly astringent taste and an acid reaction. Caustic alkalies, ammonia, and ammonic sulphide precipitate white flocculent beryllic hydrate, in the case of the last precipitant with evolution of sulphuretted hydrogen. The precipitate is soluble in excess of caustic alkali, but not in excess of ammonia. Beryllic hydrate is soluble in ammonic carbonate, and may thus be separated from alumina, along * Marignac, however, doubts whether these salts are really isomorphous. 524 INORGANIC CHEMISTRY. with which it is usually precipitated in analysis. Beryllium salts do not color the non-luminous flame. The spark spectrum contains two characteristic lines in the blue. CHAPTER XXXIY. DYAD ELEMENTS. Section 111. CADMIUM, Cd. Atomic weight =ll2. Molecular weight 112. Molecular and atomic volume I I I. 1 litre of cadmium vapor weighs 56 criths. Sp. gr. 8.6. Fuses at 320° C. (608° F.). Boils at 860° C. (1580° F.). Atomi- city Evidence of atomicity : Caclmic chloride, Cd"Cl2, Cadmic oxide, CcT'O. History.—Cadmium was discovered indepeudently and almost simul- taneously by Stromeyer and by Hermann in 1817. Occurrence.—Cadmium occurs in small quantities in many zinc ores. A fibrous zinc blende found at Przibram in Bohemia contains as much as from 2to 3 per cent, of cadmium. The rare mineral greenochite is a sulphide of cadmium (OdS"). Preparation.—In the process of extracting zinc from ores containing cadmium, the latter metal distils over first, and is for the most part oxidized by the air in the receivers. By distilling these first portions with powdered coal at as low a temperature as possible, cadmium is obtained almost pure. In order to purify it thoroughly, it is dissolved in dilute sulphuric or hydrochloric acid and precipitated from the acid solution by sulphuretted hydrogen, the zinc remaining in solution. The cadmic sulphide is redissolved in concentrated hydrochloric acid, and the cadmium is precipitated from the solution by an excess of ammonic carbonate, which dissolves any arsenic and copper that may be present. The cadmic carbonate is converted by ignition into oxide, which by dis- tillation with a tenth of its weight of powdered coal yields the pure metal. Properties.—Cadmium is a white lustrous metal, with a fibrous frac- ture. When pure it is very malleable and ductile. It loses its lustre by exposure to the air, and when heated in air burns, giving off a brown smoke of cadmic oxide. Dilute sulphuric and hydrochloric acids dis- solve it slowly with evolution of hydrogen. Nitric acid rapidly dis- solves it. Zinc precipitates it in the metallic form from the solution of its salts. COMPOUNDS OF CADMIUM. 525 COMPOUNDS OF CADMIUM WITH THE HALOGENS. Cadmic chloride, CdCl2.—A solution of the metal or of the oxide in hydrochloric acid deposits on evaporation colorless prisms of the composition CdCl2,20H2, which effloresce when exposed to the air. The water of crystallization may be expelled by heat without decompo- sition of the salt. The anhydrous chloride fuses below a red heat, and at a higher temperature may be sublimed in colorless laminae. One hundred parts of water at 20° C. dissolve 141 parts of the anhydrous salt, and the solubility scarcely varies with the temperature. It forms a number of crystalline double chlorides with the alkaline and many other chlorides. Cadmic bromide, Cdßr2, is prepared by digesting cadmium with bromine and water. On evaporation the solution yields efflorescent acicular crystals of the formula Cdßr2- 40 H2, which on heating become anhydrous. At a higher temperature the salt fuses and sublimes in colorless laminae. It forms double bromides with the bromides of the alkalies and alkaline earths. Cadmic iodide, Cdl 2, is prepared like the bromide. It crystallizes from water in fusible hexagonal plates. When heated it is decomposed with evolution of iodine. One hundred parts of water at 20° C. (68° F.) dissolve 93 parts of the salt; at 100° C. (212° F.), 133 parts. It is also soluble in alcohol. If forms numerous double iodides with the iodides of other metals. Cadmic iodide is employed in photography. COMPOUNDS OF CAD3IIUM WITH OXYGEN AND HYDROXYL. Cadraic oxide, . . . CdO. Cd=o. Cadmic hydrate, . . CdHo3. H—O—Cd—O—H. Cadmic oxide, CdO, may be prepared like the oxide of zinc by the combustion of the metal. It is thus obtained as a brown amorphous powder. When cadmic nitrate is ignited the oxide remains in the form of microscopic octahedra, which by reflected light appear blue-black, by transmitted light brown. It is insoluble in water, readily soluble in acids. It is infusible even at a white heat. When heated on charcoal before the blowpipe, it is reduced, the metal at the same time volatilizing and burning with formation of a brown incrustation of cadmic oxide on the charcoal. Cadmic hydrate, CdHoz, is obtained by precipitating the solution of a cadmium salt with potassic or sodic hydrate, and drying the precipitate at 100° C. (212° F.). It forms a white powder, insoluble in water and in solutions of potassic and sodic hydrate; readily soluble in ammonia and in acids. It absorbs carbonic anhydride from the air. At 300° C. (672° F.) it is converted into oxide. OXY-SALTS OF CADMIUM. NO Cadmic nitrate, odo//,40H2, crystallizes in deliquescent prisms, soluble in alcohol. Cadmic carbonate.—A precipitate approximating in composition to that of the normal 526 INORGANIC CHEMISTRY. salt, COCdo//, is obtained by adding in the cold a solution of a cadmium salt to an excess of an alkaline carbonate. The precipitate formed at a higher temperature, or with a smaller quantity of alkaline carbonate, is a basic salt of varying composition. Cndmie sulphate, S02Cdo", is deposited from its solutions by sponta- neous evaporation at ordinary temperatures in large colorless mono- clinic crystals of the formula 3SO2Cd0",80H2. A boiling solution containing an excess of sulphuric acid deposits warty crystals of a salt SO2Cdo//,0H2. The anhydrous salt is soluble in less than twice its weight of water at ordinary temperatures; somewhat more soluble at 100° C, (212° F.). The normal salt is converted by heating into a basic compound of the formula S02(OCdHo)2, sparingly soluble in water and crystallizing in pearly scales. Cadmic sulphate is employed in med- icine. COMPOUND OF CADMIUM WITH SULPHUR. Cadmic sulphide, CdS/r, occurs native in yellow hexagonal prisms as the rare mineral greenockite. It is obtained as an amorphous powder of a pure yellow color when a solution of a cadmium salt is precipitated with sulphuretted hydrogen or with an alkaline sulphide. It is soluble in concentrated nitric and hydrochloric acids, and in hot dilute sulphuric acid. It fuses at a white heat, and solidifies on cooling in micaceous scales. By fusing the precipitated sulphide with potassic carbonate and sulphur, extracting the cooled mass with water, or by passing the vapor of sulphur over cadmic oxide heated to the highest possible temperature, cadmic sulphide may be obtained in hexagonal crystals. General Properties and Reactions of the Compounds of Cadmium.—The salts of cadmium with colorless acids are color- less. Caustic alkalies precipitate from solutions of the salts white cadmic hydrate, insoluble in excess of the precipitant. Ammonia gives the same precipitate, readily soluble however in excess. Ammonic car- bonate precipitates cadmic carbonate, insoluble in excess. Sulphuretted hydrogen precipitates from a hydrochloric acid solution yellow cadmic sulphide, insoluble in ammonic sulphide and in potassic cyanide, but soluble in hot dilute sulphuric acid. Heated on charcoal before the blowpipe, cadmium compounds give a brown incrustation of cadmic oxide. Cadmium compounds do not color the non-luminous flame. The spark spectrum displays characteristic lines in the red, green, and blue. 527 MERCURY. MERCURY, Hg. Atomic iceight = 200. Molecular iceight 200. Molecular and atomic volume I I I. 1 litre of mercury vapor weighs 100 criths. Sp. gr. 13.59. Fuses at —39.5° C. (—39.i° F.). Boils at 357.25° C. (676.05° F.) (Regnault). Atomicity ", also a pseudo-monad. Evidence of ato- micity : Mercuric chloride, . . . . . . . Hg"Cl2. Mercuric oxide, Mercurous chloride, . ... . . . 'Hg'2CI2. Mercurous oxide, . . . 'Hgr'aO. History.—Mercury has been known from almost the earliest historic times. Occurrence.—Mercury is found native in minute globules dissemi- nated through its ores. It occurs in combination as chloride and iodide, and also with gold and silver in the form of amalgams. Its most abundant ore is mercuric sulphide or cinnabar. The most important mines are those of Idria in Carniola, Alraaden in Spain, Napa Valley in California, and at Wolfsstein and Landsberg in the Bavarian Palati- nate. Extraction.—At Idria the ore—a mixture of cinnabar with earthy matters—is placed on the top of a perforated arch, under which the furnace is situated. After closing the aperture through which the ore has been introduced the furnace is lighted. The flame, along with an excess of air which is allowed to enter by openings constructed for that purpose, plays through the perforations of the arch upon the ore, oxidizing the sulphur to sulphurous anhydried, and volatilizing the mercury. The products of combustion pass through stone chambers, in which the mercury condenses, and thence into a tower, through which a stream of water trickles, removing the last traces of mercury from the escaping gases. At Almaden, the mercury vapor, instead of passing into stone chambers, is condensed in a series of stoneware bottles termed aludels, open both at top and bottom, and so arranged that the neck of each flts into the bottom of the next. A furnace, in which from 50 to 60 tons of ore can be distilled in one operation, can be filled and the charge worked off in a day ; but four or five days must be allowed to elapse before the furnace is suffi- ciently cool to be recharged. In order to obviate this loss of time, a continuous process has been devised in which the ore, along with char- coal, is introduced from time to time at the top of the furnace whilst the ashes are withdrawn at the bottom. In the Bavarian Palatinate the ore is mixed with lime and distilled from iron retorts. Mercury passes over, and a mixture of calcic sul- phide and sulphate remains. In Bohemia the ore is distilled with smithy-scales. Mercury obtained by any of the above processes is freed from me- chanical impurities by filtering through linen. It is generally sent into the market in iron bottles. 528 INORGANIC CHEMISTRY. Preparation of Pure Mercury.—Commercial mercury is generally contaminated with small quantities of foreign metals which it holds in solution. The presence of these impurities is manifested by a diminu- tion of the fluidity of the mercury, accompanied by a tendency to ad- here to glass or porcelain ; a globule of pure mercury runs rapidly and coherently over a clean inclined surface of porcelain; but when the mercury is impure the globule becomes considerably elongated in its course, and generally leaves behind it on the porcelain a dark-colored track of oxide in which traces of the metal are retained. Mercury may be freed from these impurities by distillation, the surface of the metal being covered during the operation with a thick layer of iron- filings to diminish spirting. A very pure product may be obtained by conducting the distillation in a Sprengel vacuum. Mercury may also be purified by agitating it with dilute nitric acid, or by leaving it in shallow vessels in contact with the acid, when the impurities are dis- solved first. Mercury is also very effectively purified by leaving it for several days under a layer of concentrated sulphuric acid. Pure mer- cury ought to leave no residue when dissolved in nitric acid, evaporated, and ignited. Properties.—Mercury is a silver-white, very lustrous metal. It is liquid at ordinary temperatures, but solidifies at —39.5° C. to a tin- white, malleable, and sectile mass, crystallizing in regular octahedra. It contracts during solidification. Mercury volatilizes sensibly at ordi- nary temperatures: a piece of gold leaf suspended in a closed vessel over mercury becomes in course of time white and silvery, owing to the absorption of the mercurial vapor by the gold. Mercury boils at 357,25° C. (675.05° F.), yielding a colorless vapor. Pure mercury un- dergoes scarcely any alteration in air at ordinary temperatures, though a very thin film of mercurous oxide is formed on the surface; but at a temperature near to its boiling point it gradually absorbs oxygen with formation of red mercuric oxide. Hydrochloric acid, even when hot and concentrated, is without action upon mercury. Sulphuric acid does not attack it in the cold ; but the hot concentrated acid dissolves it with evolution of sulphurous anhydride. When the metal is present in excess, and the temperature is not allowed to rise to the boiling point of the mixture, a mercurous salt is formed ; an excess of acid leads to the formation of a mercuric salt. Cold dilute nitric acid dissolves it, yielding mercurous nitrate; when an excess of the metal is boiled with the dilute acid a basic mercurous nitrate is obtained. Hot concen- trated nitric acid in excess dissolves it with evolution of nitric oxide and formation of mercuric nitrate. When a rapid stream of water from a tap is directed from a height of three or four inches upon the surface of a large mass of mercury, bubbles of mercury are formed and float on the surface of the water. These transmit blue light through the thin metallic film, and deposit on bursting a minute globule of mercury. When mercury is triturated with sugar, grease, and various other substances, it is obtained in a very finely divided state, the union of the particles of the metal being prevented by the interposition of the foreign substance. This process is known as the “ deadening ”of the mercury. In the case of gray mercurial ointment, which is prepared MERCURY. 529 by this method, the mercury forms nearly uniform globules having a diameter of 0.001 to 0.004 mm. The vapor of mercury when inhaled acts as a*.poison, producing salivation. The finely divided mercury when taken internally has a similar action ; but liquid mercury has been swallowed without noticeable ill effects. Uses.—Mercury is invaluable to the physicist and the chemist. Many important physical observations could not have been made without the aid of apparatus of which mercury forms an essential part. The chemist employs mercury in collecting and measuring gases which are soluble in water, and also for the preparation of the mercurial compounds. It is further used in silvering mirrors, in extracting gold and silver from their ores by the amalgamation process, and in medicine. Amalgams. The alloys of the various metals with mercury are known as amal- gams. Some amalgams are formed by the direct union of their con- stituents, the combination either taking place spontaneously at ordinary temperatures, or requiring the aid of heat. Gold, silver, tin, sodium, and many other metals may be thus directly amalgamated. In other cases an indirect process must be resorted to. If the metal is more electro-positive than mercury, it may frequently be amalgamated by immersing it in the solution of a salt of mercury; in this way an amal- gam of copper may be prepared. Other indirect methods of amalga- mation are: the electrolysis of a solution of the metal, employing mercury as the negative electrode, and the action of an amalgam of so- dium upon the solution. Potassium amalgam.—Potassium and mercury combine with con- siderable rise of temperature, but without incandescence. The amalgam is solid when it contains 1 part of potassium to 96 parts of mercury, but liquid when the proportion of mercury rises to 140 parts. The solid amalgam crystallizes in cubes. Sodium amalgam.—Sodium and mercury combine violently at ordi- nary temperatures, the process being attended with a hissing noise and vivid incandescence. An amalgam containing 100 parts of mercury to 1 of sodium is viscid; with 80 parts of mercury, pasty; with 40 parts, solid; and with 30 parts, hard. The amalgams of potassium and sodium when brought in contact with water evolve hydrogen. Sodium amalgam is employed as a re- ducing agent in organic chemistry (p. 425). It is also used in the ex- traction of gold and silver (p. 449). Iron amalgam.—lron may be amalgamated by rubbing its clean surface with sodium amalgam. Ammonium amalgam.—See p. 235, Copper amalgam.— When copper is immersed in a solution of nitrate of mercury, the mercury is deposited on the surface of the copper. By treating finely divided or precipitated copper in this way, and then triturating it under warm water with the requisite quantity of mercury, an amalgam of copper may be obtained. Copper amalgam containing 30 per cent, of copper is hard enough to scratch tin, but has the re- 530 INORGANIC CHEMISTRY. markable property of becoming soft and plastic by heating to 100° C. (212° F.) and kneading in a mortar, recovering its hardness in the course of a few hours. As it has the same density in the soft as in the hard state, it may be employed to stop cavities, which it exactly fills on solidifying. In this way it has been used for stopping teeth. Cadmium amalgam possesses the same plastic properties as the above, and is also employed in dentistry. Tin amalgam.—Tin dissolves in mercury with absorption of heat. According to the relative proportions the amalgam is either liquid, or solid and crystalline. Tin amalgam is employed in silvering mirrors. COMPOUNDS OF MERCURY WITH THE HALOGENS. a. Mercurous Compounds. Mercurous chloride, Calomel, 'Hg' 2C12.—Molecular weight = 471.—This compound occurs in lustrous quadratic crystals or crystal- line crusts as the rare mineral horn-quicksilver. It is precipitated by the addition of hydrochloric acid or a soluble chloride to a solution of mercurous nitrate. It is also precipitated when a solution of mercuric chloride is saturated with sulphurous anhydride and the liquid is warmed to 70° C. (158° F.) or 80° C. (176° F.). Calomel is generally pre- pared, however, in the dry way by subliming together 4 parts of mer- curic chloride with 3 parts of metallic mercury. The sublimation is performed in a cast-iron cylinder, and the calomel vapor is passed into the upper part of a large brick chamber, where it condenses in a tine powder, as in the process of preparing flowers of sulphur. The pro- duct must be thoroughly washed with large quantities of warm water in order to remove any unaltered mercuric chloride. When the vapor is allowed to condense on a cold surface, the mercurous chloride is ob- tained as a fibrous crystalline mass of sp. gr. 7.1. Mercurous chloride assumes a gray tint under the action of light, owing to the separation of metallic mercury. When heated it sublimes without fusing. It possesses a vapor density only half of that required by the formula 'Hg' 2C12; but investigation has shown that the supposed vapor of calomel consists of a mixture of the vapors of mercuric chloride and mercury, which recombine on cooling: /Hg'2Cl2 = HgCl2 + Hg. Mercurous Mercuric 2 vols. chloride. chloride, 2 vols. It is insoluble in water, in alcohol, and in dilute acids in the cold. By boiling with hydrochloric acid it is converted into mercuric chloride, which dissolves, and metallic mercury. In contact with caustic potash it blackens, owing to the formation of mercurous oxide.—Calomel is much used in medicine. Mercurous bromide, is prepared by precipitating a solution of mercurous nitrate with hydrobromic acid or with a soluble bromide, and also by subliming a mix- COMPOUNDS OF MERCURY. 531 ture of mercnric bromide and metallic mercury. It closely resembles the chloride in its properties, and, like that compound, possesses a vapor density only half of that re- quired bv,its formula. Mercurous iodide, /Hg/212, may be obtained by triturating 10 parts of mercury with 6.7 parts of iodine, adding sufficient alcohol to moisten the mass. The product must be washed with alcohol in order to remove any mercuric iodide. It forms a yellowish- green powder, which fuses at 290° C. (554° F.), yielding a black liquid. It sublimes below this temperature, and by careful sublimation may be obtained in yellow rhombic crystals. These, when heated to 70° C. (158° F.), assume a red color, which deepens as the temperature rises, till at 220° C. (428° F.) it attains to a deep garnet-red. This change of color is not accompanied by any change in composition, and the crystals recover their original color on cooling. When quickly heated, mercurous iodide is decomposed into mercuric iodide and metallic mercury, and the same change takes place gradually at ordinary temperatures. It is only sparingly soluble in water. In contact with a solution of potassic iodide it is decomposed into mercuric iodide, which dissolves with formation of potassic mercuric iodide, and metallic mercury. Mercurous fluoride, /Hg/2F2, is prepared by dissolving freshly precipitated mercurous carbonate in hydrofluoric acid and evaporating the solution. It forms small yellow crystals, which are partially decomposed by pure water with separation of mercurous oxide. When the dry fluoride is heated in a glass vessel, mercury sublimes and the glass is corroded. b. Mercuric Compounds. Mercuric chloride, Corrosive sublimate, HgCl2.—Molecular weight = 271. Molecular volume I i I.—This compound is formed when mercury is heated in an excess of chlorine; also when mercuric oxide is dissolved in hydrochloric acid, or mercury in aqua-regia. It is generally prepared by heating a mixture of mercuric sulphate and common salt ;>the mercuric chloride sublimes and condenses as a color- less, transparent, crystalline mass in the upper part of the vessel. A small quantity of manganic dioxide is added to the mercuric sulphate in order to oxidize any mercurous salt which may be present. Mercuric chloride crystallizes from its aqueous solution in long colorless rhombic prisms, having a sp. gr. of 5.4. It fuses at 265° C. (509° F.) and boils at 295° C. (563° F.). It is soluble in from 14 to 15 parts of water at ordinary temperatures, in 2 parts of water at 100° C. (212° F.). It also dissolves in 3 parts of alcohol and in 4 parts of ether. Mercuric chloride is exceedingly stable, dissolving in concentrated nitric acid and sulphuric acid without decomposition. On heating the sulphuric acid solution the mercnric chloride sublimes out of the concentrated acid unchanged. Mercuric chloride is a violent poison. It is employed in medicine and as an antiseptic for anatomical preparations.—Mercuric chloride forms the following crystallizable double salts with the alka- line chlorides : HgCl2KC1,0H2; HgCl2,2KC1,0H2; 2HgCU,KCI,2OH2‘; HgCl2,2NaCI; 2HgCl2,2NaC1,30H2; Hgol2,NH4CI; HgCl2,2NH4C1,0H2 {sal alembroth). It also combines with hydrochloric acid to form the crystallized compounds HgCl2,HCI and 2HgCl2,HCI, both of which, on exposure to air, effloresce and part with the whole of their hydrochloric acid.—By boiling a solution of mercuric chloride with mercuric oxide, or by adding to the solution a quantity of caustic alkali insufficient for 532 INORGANIC CHEMISTRY. complete precipitation, various oxychlorides of mercury, many of which are crystallizable, may be obtained. The compound Hg403Cl2 = Cl—Hg—O—Hg—O—Hg—O—Hg—Cl forms lustrous golden-yellow scales. Mercuric bromide, Hgßr2, is obtained by the direct union of mercury with an excess of bromine. It is less soluble than the chloride, and crystallizes from water in lustrous laminae, from alcohol in rhombic needles or prisms. It sublimes readily. Mercuric iodide, Hgl2, is prepared by triturating 10 parts of mercury with 18 parts of iodine, adding sufficient alcohol to moisten the mass; or by mixing solutions of 10 parts of mercuric chloride and 12J parts of potassic iodide. The product obtained by the direct com- bination of iodine with mercury is a brilliant red crystalline powder; that prepared by precipitation is at first of a pure yellow, but speedily becomes red on standing. It is insoluble in water, readily soluble in alcohol or in solutions of potassic iodide and of mercuric chloride, yield- ing colorless solutions. From the alcoholic solution it is deposited in red quadratic octahedra. When mercuric iodide is heated to 150° C. (302° F.) it suddenly changes its color to yellow; at 238° C. (460.4° F.) it fuses to a yellow liquid and volatilizes in yellow lustrous rhombic crystals, which on standing or sometimes even during the process of cooling, are converted into aggregates of the red crystals. This change into the red modification, which is accompanied by evolution of heat, takes place instantaneously on scratching the yellow crystals. Mercuric iodide yields with potassic and ammonic iodides double salts of the formulae 2(HgI2,K1),30H2 and 2(HgI2,NH41),30H2, which crystal- lize in yellow prisms. Merc,uric fluoride, HgF2.—Mercuric oxide added to hydrofluoric acid is converted into a yellow crystalline powder consisting of mercuric oxyfluoride, HgFHo, and the solution yields on evaporation orange-colored crystals of the same compound. An excess of water decomposes the oxyfluoride, even in the cold, into hydrofluoric acid and mercuric oxide. By repeatedly treating the oxyfluoride with concentrated hydrofluoric acid, mercuric fluoride is obtained as a white crystalline mass of the formula HgF2,2OH2. The same compound is formed when mercuric oxide is added to a large excess of hydro- fluoric acid containing 50 per cent, lIF. When heated to 50° C. (122° F.) it is con- verted into the oxyfluoride. In contact with water it is decomposed into hydrofluoric acid and mercuric oxide. COMPOUNDS OF MERCURY WITH OXYGEN. Hgx Mercurous oxide, 'Hg'20. j yO. Hg/ Mercuric oxide, HgO. Hg=o. Mercurous oxide, 'Eg' 20.—This compound is obtained as a black powder by precipitating a mercurous salt with potassic or sodic hydrate. By the action of light it is decomposed into mercuric oxide and metallic mercury; for this reason it must be washed and dried in the dark. It COMPOUNDS OF MERCURY. 533 is decomposed in the same manner when heated to 100° C. (212° F.). Acids dissolve it, yielding the mercurous salts. Mercuric oxide, Hg'O, is formed as a red crystalline powder when mercury is heated in air to a temperature near its boiling-point. It is most conveniently prepared by thoroughly triturating mercuric nitrate with an equal weight of mercury and cautiously heating the mixture until acid fumes cease to be evolved. When prepared on a large scale by this method, it is sometimes obtained in small brick-red rhombic crystals, having a sp. gr. of 11.136. It is precipitated as a yellow amorphous powder when potassic or sodic hydrate is added to the solu- tion of a mercuric salt. Mercuric oxide is not quite insoluble in water, to which it imparts an alkaline reaction and a metallic taste. When carefully heated it assumes a deeper color, gradually passing into black, but recovers its original tint on cooling. At a red heat it is totally de- composed into mercury and oxygen. When heated with bodies which take up oxygen it oxidizes them : a mixture of mercuric oxide and sul- phur explodes with great violence on heating. Mercuric oxide is grad- ually blackened by exposure to light, owing to a partial decomposition. OXY-SALTS OF MERCURY. a. Mercurous Salts. Mercurous nitrate.—When mercury is dissolved in cold dilute nitric acid the solution deposits colorless monoclinic tables or prisms of the normal salt tetrahydrie mercurous dinitrate, jJqjj'^('Hg'gOg)"■ It dissolves without decomposition in water containing nitric acid, but in contact with an excess of pure cold water it is decomposed, yielding the basic salt hydric mercurous nitrate, NOHo(/Hg/202)//, as a yellow crys- talline powder which is converted on boiling with water into mercuric nitrate and metallic mercury. Other basic mercurous nitrates are known, some of which crystallize well. Thus when the crystals of the normal salt are heated with their mother liquor in contact with an excess of mercury, the solution deposits on cooling colorless, lustrous, non-efflorescent, rhombic prisms of hydric dimercurous trinitrate, N30sHo('Hg/202)//2. If, on the other hand, the crystals of the normal salt are left for some time in the cold in contact with the mother liquor along with an excess of mercury, lustrous triclinic prisms of tetrahydrie pentaraercitrous hexanitrate, N608Ho4('Hg/202)//5, are formed.—The normal mercurous nitrate forms numerous crystallizable double salts with the nitrates of other metals. f OCI Mercurous chlorate, , obtained in colorless rhombic prisms by dis- _ I OCI solving freshly precipitated mercurous oxide in chloric acid. When heated to 250° C. it decomposes into mercuric chloride, mercuric oxide, and oxygen. 534 INORGANIC CHEMISTRY. r oci o Mercurous perchlorate, , q (/Ilg Vb)/ /, G Oh is obtained in colorless deliquescent needles by dissolving mercurous oxide in an aqueous solution of perchloric acid. f OBr Mercurous bromate, J: Q(/Hg/202)//, is deposited in colorless laminae when solutions 1 OBr of mercurous nitrate and potassic bromate are mixed. An excess of water decom- poses it with formation of a basic salt. Mercurous carbonate, C0(/Hg/202)//, is precipitated as a yellow powder when a solu- tion of mercurous nitrate is poured into an excess of hydric potassic carbonate or hydric sodic carbonate. Mercurous carbonate decomposes at 130° C. (266° F.) into carbonic anhydride, mercury, and mercuric oxide. Mercurous sulphate, S02(/FFo;/202)//, is obtained as a white crystalline mass by gently heating sulphuric acid with an excess of mercury. If the temperature be raised too high, a mercuric salt is formed at the same time. Mercurous sulphate is also deposited in minute colorless prisms when dilute sulphuric acid is added to a solution of mercurous nitrate. It is only slightly soluble in water. When heated it fuses to a reddish-brown liquid which solidifies on cooling to a crystalline mass. With careful heating it may be sublimed. h. Mercuric Salts. Mercuric nitrate, j^^Hgo".—This salt is prepared by boiling mercury with an excess of nitric acid until a portion of the liquid, when removed and tested with a solution of common salt, yields no precipi- tate. The normal salt is, on account of its deliquescent nature, very difficult to obtain in a crystallized state. When the solution is evap- orated over sulphuric acid, large deliquescent crystals of dihydric dimer- curie tetranitrate, N4OrHo2Hgo"2, are obtained. A hydrated normal salt is deposited in tabular crystals of the formula N204Hgo/',BOH2, when a solution of the nitrate in nitric acid is cooled to —ls° C; the crystals fuse at ordinary temperatures. Mercuric nitrate has a great tendency to form basic salts: a solution of an excess of mercuric oxide in hot, moderately strong nitric acid, deposits on cooling colorless rhombic crystals of tetrahydric dimercuric dinitrate, N2OHo4Hgo//2. When this salt, or any of the normal salts, is treated with cold water, a still more basic salt, dihydric trimercuric dinitrate, N2OHo2Hgo//3, is formed as a white powder, and this, when boiled with an excess of water, gradually parts with all its acid and is converted into mercuric oxide. Mercuric carbonate.—The mercuric carbonates are basic compounds of ill-defined character and uncertain composition. They form brown amorphous powders. Mercuric sulphate, S02Hgo//.—This salt is prepared by heating mer- cury with one and a half times its weight of sulphuric acid until the excess of acid is expelled. It is thus obtained as a white crystalline COMPOUNDS OF MERCURY. 535 mass, which when heated turns first yellow and afterwards brown, but becomes white again on cooling. At a red heat it decomposes into mercury, oxygen, and sulphurous anhydride. When treated with a small quantity of water it forms white crystals of dihydric mercuric sulphate, SOHo2Hgo//, but an excess of water decomposes it, especially on boiling, into free sulphuric acid and a yellow insoluble basic salt, trimercuric sulphate, SHgo"3, formerly known as turpeth mineral (iurpe- tum miner ale). Mercuric orthophosphate, p[)|J^//llgo//.—This salt is precipitated as a heavy white insoluble powder when ordinary sodic phosphate is added to a solution of mercuric nitrate. Mercuric chloride cannot be substituted for the nitrate. Borates and silicates of mercury have not been prepared. COMPOUNDS OF MERCURY WITH SULPHUR. Mercuric sulphide, /Hg/2S//, is precipitated as a black powder by pouring a dilute solution of mercurous nitrate into a dilute solution of ammonic sulphhydrate. It may also be prepared by treating freshly precipitated calomel with ammonic sulphhydrate. It is a very unstable compound, and is decomposed even by a gentle heat into mercury and mercuric sulphide.* Mercuric sulphide, Cinnabar, Vermilion, HgS".—This compound occurs native in red hexagonal crystals, and also in granular masses, as the mineral cinnabar, constituting the most abundant ore of mercury. By triturating mercury with sulphur, the sulphide is obtained as a black amorphous powder; the product thus formed is known in pharmacy as Aethiops mineralis. The same black modification is obtained by pre- cipitating a mercuric salt with an excess of sulphuretted hydrogen. When the black amorphous sulphide is sublimed with exclusion of air, it is converted into the crystalline variety, which condenses on a cold surface, generally as a red fibrous mass, but sometimes in distinct crystals having the form of a native compound. By digesting with warm solu- tions of alkaline persulphides, the black sulphide is also converted into the red sulphide. The finely ground red sulphide is employed as a pigment under the name of vermilion, and is prepared on a large scale in the wet way by the following method: 100 parts of mercury are thoroughly triturated with 38 parts of flowers of sulphur, and the mass is then digested for several hours at a temperature of 45—50° C. (113- 122° F.) with a solution of 25 parts of caustic potash in 150 parts of water, renewing the water as fast as it evaporates. As soon as the ver- milion has attained the proper shade the operation is interrupted and the product is quickly washed with water, as by the further action of the potash the color changes to brown. Vermilion prepared in the wet way has a finer shade than that obtained by sublimation. Mer- curic sulphide is insoluble in hydrochloric, nitric, and sulphuric acids, but soluble in aqua-regia and in hydriodic acid. Ammonic sulphide does not dissolve it, but it is soluble in potassic and sodic sulphides in * According to some chemists the so-called mercurous sulphide is merely a mixture of mercuric sulphide and mercury. INORGANIC CHEMISTRY. presence of free alkali. The solution in potassic sulphide deposits colorless needles of mercuric dipotassic sulphide, HgKs2,50H2 The sodium compound has the formula HgNas2,80H2. Both compounds are decomposed by an excess of water with separation of black sulphide. By digesting the black sulphide with a solution of mercuric chloride, or by fusing the dry sulphide with an excess of mercuric chloride and extracting the mass with water, trimercuric disulphodichloride, ||gHgs" = Cl—Hg—S—Hg—S—Hg—Cl, is obtained as a white powder which is amorphous when prepared in the wet way, and crys- talline when prepared in the dry way. The same substance is formed as a white precipitate when sulphuretted hydrogen is passed into a solu- tion of mercuric chloride, but is converted by an excess of sulphuretted hydrogen into black sulphide. Mercuric sulphide forms numerous other double compounds with mercuric salts. Mercuric nitride, is formed when mercuric oxide, prepared by precipitation and dried at a low temperature, is heated to 100° C. in a current of ammonia: COMPOUND OF MERCURY WITH NITROGEN. 3HgO + 2NH3 = N2Hg"3 + 30 Ha. Mercuric Ammonia. Mercuric Water, oxide. nitride. The product is treated with dilute nitric acid to free it from any unaltered mercuric oxide. It forms a dark-brown powder, which explodes with great violence by heat, friction, or contact with concentrated sulphuric acid. By cautiously heating with caustic alkalies it is decomposed without detonation, yielding ammonia and sublimed mercury. AMMONIACAL MERCURY COMPOUNDS. These compounds are derived from the ordinary ammonium salts by the replace- ment of one or more atoms of hydrogen in the latter by mercury. The mercury may be present either as Hg" or as /lig/2, and each of these dyad radicals may either replace two atoms of hydrogen in a single ammonium molecule, or may replace two atoms of hydrogen in two different ammonium molecules; in the latter case uniting the two ammonium groups to a single molecule. The free mercury-ammoniums have not been prepared. Mereurosammonie chloride, NH2/Hg/2CI, is obtained as a black insoluble powder by the action of aqueous ammonia on mercurous chloride: a. Mercurosammonium Compounds. 'Hg'jCls + 2NIIS = NHj'Hg'gC 1 + NH4CI. Mercurous Ammonia. Mercuros- Ammonie chloride. ammonie chloride. chloride. Gaseous hydrochloric acid decomposes it, yielding mercurous chloride and ammonie chloride: NH2/Hg/2CI + 2HCI = 'Hg'jd, + NH4CI. Mercuros- Hydrochloric Mercurous Ammonie ammonie chloride. acid. chloride. chloride. When heated, it first gives off ammonia and nitrogen, and afterwards mercurous chlo- ride and metallic mercury. Mereurosammonie nitrate, N02(NvH2/Hg/20).—This compound, known as Mereurius solubilis Hahnemanni, is precipitated in the form of a black powder when aqueous am- COMPOUNDS OF MERCURY. 537 monia is added to a solution of mercurous nitrate. It is with difficulty obtained in a state of purity, and is generally mixed with metallic mercury. Mercuresodiammonic dichloride, jS( jj } > f°rme(i as a black powder by the action of gaseous ammonia upon mercurous chloride. When heated, or when exposed to the air, it evolves ammonia, leaving mercurous chloride. b. Mereurammonivm Compounds. NHoy/ Dimercurammonic oxide, —When mercuric oxide is treated with concen- trated ammonia, dimercurammonic hydrate is obtained as a pale yellow powder having the formula N Hg//2H0,20H2. By sudden heating or by friction it deflagrates without explosion. When heated to 80° C. (176° F.) in a current of ammonia it gives off water, leaving the pure hydrate NHg//2Ho, which at 100° C. (212° F.) parts with the elements of water, and is converted into the oxide- Dimercurammonic oxide is a brown powder, which on heating or by friction explodes violently. Mercurammonic chloride, NH2Hg//Cl.—This compound, known as infusible white pre- cipitate, is prepared by precipitating a solution of mercuric chloride with an excess of ammonia: HgCI2 + 2NII3 = NH2Hg//CI + NH4CI. Mercurous Ammonia. Mercurammonic Ammonic chloride. chloride. chloride. At a temperature below a red heat it is decomposed without fusion, yielding ammonia, nitrogen, and mercurous chloride. Water decomposes it, slowly in the cold and quickly on heating, with formation of ammonic chloride and an aquate of dimercurammonic chloride, NHg//2C1,0H2. Dimercurammonic chloride, Nidged, is obtained as a yellow powder by the action of alcoholic hydrochloric acid on dimercurammonic oxide (see above), or by treating mercurammonic chloride with water (see preceding compound). When heated to 300° C. (572° F.) it decomposes into metallic mercury, mercurous chloride, and nitrogen. Mercuridiammonic dichloride, —This compound, known as fusible white precipitate, is obtained by adding a solution of mercuric chloride to a boiling solution of ammonic chloride and ammonia as long as the precipitate which is at first formed continues to dissolve. The liquid on cooling deposits colorless regular dodecahedra, which fuse when heated, and then decompose, yielding nitrogen, ammonia, mercurous and mercuric chlorides, and ammonic chloride. # Hydroxydimercurammonic iodide, NHHg//(Hg//Ho)I, is formed by the action of an excess of aqueous ammonia upon mercuric iodide: 4NFIS + 2HgT2 + OH2 = NHHg"(Hg"Ho)I + 3NHJ. Ammonia. Mercuric Water. Hydroxydimereur- Ammonic iodide. ammonic iodide. iodide. It is most readily obtained by adding ammonia to a solution of mercuric potassic iodide containing an excess of potassic hydrate. This liquid, which is known as Nessler’s solution, is employed in testing for minute traces of ammonia. When the quantity of ammonia is too small to yield with this reagent a precipitate of hydroxy- dimercurammonic iodide, it manifests its presence by a yellow coloration. Hydroxy- dimercurammonic iodide is a reddish-brown powder, which fuses when heated, and at a higher temperature decomposes with a violent explosion. Characteristic Properties and Reactions of the Com- pounds of Mercury.—The normal salts of mercury with colorless acids are colorless; some of the basic salts are yellow. The soluble salts have an acid metallic taste, and act as irritant poisons. If a strip of copper be introduced into a solution of any mercury compound, metallic mercury is deposited on the copper. All compounds of mer- cury, when heated in a test-tube with dry sodic carbonate, yield a gray sublimate consisting of minute globules of mercury. 538 INORGANIC CHEMISTRY. a. Mercurous salts, when in solution, yield with caustic alkalies black mercurous oxide. Ammonia precipitates black mercurosarnmonium compounds (p. 536). Sulphuretted hydrogen and ammonic sulphide pre- cipitate black mercurous sulphide, insoluble in nitric acid, soluble in aqua regia. Hydrochloric acid precipitates white mercurous chloride, and potassic iodide green mercurous iodide. Stannous chloride precipi- tates mercurous chloride, which is converted by an excess of the stan- nous chloride into gray metallic mercury. , b. Mercuric salts give, with solutions of caustic alkalies, a yellow precipitate of mercuric oxide. A7nmonia precipitates a white mercur- ammonium compound (p. 537). Sulphuretted hydrogen gives a white precipitate, which passes through red to black, and then consists of mercuric sulphide; this precipitate is insoluble in nitric and in hydro- chloric acid, soluble in aqua-regia. Potassic iodide precipitates salmon- red mercuric iodide, soluble both in mercuric chloride and in potassic iodide. Stannous chloride precipitates mercurous chloride, which is then converted into metallic mercury. The mercury compounds give no flame-coloration. The spark spec- trum displays bright lines in the green and blue. COPPER, Cu. Atomic weight 63.2. Probable molecular weight = 63.2. Sp.gr. 8.9. Fuses at 1330° C. (2426° F.). Atomicity ", also a pseudo-monad. Evidence of atomicity: Cupric chloride, . . . Cu"Cl2. Cupric oxide, .... . . . Cu"0. Cuprous chloride, . . . . W2C12. Cuprous oxide, ... . . . . . 'Cu',0. History.—Copper has been known from prehistoric times. Owing to its occurring in the native state, it formed the material for tools and weapons in early ages when the metallurgical processes necessary for the extraction of iron from its ores were unknown. Occurrence.—Native copper occurs in various localities, particularly in the neighborhood of Lake Superior, where it is sometimes found in enormous masses. In the Minnesota Mine, in 1857, a mass of copper weighing 420 tons was found. In combination it occurs as cuprous oxide, 'ou'20, in red copper ore or ruby-ore ; much more rarely as cu- pric oxide, CuO, in tenorite. It also occurs as cuprous sulphide, in copper glance; as cupric sulphide, CuS", in indigo copper; as a double sulphide of copper and iron, diferric cuprous tetrasulphide, Cl/ Cupric sulphide, . . . . . CuS". Cu=S Cuprous sulphide, 'Cu'2S.—This compound occurs native as cop- per glance, and forms lead-gray rhombic tables or prisms with a me- tallic lustre, and having a sp. gr. of 5.5 to 5.8. The same compound is obtained as a black, brittle mass by heating together 4 parts of cop- per tilings and 1 part of sulphur, or by burning copper in sulphur vapor. Cupric sulphide, CuS, also occurs native as the mineral indigo- copper, but much less abundantly than the cuprous compound. It sometimes forms dark-blue hexagonal crystals with a semi-metallic lustre, but more frequently occurs massive. Its sp. gr. is 4.6. It may be obtained as a blue powder by heating finely divided copper with flowers of sulphur, avoiding a temperature higher than the boiling point of sulphur. It is obtained as a black amorphous precipitate when sulphuretted hydrogen is passed into solutions of cupric salts, and in this condition is readily oxidized if exposed to the air while still moist. The precipitated sulphide is insoluble in potassic and sodic sulphides, somewhat soluble in yellow aramonic sulphide,’ readily soluble in po- tassic cyanide and in hot nitric acid. When cupric sulphide is heated with exclusion of air, or in a current of hydrogen, it parts with half its sulphur and is converted into cuprous sulphide,—When an ammoni- acal solution of a copper salt is precipitated with sulphuretted hydro- gen a black precipitate of cupric sulphide is obtained. If this precipi- tate be washed for a very long time with sulphuretted hydrogen water, until the last traces of ammonia compounds are removed, the black sul- phide at last goes into solution, yielding a dark-brown liquid which is believed by some chemists to contain a colloidal modification of the sulphide. Solutions of salts precipitate from the liquid insoluble cupric sulphide. On evaporation the black liquid dries up to a black lustrous film. Similar colloidal modifications of sulphides have been obtained in the case of various other heavy metals.* * It is, however, probable that these so-called colloidal sulphides are nothing more than ordinary sulphides in a state of very fine subdivision. Ebell, Avho has advanced this view, has shown that the finest ultramarine, obtained by grinding and levigation, can be removed by filtration from liquids containing a salt in solution; but if the ultra- marine upon the filter be washed with pure water, it passes through the filter as soon as the salt solution has been sufficiently removed, and yields a blue liquid which to the eye is perfectly transparent, but which under the microscope is seen to contain minute sus- pended particles of ultramarine. In pure water these minute particles show no ten- dency to subside; but the addition of a small quantity of the solution of a salt pre- cipitates the ultramarine. If the salt solution be added to a drop of the blue liquid 550 INORGANIC CHEMISTRY. COMPOUNDS OF COPPER WITH NITROGEN, PHOSPHORUS, AND ARSENIC. Cuprous nitride, N2(/Cu/2)//3, is obtained as a dark green powder when gaseous am- monia is passed over finely-divided cupric oxide heated to 250° C.: 6CuO + 4NH3 = Nsf'Cu's)"# + Na + 60H2. Cupric Ammonia. Cuprous fWater. oxide. nitride. At 300° C., it is decomposed, with a slight explosion, into its elements. Cuprous phosphide, P2(/Cu/2)//3, is formed when cuprous chloride is heated in a cur- rent of phosphoretted hydrogen, or when the vapor of phosphorus is passed over copper foil heated to low redness. By fusing the compound under a layer of borax it may be obtained in the form of a silver-white regulus of sp. gr. 6.59, very brittle, and capable of taking a polish. Hydrochloric acid is almost without action upon it, but nitric acid dissolves it readily. Cupric phosphide, is prepared in a similar manner by passing phosphoretted hydrogen over heated cupric chloride. It forms a black lustrous powder, which when heated in a current of hydrogen is converted into cuprous phosphide. It is also formed as a black precipitate when phosphoretted hydrogen is passed into the solution of a cupric salt (p. 342). Cuprous arsenide, As2(/Cu/2)//3) occurs in Chili as the mineral domeykite, forming tin-white or silver-white masses. Other arsenides of copper also occur as minerals. General Properties and Reactions of the Compounds of Copper.—The soluble compounds of copper have a disagreeable me- tallic taste, and are poisonous, causing vomiting and death. a. Cuprous Compounds.—The cuprous salts are colorless. They are generally insoluble in water, but soluble in hydrochloric acid and am- monia. In solution they rapidly absorb oxygen from the air, and are converted into cupric salts. Caustic alkalies precipitate yellow cuprous hydrate, which is converted on boiling into red cuprous oxide. b. Cupric Compounds.—The cupric salts are white in the anhydrous state, blue or green when hydrated. They are nearly all soluble. The Solutions redden blue litmus. Caustic alkalies precipitate blue cupric hydrate, which on boiling is partially converted into cupric oxide and becomes black. The presence of sugar, tartaric acid, and various other organic substarices, renders the cupric hydrate soluble in an excess of alkali. Ammonia gives a similar precipitate, which is, however, soluble in excess, yielding a deep-blue liquid. Sulphuretted hydrogen precipi- tates from acid solutions brownish-black cupric sulphide, slightly soluble in yellow ammonic sulphide, readily soluble in potassic cyanide, and in hot nitric acid. Potassic ferrocyanide gives & brown precipitate, in- soluble in hydrochloric acid. From solutions of copper compounds zinc and iron precipitate me- tallic copper. All compounds of copper, when heated with sodic car- bonate on charcoal in the reducing flame of the blowpipe, yield a bead of metallic copper. A borax bead containing a copper salt, and heated under the microscope, the separate particles of ultramarine are seen to unite into aggre- gations, each consisting of a number of particles. On evaporation, the blue liquid yields a lustrous blue film adhering to the sides of the vessel. The behavior of this finely-divided ultramarine—a substance which cannot in any sense be regarded as colloidal—corresponds therefore, in all the above particulars, with that of the metallic sulphides referred to. GOLD, 551 in the oxidizing flame, is green while hot and blue when cold ; in the reducing flame the bead is colorless if the proportion of copper be small, but, if the proportion of copper be large, the bead is red from the presence of reduced copper. The compounds of copper color the non- luminous flame green or blue. Cupric chloride gives a banded flame- spectrum, this being the spectrum of the compound. The spark-spec- trum of copper contains a number of lines, among which some of those in the green are especially prominent. CHAPTER XXXY. TRIAD ELEMENTS. Section 11. GOLD, Au2? Atomic weight = 196. Probable molecular weight = 392. Sp. gr. 19.3 to 19.5. Fuses at 1240° C. (2264° F.). Atomicity ' and Evi- dence of atomicity: Aurous chloride, . . , . . . AuCl. Aurous iodide, . . . , . . . Aul. Auric chloride, . . . , . . . Au"'C13 Auric hydrate, . . . . Au'"Ho3, History.—Gold has been known and prized from the earliest his- torical times. Occurrence.—Gold occurs widely distributed, but mostly only in small quantity. It is almost always found in the native state, some- times in crystals, sometimes in dendritic forms produced by the regular aggregation of crystals, but most frequently in irregular masses termed nuggets. In matrix it is found disseminated throughout quartz veins or reefs. The alluvial deposits produced by the disintegration of the auriferous rocks form the chief sources of the metal. The principal gold-fields are those of California and Australia. Gold is still extract I from the sand of rivers in Hungary aud Transylvania, but the im- portance of these sources has diminished since the discovery of the Australian and Californian fields. Native gold generally contains more or less silver; if the percentage of silver exceeds 36 per cent, this native alloy is termed electrum. Gold is found in combination with bismuth and tellurium in a few rare minerals, and alloyed with mer- cury as an amalgam. Traces of the metal occur in many ores of silver, copper, aud lead, and in iron pyrites. In spite of the smallness of the quantity present, it is possible in some of these cases to extract the gold with profit (see p. 450). 552 INORGANIC CHEMISTRY. Extraction.—Native gold is mechanically separated from the alluvial deposits with which it is mixed by washing away the lighter earthy particles—either by the simple manual processes of pan-washing or cradle-washing, or, on a large scale, by hydraulic gold-mining. In the latter process enormous jets of water are employed to remove the whole of the alluvial deposit down to the bed-rock. The stream of water, carrying with it the disintegrated deposit, flows through a long sloping tunnel bored in the rock. Along .the bottom of the tunnel are placed “ sluice-boxes ” containing a small quantity of mercury. The particles of gold fall into the sluice-boxes and are arrested by the mercury with which they form an amalgam. The tunnel is cleared at intervals of from ten to twenty days : the amalgam of gold is removed, and the mercury expelled by distillation. In quartz-mining the auriferous quartz is stamped to a fine powder by special machinery, and the gold extracted by amalgamation. Refining.—One of the simplest and most efficient refining processes is that devised by F. B. Miller. The gold, which must not contain more than 10 per cent of silver, is melted in a clay crucible glazed in- side with borax, and a current of chlorine is passed through the molten metal. The silver is thus converted into argentic chloride, which rises to the surface and is prevented from volatilizing by a layer of fused borax; other foreign metals, such as zinc, antimony, bismuth, and tin, are volatilized as chlorides. The metal thus purified contains from 99.1 to 99.7 per cent, of gold. Pure gold may be prepared by dissolving the metal in aqua-regia, and, after expelling the excess of nitric acid, precipitating the gold by some reducing agent, such as ferrous sulphate. The finely divided gold is obtained in a coherent form by fusion with a mixture of borax and nitre. Properties.—Gold is a lustrous metal, .of a yellow color when the light is only once reflected, but red when the light is several times re- flected from the surface of the metal before reaching the eye (p. 400). It is the most malleable and ductile of the metals (pp. 409 and 410). Very thin gold leaf transmits green light. When pure it is nearly as soft as lead. It fuses at 1240° C. (2264° F.), the molten metal emit- ting a bluish-green light. At very high temperatures it is volatile. It is quite unalterable in air, oxygen, and steam, at all temperatures. No single acid, with the exception of selenic, has any action upon it; but aqua-regia, and all other liquids containing or evolving chlorine, dis- solve it with formation of auric chloride (AuCI3). It combines with chlorine and bromine at ordinary temperatures, and with phosphorus when heated in its vapor. It is precipitated from its solutions by most other metals, and by most reducing agents. Ferrous sulphate precipi- tates it as a brown powder without metallic lustre; oxalic acid, in glistening yellow scales. Uses.—Gold is employed for coinage, for ornaments, and in gilding. Non-metallic surfaces are gilt with gold-leaf. Metals are gilt by electro- deposition, employing a solution of auric chloride in potassic cyanide (the solution contains auric potassic cyanide, AuCy3,KCy) and using a gold plate as positive electrode. COMPOUNDS OF GOLD. 553 Alloys.—Pure gold is employed in the preparation of gold-leaf and of the solutions for electro-gilding, but owing to its softness is not suited for the manufacture of objects which have to resist the wear of ordinary use. For jewellery or coinage gold is therefore alloyed with copper, or with silver, or with both, these admixtures imparting to the gold the requisite hardness. The copper alloy has a reddish tinge, that with silver is whiter than pure gold. The proportion of gold in an alloy is frequently expressed in carats, or parts per 24: thus 24-carat gold is pure gold, 22-carat gold contains 22 parts of gold in 24 parts of the alloy, and so on. In most countries the composition of various standard alloys for jewellery and coinage is fixed by law. In England there are five legal standards: 22- carat—the standard gold employed for coinage, the two remaining parts in this case consisting of copper—18, 15, 12 and 9-carat gold. In the case of coinage standards, however, ihisraore usual to express the proportion of gold in parts per mille of the alloy, this expression being known as the fineness of the alloy. The English 22-carat standard gold has thus a fineness of 916.66. Most other European countries employ a coinage standard having a fineness of 900. Gold forms two classes of compounds, aurous and auric. In the first of these it is a monad, in the second a triad. COMPOUNDS OF GOLD WITH THE HALOGENS. a. Aurous Compounds. Aurous chloride, AuCl, is obtained by heating auric chloride, AuCls, to 185° C. (365° E.). It is a yellowish-white powder, which is decomposed at a higher tempera- ture into gold and chlorine. Water decomposes it into metallic gold and the trichloride. Aurous iodide, Aul, is formed by the action of hydriodic acid upon auric oxide: Au2o3 + 6HI = 2AuI + 30H2 + 212, Auric Hydriodic Aurous Water, oxide. acid. iodide. and in all similar cases when the formation of an auric iodide might be expected, the latter compound undergoing decomposition into Aul-j-I2—thus by the action of po- tassic iodide upon auric chloride : AuC13 + SKI = Aul + I 2 + 3KCI. Auric Potassic Aurous Potassic chloride. iodide. iodide. • chloride. —Aurous iodide forms a lemon-yellow powder, which is decomposed, slowly at ordi- nary temperatures, rapidly on heating, into its elements. b. Auric Compounds. Auric chloride, AuC!3.—This compound is obtained by the action of chlorine upon gold, or by dissolving gold in aqua-regia, evaporating to dryness, taking up with water, evaporating again to dryness, and heating carefully to 150° C. (302° F,). The anhydrous chloride forms a brown crystalline deliquescent mass. Though decomposed at 185° C. (365° F.), as already mentioned, into aurous chloride and chlorine, it may be sublimed in a current of chlorine at 300° C. (572° F.), and is thus obtained in long red needles. When a hot concentrated aqueous 554 INORGANIC CHEMISTRY. solution of auric chloride is allowed to cool, an aquate of the formula AuC13,20H2, is deposited in large orange-colored crystals. Auric chloride forms numerous compounds with other metallic chlorides and with hydrochloric acid. The hydrochloric acid compound, sometimes called hydrauric acid, has the formula AuCI3,HGI,3OH2, and crystallizes from the concentrated solution of gold in aqua-regia in long yellow needles. Auric potassic chloride forms two aquates— (AuC 13,KC1)2,0 Ha, crystallizing in needles, and AuCI3,HCI,2OH2, crystallizing in large rhombic tables. Auric sodic chloride, AuCI3,NaGI,2OH2, crystallizes in yellowish-red prisms. Auric amnionic chloride for ms light yellow rhombic tables, (AuC!3,NH4C1)2-r)OIl2, or monoclinic plates These double chlorides are sometimes referred to as chloraurates, thus potassic chloraurate. Auric bromide, Außr3, forms a black crystalline mass. Auric iodide, Aul;), is not known as such, but several double compounds of this iodide with iodides of other metals have been prepared. COMPOUNDS OF GOLD WITH OXYGEN AND HYDROXYL. Aurous oxide, 0Au2. Au—O—An. TAuO Auric oxide (Auric, anhydride) < O . o=Au—O—Au=C). (AuO yO H Auric hydrate, AuHo3. Au^-O—H. X)—H Aurous oxide, OAu2, is obtained as a violet-black powder by the action of dilute caustic potash upon aurous chloride. At 150° C. (302° F.) it is decomposed into its elements. With hydrochloric acid it yields auric chloride and metallic gold : 30Au2 + 6HCI 2AuCi3 + 2Au3 + 30H2. Aurous Hydrochloric Auric Water, oxide. acid. chloride. Sulphuric and nitric acids are without action upon it, but aqua-regia dissolves it readily. Auric oxide (Auric anhydride), Au203.—This compound is pre- pared by heating a solution of auric chloride with magnesia and treat- ing the precipitate, which consists of magnesia aurate, AuO Vl&° ’ with concentrated nitric acid, in which the whole dissolves. Water precipitates auric hydrate, AuHo3, as a reddish-yellow powder, which by gentle heating is converted into the oxide. It forms a brown powder which is partially decomposed at 100° C. (212° F.), wholly at 245° C. (473° F.), into its elements. It is the anhydride of auric acid, AuOHo, and dissolves in dilute caustic potash to form potassic aurate, which crystallizes in light yellow needles of the formula AuOKo,3OH2. COMPOUNDS OF GOLD. 555 —A derivative of auric anhydride is fulminating gold, a compound which is formed by the union of four molecules of ammonia with one of auric anhydride, and which may be regarded as possessing f Au(NH2)(NH/)) the constitution < O , It is best prepared by treating . Uu(NH2)(NH4O)> auric hydrate with aqueous ammonia. It forms a yellowish-brown or greenish-yellow powder, which when dry explodes with great vio- lence by heat or percussion. A similar compound, which however appears to contain chlorine, separates when ammonia is added to a solu- tion of auric chloride. Auric hydrate, AuHo3, may be obtained either as above described, or by electrolyzing dilute sulphuric acid, employing a gold plate as posi- tive electrode, when the hydrate is formed as a yellow crust on the electrode. OYY-SALTS OF GOLD. Simple oxy-salts of gold are not known. Double salts have however been prepared, such as the double thiosulphate of gold and sodium, SO2AuoAus,3SO2NaoNas,4oH2, which might also be formulated— " SHo2NaoAus - O SHo2NaoNas , ■, • 7 ,7. q q octohydnc diaurous hexasodic ~ SHo2NaoNas|’ tetrathioeulphate. 0 v SHo2NaoAus -I It is formed when a dilute neutral solution of auric chloride is added to an excess of a solution of sodic thiosulphate. A reduction of the gold from the auric to the aurous condition occurs, the red liquid which is at first formed becoming colorless. The salt is then precipitated by the addition of strong alcohol. It crystallizes in colorless needles which have a sweet taste. Neither the gold nor the thiosulphuric acid can be detected by the usual tests: the gold is not precipitated by re- ducing agents, and no separation of sulphur occurs on the addition of dilute acids. Double sulphites of gold with the alkali metals are also known. Aurous amnionic sulphite has the formula SOAmo2,350Au0Am0,30H2. Purple of Cassius.—This remarkable compound is obtained as a flocculent purple precipitate when a very dilute mixed solution of stannous, and stannic chloride is gradually added to a dilute neutral solution of auric chloride. It contains one or both of the oxides of tin. Its nature is not known with certainty, but it is supposed to be a hydrated stannous diaurous distannate, Sn202Sno" Auo2,4 0 H2. 556 INORGANIC CHEMISTRY. Its composition, however, is apt to vary with the mode of preparation. The compound is decomposed by acids with separation of metallic gold. It is insoluble in solutions of caustic potash and caustic soda, but soluble in ammonia,-yielding a deep purple liquid which is bleached by exposure to light with deposition of metallic gold and formation of ammonic stannate. Purple of Cassius is employed to impart a magnificent red color to glass. The color depends upon the presence in the glass of metallic gold in a state of minute subdivision. COMPOUND OF GOLD WITH SULPHUR. Diaurous disulphide, /S/2A.u2, is precipitated by sulphuretted hydro- gen from cold solutions of auric chloride: 2AuC13 + 3SH2 = 'S'2Au2 + 6HCI + S. Auric Sulphuretted Diaurous Hydrochloric chloride. ' hydrogen. disulphide. acid. It forms a black precipitate, insoluble in water, soluble in solutions of the alkaline sulphides, with formation of double sulphides such as SNaAu: Diaurous Disodic Sodic aurous Disodic disulphide. sulphide. sulphide. disulphide. 'S'2Au2 + 2SNa2 = 2SNaAn + 'S'2Na2. From hot solutions of gold salts sulphuretted hydrogen precipitates metallic gold. General Properties and Reactions of the Compounds of Gold.—Gold is precipitated from its solutions by most reducing agents —e.g., ferrous sulphate, mercurous nitrate, oxalic acid, formic acid, sul- phurous acid—as finely divided metallic gold. A mixture of stannous and stannic chlorides produces a characteristic precipitate of purple of Cassius (p. 555). All gold compounds are converted into metallic gold when ignited with exposure to air. The compounds of gold do not color the uon-luminous flame. Section 111. THALLIUM, Tl2? Atomic weight = 204. Probable molecular iceight = 408. Bp. gr 11.8 to 11.9. Fuses at 294° C. (561.2° F.). Atomicity' and Evi- dence of atomicity: Thallous chloride, TICI. Thallous oxide, 0T12. Thallic chloride, Tl'"^. History.—Thallium was discovered by Crookes in 1861, while exam- ining spectroscopically a seleniferous deposit from a sulphuric acid 557 COMPOUNDS OF THALLIUM. manufactory in the Harz. It was at first supposed to be a non- metal, allied to sulphur. In 1862 it was discovered independently by Lamy, who first recognized its metallic character and succeeded in iso- lating it. The name thallium, derived from OaUd<;i a green twig, was given to this element in allusion to the bright green line which constitutes its visible spectrum and by means of which it was discovered. Occurrence.—Thallium occurs widely distributed in nature, but only in small quantities. Certain varieties ef pyrites—notably Belgian, Westphalian, and Spanish pyrites—contain traces of thallium, and when such pyrites is burnt in the manufacture of sulphuric acid, the thallium condenses and collects in the form of thallous oxide, along with arsenious anhydride and other substances, as a fine dust in the flues of the pyrites burners. Salts of thallium occur in minute quantity in some mineral springs. As an essential constituent, it is found only in the rare mineral crookesite, a selenite of copper, silver, and thallium, containing from 16 to 18 per cent, of the latter metal. Preparation.—When the flue dust containing thallium is treated with dilute sulphuric acid, the thallium goes into solution as thallous sulphate, S02TIo2, and may be precipitated as sparingly soluble thallous chloride by the addition of hydrochloric acid to the filtered solution. The washed chloride is separated and reconverted into sulphate by treat- ment with sulphuric acid, heating to expel the hydrochloric acid. The sulphate is purified by crystallization, and from the solution of the pure sulphate metallic thallium is obtained by electrolysis or by precipitation with zinc. The metal, which is thus deposited in soft laminar crystals or as a spongy mass, may be obtained in a coherent form by fusion in a covered crucible under potassic cyanide. Properties.—Thallium is a heavy metal, white like tin, and soft enough to be scratched with the finger-nail. It may be distilled at a white heat in a current of hydrogen. When exposed to the air it tarnishes superficially, and is converted into thallous oxide. It does not decompose water below a red heat, and is best preserved in closed vessels under water. With access of air it slowly dissolves in water, with formation of thallous hydrate, which in solution absorbs carbonic anhydride, and is ultimately converted into carbonate. Dilute acids readily dissolve it. It is precipitated in the metallic state from its solu- tions by zinc, but it precipitates lead, copper, mercury, and silver from the solutions of their salts. Thallium forms two classes of compounds—thallous compounds, in which the metal is monadic, and thallic compounds, in which it is tri- adic. The members of the first class are the most numerous and best characterized. COMPOUNDS OF THALLIUM WITH THE HALOGENS. a. Thallous Compounds. Thallous chloride, TICI, Molecular volume I I I.—This compound is obtained as a curdy precipitate when hydrochloric acid is added to a 558 INORGANIC CHEMISTRY. not too dilute solution of thallous hydrate or a thallous salt. It is colored violet by exposure to light. It is soluble in 360 parts of water at ordinary temperature, in from 50 to 60 parts at 100° C. (212° F.). From the hot saturated aqueous solution it crystallizes in cubes. It is less soluble in water containing hydrochloric acid than in pure water. It is readily fusible, yielding a yellow liquid, which solidifies to a white crystalline mass. At higher temperatures it volatilizes. Thallous bromide, TIBr, forms a yellow precipitate. It is less soluble in water than the chloride, which it closely resembles. Thallous iodide, Til, is precipitated as a yellow crystalline powder when potassic iodide is added to the solution of a thallous salt. It is almost insoluble in water. Ex- posure to sunlight colors it green. It is readily fusible, and solidifies to a red crystal- line mass, which becomes yellow on standing. At a higher temperature it may be sub- limed with partial decomposition. Thallous fluoride, TIE, is prepared by dissolving thallous carbonate in hydrofluoric acid and evaporating. It crystallizes in colorless, very lustrous anhydrous octahedra, or, with water of crystallization, in hexagonal plates. It dissolves readily in water, and is fusible and volatile. When exposed to sunlight it becomes dark-colored. Solutions containing an excess of hydrofluoric acid deposit on evaporation over sulphuric acid in vacuo regular crystals of an acid fluoride, TIE,HE. Thallic chloride, TIC13, is formed when thallous chloride is suspended in water and chlorine passed into the liquid. On evaporation in vacuo colorless deliquescent prisms of the formula TIC13,0H2 are deposited. h. Thallic Compounds. Chlorides of thallium intermediate between thallous and thallic chloride are known. In these the thallium is in the triadic condition: TetrathaUic hexachloride, | *s when metallic thallium is strongly heated in a current of chlorine. A yellowish-brown mass is thus obtained, sparingly soluble in cold, but readily soluble in boiling water, and crystallizing from the solution in yellow laminae. f TICI Dithallic tetrachloride, ■< —When metallic thallium or thallous chloride is cautiously heated in chlorine, a compound of the above composition is obtained, which, on heating more strongly, parts with chlorine, and is converted into tetrathallic hexachloride. Thallic bromide, TlBr3, and thallic iodide, Tils, are also known. They resemble the chloride, but are less stable. COMPOUNDS OF THALLIUM WITH OXYGEN AND HYDROXYL. Thallous oxide, . . . oti2. H 1 OO 1 II HH f TIO | Tliallic oxide, . . . Ao . O (TIO 1 T1=0 Thallous hydrate, . TIHo. Tl—0—H. Thallic oxyhydrate, , . TIOHo. 0=T1—0—H. Thallous oxide, 0T12.—Metallic thallium when exposed to the air tarnishes, owing to the formation of a coating of thallous oxide. COMPOUNDS OF THALLIUM. 559 The oxide may be obtained pure by heating the hydrate to 100° C. (212° F.) with exclusion of air. It forms a black powder which fuses at 300° C. (572° F.) to a dark-yellow liquid. It attracts moisture from the air, and dissolves in water with formation of thallous hydrate. Thallic oxide, T12Os.—This oxide is formed when thallium burns in oxygen, and may also be obtained by heating thallic oxyhydrate to 100° C. (212° F.). It is a dark-red powder, insoluble in water. At a red heat it evolves oxygen, and is converted into thallous oxide. Hot concentrated sulphuric acid dissolves it with evolution of oxygen and formation of thallous sulphate. When acidulated water is electrolyzed, employing a positive electrode of thallium, the metal becomes covered with a black deposit of thallic oxide. Thallous hydrate, TlHo, is formed when thallium is simulta- neously acted upon by water and air or oxygen. It is most readily obtained pure by precipitating thallous sulphate with baric hydrate and evaporating the filtrate. It crystallizes in colorless or faint-yellow rhombic prisms, having the composition TIHo,OH2. It is readily soluble in water and in alcohol, yielding powerfully alkaline solutions. The brown stain which it produces upon turmeric paper disappears, however, after a time, owing to a peculiar destructive action which the hydrate exercises upon the coloring matter. Thallous hydrate is con- verted at 100° C., or in vacuo at ordinary temperatures, into thallous oxide. Thallic oxyhydrate, TlOHo.—This compound is produced as a brown precipitate when freshly precipitated thallous chloride is warmed with a solution of sodic hypochlorite. It is also formed by the action of a caustic alkali upon thallic chloride. It is a brown powder, which at 100° C. (212° F.) is converted into thallic oxide. OXY-SALTS OF THALLIUM. a. Thallous Salts. Thallous nitrate, N02Tlo.—This salt is obtained by dissolving the metal in nitric acid. The solution deposits opaque white rhombic prisms, soluble in about 10 parts of water at the ordinary temperature, very readily soluble in boiling water. It fuses without decomposition about 205° C; but is decomposed at a higher temperature. Thallous carbonate, COTIo2, is formed when a solution of thallous hydrate, or metallic thallium moistened with water, is exposed to the air. It is best prepared by saturating a solution of the hydrate with carbonic anhydride and evaporating to the crystallizing point. It is deposited from the aqueous solution in long, lustrous monoclinic prisms. It dissolves in 20 parts of cold water, yielding a solution with an alka- line reaction. It is fusible without decomposition, but at higher tem- peratures evolves carbonic anhydride. Thallous sulphate, S02TIo2, crystallizes in rhombic prisms, and is isomorphous with potassic sulphate. It is soluble in 20 parts of water at ordinary temperatures and in 5 parts at 100° C. (212° F.). 560 INORGANIC CHEMISTRY. When air is excluded, it fuses at a red heat without decomposition; but when heated in air, it is decomposed with evolution of sul- phurous anhydride.—llydric thallous sulphate is deposited from so- lutions containing a large excess of sulphuric acid. It crystallizes in short thick prisms, having the formula SO2HoTlo,3oH.—With the sulphates of the dyad metals thallous sulphate forms double salts, ( S02TIo such as the double sulphate of zinc and thallium, < Zno" ,60H2, t S02TIo corresponding with the double sulphates of ammonium and potassium with the dyad metals, and like these, containing 6 molecules of water of crystallization. Thallous phosphate, POTIos, is obtained as a white crystalline precipitate when a thallous salt is added to a solution of ordinary sodic phosphate containing ammonia. It dissolves in 200 parts of cold and in 150 parts of boiling water. It is soluble in so- lutions of ammonia salts.—Ilydric dithallous phosphate, POHoTIo2,OH2, is prepared by neutralizing a solution of phosphoric acid with thallous carbonate. The solution deposits on evaporation rhombic crystals, which part with their water of crystallization at 200° C, and at a red heat are converted into a vitreous mass of thallous pyrophosphate, P‘203TIo4.—Dihydrie thallous phoshate, POHo2TIo, is prepared by adding to a solution of thallous carbonate sufficient phosphoric acid to produce a distinctly acid reaction, and then evaporating. It forms nacreous monoclinic prisms or laminae, readily soluble in water. At a red heat it is converted into metaphosphate. fNO,—| Thallic nitrate, -1 N02—Tlo///,80H2, is deposited in colorless crystals from the solu- . I N02—1 tion of the thallic oxide in concentrated nitric acid. Excess of water decomposes the salt with separation of thallic oxyhydrate. Thallic sulphate, S306Tk///2.70H2, crystallizes in thin colorless laminae from a solu- tion of thallic oxide or hydrate in warm dilute sulphuric acid. Water decomposes it in the cold. When heated it gives off sulphuric acid, sulphuric anhydride, and oxy- gen, and is converted into thallous sulphate. b. Thallic salts. COMPOUNDS OF THALLIUM WITH SULPHUR. Thallous sulphide, ST12.—This compound is obtained as a brownish- black amorphous precipitate when sulphuretted hydrogen is passed into an alkaline or acetic acid solution of a thallium salt. From a solution of thallous sulphate containing a trace of free sulphuric acid, it is deposited in minute, lustrous, dark-blue tetrahedra. It may be ob- tained as a black, lustrous, crystalline mass by fusing thallium with sulphur in absence of air.—Thallous sulphide is insoluble in water, in alkalies, or alkaline sulphides, and in potassic cyanide, soluble with difficulty in acetic acid, readily soluble in sulphuric and in nitric acid. The precipitated sulphide, when exposed to the air in a moist state, un- dergoes oxidation to sulphate. By heating in a current of hydrogen, thallous sulphide is reduced to metallic thallium. Thallic sulphide, TI2S//3, is prepared by fusing thallium with an excess of sulphur, expelling this excess at a low temperature with exclusion of air. It is a black amor- phous readily fusible substance. In warm weather it is soft like pitch, but below 12° C. INDIUM. 561 it is brittle. Hot dilute sulphuric acid dissolves it without separation of sulphur Thallic sulphide is the anhyride of a sulpho-acid, Tl3//Hs. The potassium salt of this acid, potassic sulphothallate, TIS//Ks, is obtained by fusing together 1 part of thallous sulphate with 6 parts of potassic carbonate and 6 parts of sulphur, extracting the cooled mass with water. The sulphothallate remains behind as a dark cochineal-red powder, consisting of microscopic quadratic plates. General Properties and Reactions of the Compounds of Thallium.—The salts of thallium are generally colorless. They have a disagreeable metallic taste and are poisonous. Zinc precipitates metallic thallium from solutions of the salts. Sulphuretted hydrogen precipitates neutral or slightly acid solutions of thallium salts only par- tially, and solutions containing an excess of a mineral acid not at all. Ammonic sulphide precipitates the whole of the thallium as brownish- black thallous sulphide, insoluble in alkaline sulphides. Thallous salts yield precipitates with the hydracids and soluble haloid salts (see p. 558). Thallium compounds impart to the non-luminous flame a magnificent emerald-green coloration. The spectrum of the thallium flame consists of one bright green line. INDIUM, In2? Atomic weight 113.4. Probable molecular weight 226.8. Sp. gr. 7.3 to 7.4. Fuses at 176° 0.(348.8° F.). Atomicity Evidence of atomicity: Indie chloride, In;"C]3. Indie hydrate, In'^Hoj. History.—lndium was discovered in the year 1863 by Reich and Richter in the zinc blende of Freiberg by means of the spectroscope. It received its name from the characteristic indigo-blue line which its spectrum exhibits. Occurrence.—Indium occurs in minute traces in various zinc blendes, particularly in that of Freiberg. The best source of the metal is the zinc from Freiberg, which contains on an average 0.05 per cent, of in- dium. Preparation.—Freiberg zinc is treated with a quantity of dilute hy- drochloric acid or sulphuric acid not quite sufficient to dissolve it, and is boiled with the liquid until gas ceases to be evolved. In this way any indium which may have gone into solution is precipitated upon the undissolved zinc. The spongy metallic mass which remains, and which, in addition to indium and zinc, usually contains lead, arsenic, cadmium, copper, tin, and iron, is dissolved in nitric acid and the solution boiled down with sulphuric acid until all the nitric acid is expelled, after which it is diluted with water, filtered from plumbic sulphate, and precipitated with a large excess of ammonia. The precipitate, which contains all the indium and iron, along with traces of the other metals present, is washed, dissolved in a small quantity of hydrochloric acid, and, after adding hydric sodic sulphite, boiled until the smell of sulphurous anhy- 562 INORGANIC CHEMISTRY. dride lias disappeared. In this way the whole of the indium is precip- itated as basic indie sulphite hydrate (see Indie Sulphite). It is, how- ever, still contaminated with lead, and, in order to free it from this impurity, it is dissolved in aqueous sulphurous acid, separated by filtration from undissolved plumbic sulphite and reprecipitated by boiling, when the pure basic sulphite is obtained. In order to prepare metallic indium, the sulphite is dissolved in hot hydrochloric acid, the solution precipitated with ammonia, and the precipitate of indie hydrate ignited and afterwards reduced in a current of hydrogen. Properties.—lndium is a non-crystalline, silver-white, lustrous metal. It is softer than lead and very malleable. It undergoes no change in air at ordinary temperatures, but when strongly heated in air, burns with a blue flame, giving off a brown smoke of indie oxide which con- denses on a cold surface as a yellow incrustation. Water, even at its boiling-point, is without action upon the metal. Dilute hydrochloric and sulphuric acids dissolve it slowly with evolution of hydrogen ; nitric acid dissolves it readily. COMPOUNDS OF INDIUM WITH THE HALOGENS. Indie chloride, InCl3.-—Molecular volume I I I.—This compound is prepared by heating the metal, or a mixture of the oxide with carbon, in a current of chlorine. It sublimes, without previous fusion, in soft, colorless laminae. It is deliquescent, and hisses when thrown into water, evolving great heat. The solution may be evaporated on the water-bath without decomposition, but on heating to a higher tem- perature to expel the last traces of water, hydrochloric acid is evolved and oxychlorides are formed. The bromide and iodide, which resemble the chloride in their properties, may be obtained by the direct union of their elements. COMPOUNDS OF INDIUM WITH OXYGEN AND HYDROXYL. In=o fIn° 1 Indie oxide, . . . .< O . O UnO | ln=o yO H Indie hydrate, . . . InHo3. In(—O—H X)—H Indie oxide, ln203, is formed as a pale yellow powder when the rnetai is burned in air or oxygen. It may be prepared by heating the hydrate or the nitrate. When heated it becomes reddish-brown, but recovers its original color on cooling.—By heating the oxide to 300° C. (572° F.) in a current of hydrogen a black powder is obtained which, unless COMPOUNDS OF INDIUM. 563 allowed to cool thoroughly before bringing it in contact with air, is pyrophoric. It appears to contain the lower oxide f1n"20.2. Indio hydrate, InIIo3, is obtained as a white gelatinous precipitate when ammonia is added to the solution of an indium salt. After dry- ing at 100° C. it forms a white horny mass, which at a higher temperature is converted into the oxide. The freshly precipitated hydrate is soluble in excess of potash and soda, but not in ammonia. It separates from the alkaline solution, slowly on standing, rapidly on boiling, or on the addition of amnionic chloride. OXY-SALTS OF INDIUM. Indie nitrate, crystallizes from its neutral aqueous solution with difficulty. From solutions containing an excess of nitric acid it is deposited in tufts of deliquescent needles. Indie sulphate, does not crystallize. By evaporation of its solution to dryness and heating to 100° C. (212° F.) it is obtained as a gummy mass having the composition S,jo«lno///2,90112; this when heated to 300° 0.(572° F.) is converted into the anhydrous salt. When a solution of indie sulphate containing an excess of sul- phuric acid is evaporated in vacuo, deliquescent crystals of dihydric di-indic tetrasulphate, S4OBHo2Tn///2,50H2, are deposited Diammonic di-indic tetrasulphate {lndiumammonia alum). S/VNH4O)2Ino///2 2401T2, crystallizes from mixed solutions of indie and ammonic sulphates in well-defined, colorless, regular octahedra. These dissolve in half their weight of water at 16° C., and in a quarter of their weight at 30° C. (86° F.). • At 36° C. (96.8° F.) the crystals fuse in their water of crystallization, and from the solution an octo-aquate is deposited in monoclinic crystals. Similar octo-aquates of the double sulphates of indium with sodium and potassium have also been prepared, but the aquates with 24 aq., or alums, are not known. Indie sulphite.—A basic indie sulphite of the formula SisOstOJnHobV OInITo2)2,50H2, tetrindic trisulphite hexahydrate, is deposited as a white crystalline powder when the solution of an indium salt is boiled with hydric sodic sulphite. It is insoluble in water, but readily soluble in acids. It dissolves in aqueous sulphurous acid, but is reprecip- itated from this solution by boiling. This property is turned to account in the separa- tion of indium from other metals (p. 562). COMPOUNDS OF INDIUM WITH SULPHUR. Indie sulphide, is obtained as a brown infusible mass by the direct union of its elements at a red heat. It is precipitated as an amorphous yellow powder when sulphuretted hydrogen is passed into the solution of an indium salt, but the precipita- tion is complete only when the liquid is kept neutral during the whole operation, or when sodic acetate has been added.—Ammonic sulphide produces in solutions of indium salts a white precipitate of a sulphhydrate which dissolves in an excess of the precip- itant on heating and separates out again on cooling.—Indie sulphide is the anhydride of a sulpho-acid, sulphindic acid, InS//Hs. Potassic sulphindate, InS//Ks, is prepared by heating together 1 part of indie oxide, 6 parts of potassic carbonate, and 6 parts of sulphur, at first at a gentle heat, afterwards more strongly. On extracting the cooled mass with water the sulphindate remains behind in the form of bright hyacinth-red, quadratic plates. Acids readily decompose it. General, Properties and Reactions of the Compounds of Indium.—The salts of indium with colorless acids are colorless. Zinc precipitates the metal from the solutions of its salts. Caustic alkalies precipitate white gelatinous indie hydrate, slightly soluble in excess, but reprecipitated on boiling. Sulphuretted hydrogen gives no precipitate in solutions containing an excess of mineral acid; from acetic acid solu- tion indie sulphide is precipitated. The same precipitate is produced 564 INORGANIC CHEMISTRY. by ammonia sulphide. The compounds of indium color the non-luminous flame dark-blue. The.spectrum exhibits an intense line in the indigo and a less marked line in the violet. CHAPTER XXXVI. TETRAD ELEMENTS. Section 11. ALUMINIUM, Al. Atomic weight 27. Molecular weight unknown. Sp. gr. 2.67. Fuses about 700° C. (1292° F.). Atomicity iv, but is always a pseudo-triad. Evidence of atomicity : analogy of iron and chromium. History.—Aluminium was first isolated by Wohler in the year 1827, but it was first obtained in the massive form by Deville in 1854. Occurrence.—Aluminium is, with the exception of oxygen and sili- con, the most abundant and widely distributed of the elements. It is always found in combination with oxygen. The oxide A1203 occurs as corundum, ruby, or sapphire ; the hydrate as hydrargillite, diaspora, and bauxite ; whilst the compound silicates of aluminium with other metals form a vast number of important minerals which are among the proxi- mate constituents of the various rocks (see Silicates, p. 319). Preparation.—Aluminium cannot be reduced directly from its oxide. It may be obtained by passsing the vapor of the chloride over heated potassium or sodium, and by the electrolysis of fused sodic aluminic chloride, Al2Cl6,2XaCI. On a large scale aluminium is prepared from bauxite, a native aluminic oxyhydrate of the formula Al2OHo4, in which a portion of the aluminium is isomorphously replaced by iron. This mineral contains about 50 per cent, of alumina. When heated with caustic soda in a reverberatory furnace the alumina forms sodic aluminate, Al202Xao2, which can be extracted with water, whilst the iron remains behind as insoluble ferric oxide. By passing carbonic an- hydride through the solution of the aluminate, aluminic hydrate is pre- cipitated, which by drying and heating is converted into alumina. This is mixed with powdered coal and common salt, and the mixture is made into balls, which are introduced into a fire-clay retort and heated to whiteness, while a current of dry chlorine is passed over them. The following reaction occurs: A1203 + 3C + 3C12 = A12C16 + 3CO. Alumina. Aluminic Carbonic chloride, oxide. The aluminic chloride volatilizes along with the sodic chloride as sodic aluminic chloride, which is condensed. It is now only necessary to ALUMINIUM. 565 reduce this double chloride with sodium. For this purpose the double chlorides is heated with sodium and cryolite (a native sodic aluraiuic fluoride of the formula Al2F6,6NaF), this last acting as a flux. In prac- tice 100 kilos, of the double chloride, 35 kilos, of sodium, and 40 kilos, of cryolite are employed in one operation. This mixture is heated, with gradual rise of temperature, on the hearth of a reverberatory furnace. The reduced aluminium fuses and collects on the hearth, whence it is drawn off and cast into ingots. The metal thus obtained contains iron and silicon. Aluminium may also be prepared from cryolite by mixing the finely powdered mineral with sodic and potassic chloride and heating the mixture in a crucible with sodium. The yield by this method is small and the metal impure. Properties.—Aluminium is a white metal, closely resembling zinc in color and hardness. It may be rolled into very thin foil or drawn into fine wire, and possesses at the same time great tenacity. It is most readily worked at a temperature between 100° C. (212° F.) and 150° C. (302° F.), It is not volatile at the highest temperatures that can be artificially produced. It is not oxidized by exposure to the air at ordinary temperatures, and is only superficially oxidized when fused in oxygen ; but in the form of foil or wire it may be burnt in oxygen, and emits a dazzling white light. Aluminium, when pure, does not decompose water, even at a red heat, but does so at 100° C. (212° F.) if the aluminium contains traces of sodium. It is soluble in caustic alkaline solutions and in hydrochloric and sulphuric acids. Nitric acid in all degrees of concentration is without action upon it. Organic acids alone scarcely attack it, but dissolve it rapidly in presence of chlorides, such as common salt; a fact which precludes its employment in the manufacture of utensils which have to come in contact with food. Uses.—lts lightness, tenacity, unalterability in air, and other valu- able properties, together with the abundance of its occurrence in nature, would probably render aluminium one of the most useful of metals, were it not for the difficulties attending its production in large cpiantity. For many purposes it might, for example, replace zinc and iron. At present, however, it is chiefly used in the manufacture of various physi- cal instruments, especially beams of delicate balances, in which a com- bination of lightness and inflexibility is essential. Aluminium bronze.—Aluminium forms alloys with most of the other metals; those with copper are the most important. Aluminium bronze is an alloy containing 90 parts of copper to 10 parts of aluminium, and is prepared by fusing the two metals together. Electrolytic copper is generally employed for this purpose, the quality of the alloy being de- pendent on the purity of the copper. The presence of iron is especially prejudicial. The alloy is brittle at first, but by repeated fusion becomes malleable. It has the color of gold, and resists the action of the air. It yields sharp castings, and is more easily worked than steel. Its tenacity is equal to that of cast steel, and more than twice that of gun-metal, whilst its resistance to flexure is thrice that of gun-metal. It is employed in the manufacture of imitation gold ornaments and of physical instruments. Alloys of aluminium with silver and with tin have also found appli- cation in the arts. 566 INORGANIC CHEMISTRY. COMPOUNDS OF ALUMINIUM WITH THE HALOGENS. Aluminic chloride, AI2C16.—Molecular volume I I I.—This com- pound is formed when aluminium is heated in chlorine. (Preparation, see p. 56-1.)—If contaminated with ferric chloride, which imparts to it a yellow color, it may be purified by mixing it with iron filings, or better with aluminium filings, and re-subliming. In either case the ferric chloride is converted into the much less volatile ferrous chloride. Aluminic chloride when perfectly pure is a white crystalline substance. It sublimes readily at ordinary pressures without fusing, but may be fused under the pressure of its own vapor, or when rapidly heated in large quantity. By sublimation it is sometimes obtained in hexagonal tabular crystals. It attracts moisture from the air, and evolves hydro- chloric acid. The solution of the metal or the oxide in hydrochloric acid deposits on concentration colorless needle-shaped crystals of the aquate A12C16,120H2, which on heating are decomposed into water, hydrochloric acid, and alumina. Aluminic chloride forms a large num- ber of compounds with the chlorides of other elements. Fotassic alu- minic chloride, A12C18,2KC1, and sodic aluminic chloride, Al2Cl6,2NaCI, are formed when aluminic chloride is heated with potassic and sodic chlorides. The sodium compound fuses without decomposition at 185° C. (365° F.), and is volatile at a red heat. It is employed in the prepa- ration of aluminium. Aluminic bromide, Al2Br6.- Molecular volume [ | | —Aluminium and bromine unite with incandescence to form this compound. It may be most readily obtained by pass- ing bromine vapor over a red-hot mixture of alumina and carbon. It may be purified by repeated sublimation with aluminium in a sealed tube. It forms deliquescent, colorless, lustrous laminae, fusing at 90° C. (194° F.), and boiling between 265° C. (509° F.) and 270° C. (518° F.). Concentrated aqueous solutions deposit colorless needles of the aquate Al2Br6 120H2, which on heating are decomposed like the cor- responding chlorine compound. Aluminic bromide forms fusible double bromides with the bromides of the alkali metals: thus, potassic aluminic bromide, Al2Br6,2K8r. Aluminic iodide, A1216.—Molecular volume m- This compound is formed with in- candescence when aluminium and iodine are cautiously heated together in a sealed tube. It is also formed when argentic iodide is heated with aluminium filings.— Aluminic iodide is a white crystalline mass, fusing at 185° C. (365° F.), and boiling at 360° C. (680° F.). Its vapor is combustible, and forms an explosive mixture with air. The products of combustion are alumina and iodine. It is decomposed in the same way when heated in contact with air. When exposed to the air it fumes and deli- quesces. It is readily soluble in water, alcohol, and bisulphide of carbon. It forms an aquate, Al2Ie,120142, and unites with the alkaline bromides to form double iodides, all of which compounds closely resemble the corresponding chlorides and bromides. Aluminic fluoride, Al2F6, is formed by the action of gaseous or aqueous hydrofluoric acid upon alumina or aluminic hydrate. At a bright red heat it sublimes in colorless rhombohedra, closely approximating to cubes. It is insoluble in water, and is not de- composed by acids.—Aluminic fluoride forms insoluble double fluorides with the fluo- rides of the alkali metals. The most important is aluminic sodic fluoride, Al2Ffi,6NaF, which occurs as the mineral cryolite in enormous deposits on the coast of Greenland. It may be artificially obtained by fusing together its component fluorides. It forms a white, translucent mass. It is.decomposed by sulphuric acid with evolution of hydro- fluoric acid. Boiling with caustic alkalies, or with calcic hydrate and water, also decomposes it. In the decomposition with calcic hydrate insoluble calcic fluoride is formed, whilst sodic aluminate goes into solution ; AL Fe,6N aF 6CaHo2 = 6CaF2 + Al2Nao6 -f- CO H2. Sodic aluminic Calcic Calcic Sodic Water, fluoride. hydrate. fluoride. aluminate. 567 COMPOUNDS OF ALUMINIUM. On this reaction is based an industrial process for the preparation of soda and aluminium salts from cryolite. COMPOUNDS OF ALUMINIUM WITH OXYGEN AND HYDROXYL. Aluminic oxide [Alumina), f aioq [ Al(j ‘ O 0=M—A\=0. H—O O—H Aluminic hydrate [Hy dr argillite), < f A1Ho3 [ A1Ho3‘ | H—O—Al—Al—O—H. 1 1 r AlOHo [ AlOHo- H—O O—H Aluminicoxydihydrate [Diaspore) j 0=A1—A1=0 1 i 0 O 1 1 H H Aluminic oxide [Alumina), A1203.—This oxide occurs native in hexagonal crystals, sometimes colorless, sometimes variously colored owing to the presence of other oxides. Crystallized alumina is harder than any known substance with the exception of the diamond and crys- tallized boron. The colorless or gray crystals are known as corundum; the red crystals, the color of which is due to chromium, constitute the gem ruby; whilst sapphires are crystals of alumina colored blue, probably by cobalt. In an impure state, contaminated with iron and silica, alumina occurs in large masses as emery. The latter mineral, when powdered and levigated, is employed for grinding and polishing surfaces of glass and metal, purposes for which from its hardness it is admirably suited. Alumina is obtained as a white amorphous powder by heating the hydrate or ammonia alum; in the latter case it is diffi- cult to expel the last traces of sulphuric acid. It may be obtained in the crystallized condition by the action of aluminic fluoride upon boric anhydride at a high temperature. Fremy arid Feil have prepared crystallized alumina on a large scale by heating together equal weights of alumina and red-lead in a clay crucible to bright redness for a con- siderable time, sometimes as much as twenty days. The cooled mass consisted of two layers: one a vitreous mass of plumbic silicate, the silica of which had been derived from the material of the crucible; the other crystalline, and containing cavities which were filled with well-formed crystals of corundum. By the addition of from 2 to 3 per cent, of potassic dichromate to the above mixture crystals of ruby were obtained ; the color of sapphires was produced by adding a small quantity of cobaltous oxide, together with a trace of potassic dichromate. By heating a mixture of equal weights of alumina and baric fluoride, with a small quantity of potassic dichromate for a length of time to a very high temperature in a glass furnace, magnificent crystals of ruby 568 INORGANIC CHEMISTRY. were obtained. The reaction in this case depends upon the formation of aluminic fluoride which is then decomposed by the furnace gases. The crystals of ruby are deposited in the upper part of the crucible.— Crystallized or strongly ignited alumina is insoluble in acids at ordi- nary pressures, but dissolves in concentrated sulphuric acid when heated with it in sealed tubes. It is also attacked by fusion with hydric po- tassic sulphate or potassic hydrate, after which treatment it dissolves in water. Alumina is fusible in the oxyhydrogen flame. Aluminic hydrate, Al2Ho6, occurs as hydrargillite in small hexagonal crystals. When ammonia is added to the solution of an aluminium salt a white gelatinous precipitate is formed, which after drying at ordinary temperatures has the composition AIJIo6201I2. This when heated slightly above 300° C., is converted into aluminic oxydihydrate, A12G2Ho2, a compound which occurs in nature as the mineral diaspore in rhombic crystals. An aluminic oxyhydrate, corresponding with the formula Al2OHo4, aluminic oxytetrahydrate, occurs as the mineral bauxite, but a portion of the aluminium in this compound is isomor- phously replaced by iron. All the aluminic hydrates are converted into the oxide by heating.—Aluminic hydrate is insoluble in ammonia, but when freshly precipitated dissolves readily in acids and in solutions of potassic and sodic hydrate. When dried by a moderate warmth, or when allowed to stand under water, it becomes difficultly soluble in acids and alkalies.—Freshly precipitated aluminic hydrate dissolves in a solution of aluminic chloride, and if the liquid thus obtained be sub- jected to dialysis, hydrochloric acid passes through the dialyser, till at last only a neutral tasteless solution of colloidal aluminic hydrate remains. This soluble modification of aluminic hydrate is very un- stable: the solution coagulates after standing for some days, and the same change takes place immediately on the addition of traces of acids, alkalies, or salts. Aluminic hydrate possesses the property of pre- cipitating organic coloring matters from their solutions. Upon this property the application of the salts of alumina as mordants in dyeing and in the preparation of the so-called lakes depends. Aluminates.—Aluminic oxydihydrate behaves towards stronger bases like a weak acid. Its salts, in which both the hydrogen-atoms of the oxydihydrate are replaced by metal, are known as aluminates. The aluminates of potassium and sodium are pre- pared by dissolving aluminic hydrate in caustic potash or soda; by evaporation in vacuo, the potassic alurainate may be obtained in hard lustrous crystals of the formula Al202Ko2,301b- Sodic aluminate, has not been obtained in the crystal- lized state. It is used as a mordant, Beryllie aluminate, AI,02Beo//, occurs in nature as the mineral chrysoberyl in green rhombic crystals. The aluminates of the metals of the magnesium group occur in nature as the spindles, crystallized in forms belong- ing to the regular system. Examples of these are; magnesic aluminate or spindle, Al20,2Mgo//, and zincie aluminate or zinc spindle, The two latter com- pounds may be prepared artificially by passing the vapor of aluminic chloride over strongly heated magnesia or zincic oxide, or by heating alumina and boric anhydride with these oxides to a white heat for several days. OXY-SALTS OF ALUMINIUM. Aluminic nitrate, Ne012(/Al///206)''',lBOll2> crystallizes from a concentrated solution of the hydrate in nitric acid in deliquescent monoclinic prisms. On heating to 150° C. (302° F.) the salt is decomposed, leaving a residue of alumina. It is employed in calico-printing as a mordant. 569 COMPOUNDS OF ALUMINIUM. Aluminic sulphate, S3O6(/A1"/2O6)vI,IBOH[2, occurs as the mineral Iceramohalite. It is prepared on a large scale by dissolving aluminic hydrate, obtained from cryolite or bauxite and as free from iron as possible, in sulphuric acid; or by decomposing china clay, a hydrated aluminic silicate, with sulphuric acid. The solution is evaporated till it solidifies on cooling. A soft mass is thus obtained which is cut into blocks. It is difficultly crystallizable, and forms thin, flexible, na- creous laminae. It dissolves in twice its weight of cold water. When heated, it first fuses in its water of crystallization, then swells up, and is converted into a white porous mass of the anhydrous salt. Aluminic sulphate is employed as a mordant and in weighting paper.—Basic sul- phates are formed when a solution of the normal sulphate is precipitated with an insufficiency of ammonia, or by boiling its solution with the freshly precipitated hydrate. A compound of this class, aluminic sul- phate tetrahydrate, SO2(/Al///2O2Ho4)//,70H2, occurs in nature as the mineral aluminite. The Alums. Among the most important salts of alumina are the double sul- phates which it forms with the alkalies, known as the alums. Of these the principal are potash alum or common alum, dipotassic S02Ko-. gQ I aluminic tetrasulphate, gQ2 ('AI///206)v 1,240112, and ammonia alum, so2kOJ in which the potassium of the preceding compound is replaced by ammonium. The object of preparing these salts, which are exten- sively used by the dyer and calico-printer, is to obtain compounds of alumina in a very pure form, and especially as free from iron as possi- ble. The alumina is alone valuable. The name alum is not restricted to compounds of alumina: it is employed to desig- nate a class of tetrasulphates which, like potash alum, contain in their molecule two atoms of a monad metal (or the equivalent of a monad metal, such as NH4) together with one hexadic metallic group—of which /Al///2 may be taken as a type—and which crystallizes with 24 aq. in regular octahedra. Almost any monad metal may enter into the composition of an alum; thus, besides the alums above mentioned, alums have been prepared containing sodium, caesium, rubidium, silver, and thallium. The hexadic group /Al///2 may be replaced by the isomorphous groups /Cr///2, /Fe///2, and /Mn///2. It even appears that this group of two pseudo-triads may be replaced by two true triads: thus an ammonia indium alum has been prepax-ed containing the hexadic group In2 (p. 563). The following system of nomenclature is as a rule applied to these compounds. It the monad metal be potassium, the name of this metal is not introduced into the name of the compound : thus chrome alum means potassium chromium alum. If the hexadic group be /Al///2, aluminium is not named : thus by ammonia alum is understood ammo- nia aluminium alum. If the alum contain neither potash nor aluminium, both metals present must be named : thus, ammonia chrome alum. Selenic acid forms a similar series of alums. These may be regarded as sul- phuric alums in which sulphur has been replaced by the isomorphous selenium. Se02Ko-j The potash alum of this series has the formula (/Al///206)t1,240H2- Se02KoJ A class of pseudo alums also exists in which the two monad atoms are replaced by one dyad atom. These pseudo alums also contain 24 aq. in the molecule, but do not crystallize in the regular system (see Pseudo Alums). 570 INORGANIC CHEMISTRY. A solution containing two or more octahedral alums deposits octahedral crystals, in which the various alums present may be contained in any proportion. Potash alum crystallizes from mixed solutions of aluminic and potassic sulphates. It is formed in nature, especially in volcanic districts, by the action of sulphurous acid and oxygen upon rocks containing potassic and aluminic silicates. In the neighborhood of Naples and at Solfatara it occurs in quantity sufficient to render its extraction profitable. Large quantities of very pure alum, the so-called Roman alum, are obtained from the mineral alum stone or alunite, a basic double silicate of potash and alumina of the formula SOH6KO('A1"'«HoA)"' 2 (7A1"'2H0403)'/ , which occurs at Tolfa and in Hungary. ioSoko('Ai;"°H°A)'" The mineral is mixed with fuel and roasted, either in heaps or in kilns, after which it is moistened and exposed to the air for several weeks. The mass gradually disintegrates, and is then extracted with water, when alum goes into solution and alumina remains behind. The liquid is concentrated and allowed to crystallize.—Alum is, however, more frequently prepared from alum shale, a bituminous shale containing iron pyrites disseminated through its mass. The shale is exposed in heaps to the air, by which means the iron py- rites (FeS"2) is gradually oxidized to ferrous sulphate and sulphuric acid, the latter of which then decomposes the aluminic silicate present in the shale. The process is generally shortened by first roasting the shale, in order to effect a partial oxidation, after which the roasted shale is moistened and exposed to the air as above described. The oxidized shale is lixiviated with water and the solution evaporated. A considerable quantity of the ferrous sulphate present crystallizes out and is removed. If, however, the shale has been- exposed to the air for a sufficient length of time, the ferrous sulphate is oxi- dized to ferric sulphate, the presence of which is less objectionable. The concentrated mother liquor containing aluminic sulphate is now heated to boiling, and solid potassic sulphate is dissolved in it. The potassic sulphate combines with the alumnic sulphate to form alum. If any considerable quantity of ferric sulphate is present it is advantageous to add, along with the potassic sulphate, a quan- tity of potassic chloride equivalent to the ferric sulphate, the two latter salts yielding by double decomposition potassic sulphate and the very soluble ferric chloride. The presence of ferrous sulphate is objectionable, as a loss of potassium salt is occasioned by the ( SQ2Ko formation of ferrous dipotassic disulphate, < Feo" . The hot solution, f SG2Ko which now contains the alum, is well stirred till cold. In this way the alum is deposited in small crystals, which are less apt to retain impurities from the mother liquor than the large crystals which would be formed were the liquid permitted to cool undisturbed. The small crystals, known as alum meal, are washed with cold water, dis- solved in boiling water, and the solution allowed to crystallize in large 571 COMPOUNDS OF ALUMINIUM. barrels with movable staves, which are afterwards taken to pieces in order to remove the large crystals of alum which line their sides.—Alum crystallizes in large colorless transparent regular octahedra, which as a rule also exhibit subordinate cubical faces. From solutions containing free caustic alkali, or basic alum, the alum crystallizes by spontaneous evaporation in cubical crystals, which have exactly the same composi- tion as octahedral alum. The crystallized alum is soluble in 7 parts of water at 20° C. (68° F.), and in less than £ part at 100° C. (212° F.). The solution has a faint acid reaction and a sweet astringent taste. The crystals are insoluble in alcohol. When heated they fuse in their water of crystallization, which is expelled by continued heating, leaving a white porous mass known as burnt alum. This dissolves slowly in water. Anhydrous alum may be obtained in six-sided crystals by fusing alumina with hydric potassic sulphate, and removing the excess of this salt from the fused mass with hot water. Ammonia alum (diammonie aluminic tetrasulphate), S02(NH40)-| SO* ('AI" W,240H2 SC&NH.O)J This compound was formerly prepared from alum shale by methods similar to those employed in the manufacture of potash alum. The roasted shale was treated with sulphuric acid, and into the acid solution of aluminic sulphate, ammonia, obtained from the amraoniacal liquors of the gas-works, was passed. The alum was purified by crystalliza- tion. Since the introduction of cheap potash salts from the Stassfurt beds, the manufacture of ammonia alum in England has practically ceased,—Ammonia alum crystallizes in large colorless octahedral crystals, in appearance indistinguishable from the potash salt. Its solubility is also almost the same as that of potash alum. Soda alum, S, 08 Nao2 (/ AIx 772 0(i)v 1,210 tl ,is difficult to purify on account of its great solubility. It dissolves in its own weight of water at ordinary temperatures. It is not manufactured. Aluminic phosphates.—The normal orthophosphate, aluminic PO diphosphate, pq(/Al///206)tl, is obtained as a hydrated gelatinous pre- cipitate when hydric disodic phosphate is added to the neutral solution of an aluminium salt. It is soluble in alkalies, but not in ammonia; and in mineral acids, but not in acetic acid.—Various basic phos- phates of alumina occur in nature. The mineral wavellite, which forms rhombic crystals or radiating masses, is a basic phosphate of the formula P,O(’A 1 ///2Ofi)V13,120H2. Calaite, which when colored greenish blue by copper constitutes the gem oriental turquoise, has the formula PO('AF"2Ho3O3)'",OH2. Aluminic silicates.—The silicates of alumina, both simple and compound, form a large class of important minerals. A detailed de- scription of these belongs rather to mineralogy than to chemistry; but the names and formulae of some of the more important may be here given. 572 INORGANIC CHEMISTRY. Andalasitc(phiastolitecyanite,fhro- \ a V” n \" . lite, sillimanite, j 2 4' * SV/AJ/// Q \vi Bucholzite (xenotime), .... Si), * ~,,2n6(vi. gjV 2^o/ Miloschine, SiHo2(/Al"'2Ho402)" Alophane, SiEto2(/Al//'2Ho402)//,(2 or 4)0H2, Colly rite, Ho60)2,40H2. Porcelain clay of Passau, . . . g^^'AF'^HogOJ^. ( SiHo-, Kaolin {porcelain clay, china clay), < O (/Al///2Ho204)iv. t SiHoJ SiPIo2n Kazoumoffskin, SiHo^'AP'^CUA SiHo2J! Wnrtlij+P sin^Ab'^HoOg)*. VVormite? gJU(/Al"/2Ho05)7 SiHo,-, Cimolite (kaolin of Ellenbogen), . (/A1"'206'206)Yi. SiHo3-J BiOHo-| Agalmatolite, ('AP'^Og)™. SiOHo-1 Malthacite, Si8011Ho4(/Al///206)vl. Malthacite, f SiHo2Nao )°- Analcime, i g|(/Al///206)vi. \o ( SiHo2Nao SiOKao . ( SiO 1 I \o I I Albite, rl8H('Al",*°«)Tl- • \o 1 l (SiO—J I SiONa—J SiONa— Lepidolite, Si908Ko2Lio4('Al"'206)Ti2('Al"'2F402)". Petalite, Si30O45Nao2Lio4(/Al///2O6)vi4. PORCELAIN AND POTTERY. 573 Spoduraene, . . . . Si1A5LH('Ai"WV Wernerite, . . . . Si2Cao//(/Al"/2Oc)vi. Prehnite, . . . . . Si3H02Ca0"2(/Al///206)vi. / f'Al'"Oa * Zoisite, .... . . Si4Cao" { { O \\'AV"aoJ Saponite, . . . . . Si7(MgO//6Ho10(/Al///2O6)vi. Topaz, .... . . Sis(/Al///205F)T('Al///201F2)iT(/Al///204F)///. (See also Silicates, p. 319.) Ultramarine. Various native double silicates of aluminium with other metals con- tain sulphur as an essential constituent. One of these, a double silicate and sulphide of aluminium and sodium, forms the mineral lapis lazuli, prized for its splendid blue color, and employed as a material for vases and inlaid or mosaic work. It is sometimes found crystallized in dode- cahedra, but generally occurs massive. It has not as yet been found possible to express the composition of this mineral by means of a formula. The powdered mineral was formerly employed as a valuable blue pigment under the name of ultramarine, a substance which is now prepared artificially. For this purpose china clay {infra) is heated in crucibles along with sodic sulphate and charcoal. The sodic sulphate is reduced to sodic sulphide, which then combines with the aluminic silicate. The product is a white mass, which, however, speedily be- comes green. This substance, known as green ultramarine, is also employed as a pigment. When green ultramarine is heated with sul- phur, allowing the sulphur to burn off in air, it assumes a blue color, and is thus converted into the ordinary blue ultramarine of commerce. The same change is effected when green ultramarine is heated with am- monic chloride, or when chlorine is passed over it, but the sulphur method is employed in practice. No difference in chemical composition can be detected between the green and the blue modification. When ultramarine is treated with hydrochloric acid, it is decolorized with evo- lution of sulphuretted hydrogen and separation of amorphous silicic acid. It is used in paper-staining, in calico-printing, and as an oil paint. Porcelain and Pottery. Porcelain and pottery in all their forms are manufactured primarily from clay, an aluminic silicate. This material possesses sufficient plasticity to allow of its being moulded into any desired form, whilst by the action of heat it is rendered .sufficiently hard and tenacious to resist the wear of every-day use. The I Si Ho-, purest clay is kaolin or china clay, < O (/Al///2H0204)iy, which is t SiHoJ formed from felspar, Sig08Ko2(b41///206)vi, by weathering, the grad- 574 INORGANIC CHEMISTRY. nal action of water removing the potash’ together with a portion of the silicic acid, and leaving an alurainic silicate. Kaolin sometimes occurs in six-sided tablets, but generally forms a white or yellowish-white mass. The commoner clays consist of kaolin with various impurities —calcic and magnesic carbonates, ferric oxide, sand, and organic mat- ter. Kaolin does not fuse when heated, but bakes together into a hard porous mass; in order, therefore, to increase the durability of utensils manufactured from it, the kaolin is mixed with some fusible material, technically known as a frit, which by its fusion binds the whole together. The materials added are ground feldspar, quartz-sand, chalk, gypsum, bone-ash, and sodic or potassic carbonate—the nature of the frit varying with the quality of the ware required. The mate- rials are carefully ground under water and mixed. The mixing is an operation of great nicety, inasmuch as it is necessary to preserve the same composition of the mixture for a given kind of ware; and as the composition of the clay is apt to vary, this constancy of composition can only be attained by suitably varying the proportions of the other ingredients: thus, if the clay should happen to contain a larger quan- tity of silicia, less quartz-sand will have to be added, and so on. The presence of organic matter is objectionable, as organic substances disen- gage gas during the firing, and are thus liable to spoil the work. By allowing the mixture to stand in a moist state for a*considerable length of time, the organic matter undergoes putrefaction, and is thus got rid of. The plastic mass is then moulded into the required form, either on the potter’s wheel, or by means of moulds. The articles are then allowed to dry at ordinary temperatures, and are then in some cases subjected to a preliminary process of firing at a relatively low tempera- ture, known as baking, after which they are glazed. The glaze is of various kinds, according to the nature and quality of the ware; but in every case it consists of some material which in the subsequent firing fuses, and imparts to the porous ware a smooth vitreoussurfa.ee, imper- meable to liquids. The glaze is generally employed in the form of a fine powder, which is either suspended in water, into which the baked articles are dipped, or is dusted upon their surface. Another mode of glazing consists in volatilizing in the porcelain kiln some material which is thus deposited on the surface of the articles, and forms with the silica which they contain a fusible glaze (salt-glazing). The finer sorts of porcelain and earthenware are not exposed to the direct action of the flame in firing, but are inclosed in fire-clay crucibles, known as saggers, by which means they are protected from the action of the smoke and ash. The porcelain kiln consists of a tall reverberatory furnace, divided usually into three stories or floors, through which the flame passes. The upper story is employed for baking, the two lower for firing. The firing is continued during eighteen hours, after which the kiln is allowed to cool slowly for three or four days in order to anneal the ware. Porcelain or China.—This is the finest description of ware. It was manufactured in China before the Christian era; but the art of making true porcelain was not discovered in Europe till the commencement of the 18th century. There are two chief classes of porcelain: hard'por- celain, to which class the Chinese, German and Sevres porcelain belong; PORCELAIN AND POTTERY. 575 and tender 'porcelain, produced especially in England. Great care has to be exercised in the selection of the materials in order that the result- ing porcelain may be colorless. The presence of ferric oxide and organic matter is to be avoided. The purest kaolin forms the basis of all porcelain ; and upon the nature of the frit the difference in proper- ties of the various kinds of porcelain depends. In the case of hard porcelain, the frit consists of calcic and potassic silicates : thus the paste employed at Sevres for ornamental porcelain has the composition : Washed kaolin, 62 parts; chalk, 4 ; quartz-sand, 17 ; felspar, 17. The glaze for this porcelain consists of a mixture of felspar and quartz. In the case of English porcelain, a frit consisting of bone-ash or a mineral phosphate, together with borax, is employed. This frit is much more fusible than the preceding, and the porcelain thus obtained is softer. This porcelain is glazed with an easily fusible mixture of bone-ash, plumbic oxide, potashes, sand, and borax. Tender porcelain must be baked before applying the glaze, and then fired; hard porcelain i£ sometimes glazed after drying at ordinary temperatures. The reason for this difference in treatment is to be found in tiie fact that in the case of tender porcelain the glaze is very much more fusible than the mass, whilst with hard porcelain this is not the case. Porcelain forms a white, translucent, homogeneous mass. Hard porcelain resists sudden changes of temperature and the action of acids and alkalies much better than glass, and is for this reason employed in the manufacture of laboratory vessels. Stoneware differs from porcelain in being always opaque and generally more or less colored. The materials employed are not so pure, and generally contain ferric oxide. It is more fusible than porcelain. In order to glaze this ware, the process known as salt-glazing is employed. The articles to be glazed are dipped in sand and water, and then grad- ually heated to a very high temperature in the kiln. A quantity of common salt is then thrown into the kiln. The salt volatilizes, forming with the sand a fusible sodic silicate, which combines with the other silicates present to yield a glass or glaze, and coats the ware, rendering it impervious to water. The explanation of the process is as follows: Silicic anhydride alone is not capable of decomposing sodic chloride at any temperature; but when the two substances are strongly heated together in presence of the vapor of water, hydrochloric acid is ex- pelled and sodic silicate formed: Si02 + 2NaCI + 0H2 = SiONao2 + 2HCI. Silicic Sodic Water. Sodic Hydrochloric anhydride. chloride. silicate. acid. The water is furnished by the combustion of the fuel. At the same time another portion of sodic chloride acts upon the ferric silicate con- tained in the clay, yielding sodic silicate and volatile ferric chloride: Si303(/Fe'"206)Ti -f 6NaCI = 3SiONao2 + Fe2Cl6. Ferric silicate. Sodic chloride. Sodic silicate. Ferric chloride. The iron present on the outer surface of the ware is thus removed. 576 INORGANIC CHEMISTRY. Earthenware.—This ware differs from the two preceding varieties, inasmuch as no fusion or vitrification occurs during firing, and the body of the ware remains porous. A piece of unglazed earthenware adheres to the tongue. In the manufacture of fine earthenware a paste is employed consisting of a mixture of fine plastic clay and ground flints. This mass burns white on firing, and is afterwards glazed with an opaque lead glaze. Common earthenware is prepared from inferior clay. In the manufacture of common 'pottery ware—bricks, flower-pots, etc. —impure clays are employed. The color, red or yellow, is due to the presence of ferric and other oxides in the clay. Fire-bricks, melting crucibles, and other articles which are required to resist a high temperature, are prepared from a pure clay rich in silica. In order to lessen the shrinkage which this clay suffers in firing, a quantity of finely powdered burnt clay (broken pots of the same ma- terial) is added. COMPOUND OF ALUMINIUM WITH SULPHUR. Aluminic sulphide, Al'2S//3, is formed as a black mass, which acquires metallic lustre under the burnisher, by the union of aluminium with suiphur at a red heat, and may also be obtained as a white vitreous substance by passing the vapor of carbonic disul- phide over alumina heated to whiteness : 2AIA + 3CS"s = 2Al2S//3 + 3C02. Aluminic Carbonic Aluminie Carbonic oxide. disulphide. sulphide. anhydride. Water decomposes it, yielding aluminic hydrate and sulphuretted hydrogen.—Alka- line sulphides and sulphhydrates precipitate aluminic hydrate from solutions of alumi- nium salts. General Properties and Fractions of the Compounds of Aluminium.—The salts of aluminium with colorless acids are colorless. They have a sweet but very astringent taste. Their solutions redden blue litmus. Caustic alkalies, ammonia, ammonic carbonate, baric car- bonate, and ammonic sulphide, all precipitate aluminic hydrate—in the case of the carbonates with evolution of carbonic anhydride, and in the case of ammonic sulphide with evolution of sulphuretted hydrogen. The precipitate is readily soluble in caustic alkalies, only very spar- ingly soluble in ammonia. If aluminium compounds be ignited before the blowpipe, then moistened with cobaltous nitrate and again ignited, a pale blue mass (Thenard’s blue) is obtained. Aluminium compounds do not color the non-lumiuous flame. The spark-spectrum of alumi- nium is very complex. GALLIUM, Ga. Atomic weight 68.8. Molecular weight unknown. Bp.gr. 5.9, Fuses at 30.1° C. (86.2° F.). Atomicity IV, but is always a pseudo-triad. Evidence of atomicity: analogy with aluminium. History.—Gallium was discovered in 1875 by Lecoq de Boisbaudran with the aid of the spectroscope. GALLIUM. 577 Occurrence.—Gallium is one of the rarest elements. It occurs in minute traces in the zinc blende from Pierrefitte in the Pyrenees, from Austria and from Bensberg. The blende from the latter source, which is the richest in gallium, contains only 0.0016 per cent, of this metal. Extraction.—The zinc ores containing gallium are dissolved in acid —hydrochloric acid, sulphuric acid, or aqua-regia, according to the nature of the ore—and the solution is partially precipitated with metallic zinc. The gallium, along with the other foreign metals originally con- tained in the zinc ore, is precipitated upon the zinc. The precipitate is redissolved in hydrochloric acid and the solution again treated with metallic zinc. This precipitate is again dissolved in hydrochloric acid, and sulphuretted hydrogen is passed into the solution. The liquid is filtered from sulphides, and, after expelling the sulphuretted hydrogen by boiling, fractionally precipitated with sodic carbonate as long as spectroscopic examination shows the presence of gallium in the fractions. The various fractions are dissolved in sulphuric acid, the solution evaporated to dryness, and the residue heated so as to expel the excess of acid. On treating with hot water, basic gallic sulphate separates, and must be filtered off hot. The basic sulphate is dissolved in the smallest possible quantity of sulphuric acid, and, after adding amnionic acetate, the gallium is precipitated from the solution as sulphide by means of sul- phuretted hydrogen. In order to obtain metallic gallium the sulphide is again dissolved in sulphuric acid and, after adding an excess of caustic potash, in which the gallic hydrate is soluble, the liquid is subjected to electrolysis, employing electrodes of platinum. The electrolytically deposited gallium is washed with dilute nitric acid, and is then pure. Properties.—Gallium is a bluish-white metal of sp. gr. 5.9. It fuses at the low temperature of 30.1° C. (86.2° P.), and remains for a long time in a state of superfusion, even at 0° C., but when touched with a piece of the solid metal instantly solidifies in pyramidal crystals. The metal when fused is silver-white and more lustrous than in the solid state. It dissolves with evolution of hydrogen in hydrochloric acid and in caustic potash. Nitric acid is almost without action upon it in the cold, but dissolves it on heating. When a solution of gallic chloride is warmed with metallic zinc, gallic oxide or a basic salt is precipitated. COMPOUNDS OF GALLIUM. Gallic chloride, /Ga///2C16, forms soluble, deliquescent, colorless needles. Excess of water decomposes it with separation of an oxy-chloride. Gallic oxide, /Ga///203, is a white precipitate insoluble in water, but soluble in caus- tic alkalies and in ammonia. S02-n Gallic sulphate, S02-(/Ga///206)VI, is very soluble. From mixed solutions of SO,-- this salt with amnionic sulphate, regular crystals of ammonium gallium alum, S02Amo-j qq2 (/Ga///206)v 1,210112, are deposited. S02Amo-^ General Properties and Reactions of the Compounds of Gallium.—Gallium is most readily recognized by means of its spark 578 INORGANIC CHEMISTRY. spectrum, which consists of two lines in the violet. The flame spec- trum shows only one of these lines, and that but faintly. The other characteristic properties of the gallium compounds are given above. CHAPTER XXXYII. METALS OF THE RARE EARTHS. The metals of this group occur, generally together, in a few rare minerals. Their separation is a matter of extreme difficulty owing to the similarity of their compounds. Indeed it is doubtful in the case of most of them whether pure compounds have ever been obtained—a fact pointed to by the discrepant results arrived at by careful experimenters in the determination of the atomic weights of these elements. The most important metals of this group are cerium, lanthanum, didyrnium, yttrium, and erbium. TETRAD ELEMENTS, Section 111. CERIUM, Ce. Atomic weight = 140.5. Molecular weight unknown. Sp. gr. 6.728. Atomicity iY, also a pseudo-triad. Evidence of atomicity: f gpr//n Cerous chloride, . . . . j Ce'^Cl3, Ceric oxide, CelT02. History.—Ceria was discovered by Klaproth in 1803, but was first recognized as the oxide of a new metal by Berzelius and Hisinger. Occurrence.—Cerium always occurs together with lanthanum and didyrnium. The most abundant source of these three metals is the mineral cerite, in which they occur as silicates. They also occur in monazite as phosphates, and in fiuocerite as fluorides. Preparation.—Separation of Cerium, Lanthanum, and Didyrnium. —Finely powdered cerite is mixed with concentrated sulphuric acid so as to form a thick cream, and the mixture is heated in a Hessian cruci- ble—first gently in order to expel the acid, finally to low redness. The cooled contents of the crucible are powdered and added in small por- tions at a time to water at 0° C., great care being taken to avoid any rise of temperature. The solution, after filtering from sand and other in- soluble matters, is treated with sulphuretted hydrogen in order to pre- cipitate copper, bismuth, molybdenum, and lead. After removing these, chlorine is passed in to reoxidize the iron and, after acidifying with CERIUM 579 hydrochloric acid, oxalic acid is added in excess. In this way the cerium, lanthanum, and didymium—together with any yttrium and erbium, if present—are precipitated as oxalates. The precipitate is strongly ignited, by which means the oxalates are converted into oxides. These are dissolved in nitric acid and the solution evaporated to a syrup. The syrupy solution is then diluted with water and poured into a large excess of boiling water containing 2 c.c. of sulphuric acid to the litre. The cerium is thus precipitated as a basic ceric sulphate. This precipi- tate is dissolved in sulphuric acid and re-precipitated as basic sulphate by again pouring into boiling water, repeating these operations until the solution of the cerium salt in sulphuric acid no longer shows the absorption spectrum of didymium. The cerium compound may then be regarded as pure. In order to obtain the lanthanum and didymium from the filtrate from the first precipitation of basic ceric sulphate, this liquid is first boiled with pulverized magnesite, which precipitates the rest of the cerium as oxide, whilst the lanthanum and didymium remain in solu- tion. The lanthanum and didymium are then precipitated by the ad- dition of oxalic acid to the solution acidified with hydrochloric acid, the oxalates are converted as above into oxides, these are dissolved in sul- phuric acid, the solution is evaporated to dryness, and the salt heated to low redness. The anhydrous sulphates thus obtained are dissolved in five times their weight of ice-cold water, adding the salt to the water in small quantities at a time, and never allowing the temperature to rise above 5° C. (41° F.). On warming the solution, the greater part of the lanthanum separates out as a sulphate of the formula carrying down with it, however, a small quantity of didymium. This precipitate is filtered off hot, employing a hot-water funnel; the solu- tion is reserved for the preparation of a pure didymium compound. In order to free the precipitate of lanthanous sulphate from didymium, it is necessary to repeat the operations of dehydrating at low redness, dis- solving in ice-cold water and precipitating by warming the solution, until the solution no longer shows the spectrum of didymium. For this purpose from six to eight repetitions of this series of operations are generally necessary. In order to obtain a pure didymium salt the mother liquor from the first precipitation of the lanthanous sulphate is fractionally precipitated with oxalic acid. When the oxalic acid is very gradually added, the precipitate which is at first formed redissolves; but at length a point is reached when a permanent precipitate of crystalline, pink-colored didy- rnous oxalate separates out. This oxalate is converted into oxide, then into sulphate, which js fractionally precipitated in the same way.. After several repetitions of this treatment a product is obtained, from the spark-spectrum of which the lanthanum lines are absent. The metals were originally prepared by heating the chloride with sodium. They may, however, be obtained more readily and in a state of greater purity by the electrolysis of the fused chlorides. Properties.— Metallic cerium possesses the color and lustre of iron. It is malleable and ductile. It tarnishes in moist air. Its fusing-point lies between those of antimony and silver. When heated in air it burns 580 INORGANIC CHEMISTRY. even more brilliantly than magnesium. It slowly decomposes cold water. Dilute sulphuric and hydrochloric acids rapidly dissolve it with evolution of hydrogen; but cold concentrated sulphuric acid and concentrated nitric acid are without action upon it. COMPOUNDS OF CEEIUM. Cerous chloride, 'Cer//2C16.—Finely-divided metallic cerium inflames when thrown into chlorine, yielding a yellowish-white deliquescent mass of cerous chloride. This compound is also formed when chlorine is passed over a strongly heated mixture of cerous oxide and carbon. When a solution of the oxide in hydrochloric acid is evaporated over sulphuric acid, an aquate of the formula /oe'//2C16,50H2 is obtained in crystals. On heating, this salt is decomposed with formation of an oxy- chloride, but by the addition of amraonic chloride this decomposition may be prevented and the anhydrous chloride obtained. The bromide and iodide have also been prepared. Cerous fluoride, /Ce///2F6, is a white precipitate.—Ceric fluoride, CeF4,0H2, is a brownish powder obtained by the action of hydrofluoric acid upon ceric hydrate. When cautiously heated it first loses water and a part of its fluorine as hydrofluoric acid; on heating more strongly, a gas is given off’ which smells like chlorine and lib- erates iodine from a solution of potassic iodide—probably free fluorine (Brauner), Cerous oxide, is formed when the oxalate, or the carbonate, or ceric oxide, is heated in a current of hydrogen. It is a bluish-green powder, which absorbs oxygen from the air, and is converted into ceric oxide.— Cerous hydrate is thrown down as a bulky white precipitate when a caustic alkali is added to the solution of a cerous salt. Expo- sure to the air colors it yellow, owing to oxidation. Ceric oxide, 0eO2, is obtained by heating the oxalate or the nitrate in air or oxygen. Thus prepared it forms a colorless or faint-yellow powder, but by heating cerous chloride with borax in a wind furnace for forty-eight hours, it may be obtained in crystals belonging to the regu- lar system. On heating, it becomes darker in color, but resumes its original tint on cooling. Hydrochloric act dissolves it, yielding a yel- low solution, which when warmed evolves chlorine, and then contains cerous chloride. With concentrated sulphuric acid it also yields a yellow solution, which possesses oxidizing properties and evolves ozonized oxygen.—The hydrate has the formula Ce2OHo6. Cerous nitrate, N6012(/Ce///206)v1,120fT2, is best prepared by dissolving ceric oxide in nitric acid with the addition of alcohol, the latter substance acting as a reducing agent. It forms a crystalline mass. Ceric nitrate, NJ0BCeolv, is formed when ceric oxide is dissolved in concentrated nitric acid. It is soluble in strongly acid solutions, but excess of water decomposes it with separation of a basic salt. It forms double salts with other nitrates. Cerous sulphate, S3O6('Ce///2O(;)v 1,90 is deposited in large octahedra or prisms when a solution of ceric oxide in sulphuric acid is mixed with alcohol or sulphurous acid and allowed to evaporate spontaneously. Hot solutions deposit the anhydrous salt in minute crystals, which are COMPOUNDS OF DIDYMIUM. 581 soluble in six parts of cold and sixty parts of boiling water.—Cerous potassic sulphate, S(;012Ko6('Ce"/206)vl, separates as a white crystalline powder when an excess of potassic sulphate is added to a solution of the preceding salt. It is sparingly soluble in water, and almost insoluble in a concentrated solution of potassic sulphate. Cerous sulphate forms similar double salts with the sulphates of sodium and ammonium. Ceric sulphate, S2O4Ceoiv,7oH2, is a yellow crystalline mass. Cerous phosphate, P202(/Ce///206)vi, occurs as monazite. A portion of the cerium in this mineral is isomorphously replaced by lanthanum and didymiurn. PENTAD ELEMENTS. Section 11. DIDYMIUM, Di. Atomic weight = 146. Molecular weight unknown. Sp. gr. = 6.514 Atomicityand v. History.—Didymiurn was discovered by Mosander in 1841. Occurrence and Preparation.—See Cerium, p. 578. Properties.—ln its properties didymiurn resembles the two foregoing metals, except that it has a slightly yellow tint. COMPOUNDS OF DIDYMIUM. Didymous chloride, DiCl3, is a rose-colored crystalline mass. Its solutions deposit rose-red crystals with 6 aq. Didymous oxide, Di203, is prepared by igniting the oxalate or the hydrate. It forms a white or bluish powder, neither fusible nor vola- tile, which when strongly ignited gives a continuous spectrum inter- sected by bright bands, corresponding in position with the dark bands of the absorption spectrum of the didymiurn salts (cf, Erbia, p. 584).— Didymous hydrate, Dillo3, is obtained as a pale pink-colored precipitate by adding a caustic alkali or ammonia to the solution of a didymous salt. Didymic oxide, Di2G5, is obtained as a chocolate-colored mass by heat- ing the basic nitrate of didymiurn to dull redness in a current of oxygen. Didymous nitrate, N306 Dio///,001I2, forms large rose-red deliquescent crystals. Didymous sulphate, S:i()3Dio///2,50 i 12,I2, crystallizes in soluble, rose-red monoclinic prisms. 582 INORGANIC CHEMISTRY. TRIAD ELEMENTS.* Section IV. LANTHANUM, La. Atomic weight = 138.5. Molecular weight unknown. Sp. gr. 6.163 Atomicity f History.—Lanthanum was discovered by Mosander in 1839. Occurrence and Preparation.—See Cerium, p. 578. Properties.—Lanthanum is a malleable metal of an iron-gray color. The freshly cut surface is very lustrous, but speedily tarnishes on ex- posure to air. In its behavior towards water and acids it resembles cerium, except that it is attacked in the eold both by concentrated and by dilute nitric acid. COMPOUNDS OF LANTHANUM. Lanthanous chloride, LaCl3, is prepared like cerous chloride, which it resembles in its properties. Lanthanous oxide, La2Os, is obtained as a white powder by heating the oxalate or the nitrate. It combines with water with evolution of heat, and is converted into the hydrate LaITo3. Lanthanous nitrate, NaOgLao ///, 6 OIT2, forms colorless, deliquescent, tabular crystals. Jjcdhanous sulphate, S;!o(.Lao'/;2,9 0II2, crystallizes in six-sided prisms. The anhydrous salt is readily soluble in ice-cold water, but on gently warming the solution the above aquate separates in microscopic star- shaped crystals, which at 13° C. dissolve in less than 6 parts of water, but at 100° C. require 115 parts for their solution. (See Separation of Lanthanum, p. 579.) YTTRIUM, Y. Atomic weight 89.8. Molecular weight unknown. Atomicity f History.—The earth yttria was discovered by Gadolin in 1794. Occurrence.—This element occurs, always accompanied by erbium, in a few very rare minerals: thus as silicate in gadolinite and orthite (along with cerium, lanthanum, didyrnium, beryllium, iron, and other metals); also as tantalate, niobate, and phosphate. Recently, however, the spectroscope has shown yttrium to be a very widely diffused element (Crookes). * The remaining elements of this group have been classed as triadic; but it is quite possible that they may be only pseudo-triadic. YTTRIUM. 583 Preparation. Separation of Yttrium and Erbium.—Gadolinite is decomposed with hydrochloric acid and evaporated to expel the excess of acid. The residue is extracted with dilute hydrochloric acid, and the sohition is heated to boiling and precipitated with oxalic acid. The precipitate, which contains, in the form of oxalates, ail the yttrium and erbium, along with calcium, cerium, lanthanum, didymium, and traces of manganese and silica, is washed by decantation and heated in an open platinum dish, until the oxalic acid is totally destroyed. The mixed oxides thus obtained are dissolved in nitric acid, and a concen- trated solution of potassic sulphate is added, which precipitates the cerium, lanthanum, and didymium as double sulphates of these metals with potassium. From the filtrate the yttrium and erbium are again precipitated as oxalates, the oxalates converted by heating into oxides, the latter redissolved in nitric acid, and the solution examined with the spectroscope for didymium, the presence of which metal can be readily detected by its characteristic absorption spectrum. If didymium is present, the precipitation with potassic sulphate and the other operations must be repeated until a solution is obtained which does not give the didymium spectrum. A trace of calcium is got rid of by precipitating the yttrium and erbium as hydrates by ammonia. In order to separate the yttrium and erbium, the pure hydrates are dissolved in nitric acid, and the mixed nitrates are carefully heated in a platinum dish over a small flame until the first bubbles of nitrous anhydride begin to make their appearance. The moment this point is reached, the dish is rapidly cooled in order to prevent further decomposition, and the residue is dissolved in a quantity of warm water just sufficient to prevent the solution from becoming turbid on boiling. This solution deposits on cooling needles of a basic nitrate of erbium, which is, however, still contaminated with yttrium. Further crops of this salt, but still less pure, are obtained from the mother liquors. The purer crops are mixed, dissolved in nitric acid, again heated to incipient decomposition, and treated as above, repeating this operation until a pure erbium salt is obtained. In order to separate the yttrium in a state of purity from the erbium, with which it remains mixed in the mother liquors in the form of nitrate, the solution is evaporated to dryness, the residue heated to redness, and, after cooling, extracted with water; the solu- tion thus obtained is again evaporated to dryness, heated, and the residue extracted with water, repeating these operations until a solution is obtained which no longer gives an absorption spectrum of erbium. From this solution, which contains a basic yttric nitrate, the yttrium is precipitated by oxalic acid. The pure oxalate of yttrium is converted by ignition into the oxide. Properties.—Pure metallic yttrium and erbium have not been pre- pared. By heating the mixed chlorides of the two metals with sodium, a black powder has been obtained, which assumes a metallic lustre under the burnisher. This metallic substance burns brilliantly when heated in air. Water decomposes it slowly at ordinary temperatures, more rapidly on boiling. Acids dissolve it readily, with evolution of hydrogen. The attempt to prepare yttrium and erbium by the electrolytic de- composition of the chlorides has not proved successful. 584 INORGANIC CHEMISTRY. COMPOUNDS OF YTTRIUM. Yttrous chloride, YC13.—When the above described impure yttrium is heated in chlorine, it is converted into a non-volatile chloride. By dissolving the oxide in hydrochloric acid and evaporating, an aquate of the formula YC13,60H2 is obtained, which when heated evolves hydro- chloric acid. By heating the aquate with ammonic chloride anhydrous yttrous chloride may be obtained. The bromide and iodide closely resemble the chloride. Yttrous fluoride occurs in combination with the fluorides of cerium and calcium in the mindral yttrocerite. Yttrous oxide (Yttria), Y2Os, is obtained as a yellowish-white powder by igniting the oxalate (see p. 583). It is neither fusible nor volatile. When strongly heated it emits a pure white light, which when exam- ined by means of the spectroscope, gives a perfectly continuous spectrum, without any trace of lines or bands, a behavior which affords a means of distinguishing this oxide from that of erbium. Water neither dis- solves it nor converts it into hydrate. Mineral acids slowly dissolve it, yielding salts.— Yttrous hydrate, YHo3, is obtained as a gelatinous pre- cipitate when alkalies are added to solutions of yttrium salts. Yttrous nitrate, N306Yo// b6 O H2, is readily soluble, and forms long needles permanent in air. A basic nitrate of the formula N306Yo///,Y'H03,30H2, is obtained by heating the normal nitrate to incipient decomposition and crystallizing from a small quantity of water. Yttrous sulphate, S306Yo///2,80H2, is deposited from its solutions in well-formed crystals, which become anhydrous only at a high temperature. The anhydrous salt is much more soluble than the crystallized aquate. A saturated solution of the anhy- drous-salt prepared at 15° C. (59° F.) deposits a portion of the salt in the hydrated state on warming. ERBIUM, Er. Atomic weight = 165.9. Molecular weight unknown. Atomicity . History.—Erbium was discovered in 1843 by Mosander. Occurrence, Preparation, and Properties.—See Yttrium, p. 583. COMPOUNDS OF ERBIUM. These resemble the compounds of yttrium. Erbous oxide (Erhia), Er203, is obtained by igniting the oxalate or nitrate. It forms an amorphous mass of a yellowish color. It does not fuse at the highest temperatures, but, when strongly heated, emits a greenish light, which, when examined spectroscopically, gives a con- tinuous spectrum, intersected however by bright bands, the position of which agrees with that of the dark bands in the absorption spectrum of the solutions of erbium salts. Towards acids erbia behaves like TERBIUM 585 yttria.—JErbous hydrate, ErHo3, is precipitated by alkalies from the solutions of the salts of erbium. Erbous nitrate.—A basic nitrate of the formula N306Ero///,ErH03.30E12, obtained like the corresponding yttrium salt, is employed in the separation of erbium from yttrium. Erbous sulphate, is deposited from its solutions at 100° C. in well- formed crystals. It closely resembles in its properties yttrous sulphate. All the salts of erbium when in solution display a spectrum with characteristic absorption bands. Atomic weight 148.8 (?). TERBIUM, Tr. Very little is known concerning this element, which occurs along with yttrium and erbium in samarskite. The metal has not been isolated, and even its compounds have not been obtained free from erbium. The above atomic weight is therefore to be re- garded only as an approximation. Another metal, ytterbium {atomic weight 172.8) has lately been added by Marignae to the list of the metals of the rare earths. It occurs in crude erbia. Its oxide is white and gives no absorption spectrum. Scandium, Sc (atomic weight 44).—Very little is yet known concerning this rare element, which was discovered by Nilson in 1879. The metal has not yet been iso- lated. It occurs along with the other rare earths in gadolinite and euxenite. It is separated by means of the property which its nitrate possesses of undergoing decompo- sition at a relatively low temperature. Scandous oxide, 8C203, is a white infusible powder. Its salts closely resemble those of the other metals of this group. Samarium, Sm (atomic weight = 150), was discovered by Lecoq de Boisbaudran in samarskite. It is easily recognizable by means of its characteristic spectrum. The com- pounds of this element resemble those of didymium. Samarous chloride, SmCRbOHa, forms large tabular deliquescent crystals. Samarous oxide, SiruOa, is a white or faint- yellow powder. The solutions of its salts have a deep yellow color. Decipium, Dp (atomic weight 159 ?) was discovered by Delafontaine in the samars- kite of North Carolina. It has not yet been found possible completely to separate its compounds from those of .didymium. The solutions exhibit a characteristic absorption spectrum. General Properties and Reactions of the Compounds of the Rare Earth Metals.—The corresponding compounds of these various metals are characterized by their great similarity, so that their separation is generally a matter of difficulty. The methods for the separation of the principal members of the group—yttrium, erbium, cerium, lanthanum, and didymium—have already been given (pp. 578 and 583). 586 INORGANIC CHEMISTRY. CHAPTER XXXVIII. TETRAD ELEMENTS. Section IV. PLATINUM, Pt. Atomic weight 194.4. Molecular weight unknown. Sp. gr. 21.5. Fuses about 2000° C. (3632° F.). Atomicity " and IV. Evidence of atomicity: Platinous chloride, .... .... Pt"Cl,. Platinous oxide, .... .... PP'O. Platinic chloride, .... .... PtivCl4. Platinic oxide, . . . . Pf'O,. History.—Platinum was first recognized as a distinct metal in the eighteenth century, though it was known as a refractory metallic sub- stance a couple of centuries earlier. Occurrence.—Platinum occurs only in the native state. Native plat- inum is never pure: it contains from 50 to 80 per cent, of platinum, the remainder consisting,of iridium, palladium, rhodium, osmium, and ruthenium—which, together with platinum, constitute the so-called platinum metals—also gold, iron, and copper. This impure metal, or platinum “ore,” usually forms minute grains, although larger masses or nuggets are also found. It occurs most frequently in the sand of rivers and in alluvial deposits. The chief localities from which platinum is obtained are the Urals, Borneo, California, Brazil, and Peru. Traces have been found in the auriferous sands of the Rhine. The supply from Russia is ten times as great as that from all the other parts of the world taken together, and amounts to about 800 cwt. yearly. jExtraction.—The following is the method employed at St. Peters- burg in treating the platinum ore: The ore is first digested with dilute aqua-regia, which extracts the gold ; then with concentrated aqua-regia, as long as anything dissolves. The insoluble portion consists of grains of a native alloy of osmium and iridium. The solution contains the platinum as platinic chloride, along with small quantities of other metals. Amnionic chloride is now added to the solution, and in this way the platinum is precipitated as amnionic platinic chloride (PtCI4,2NH4CI) along with the small quantity of iridium which is present. The pre- cipitate of arnmonic platinic chloride is decomposed by heat, employing as low a temperature as possible, in order that the platinum may be ob- tained in a finely divided state. The metallic powder is formed into a cake by pressing it into a conical mould of brass, after which the cake is heated to whiteness, and welded into an ingot by hammering. In this form the platinum may be drawn into wire or rolled into plate, and otherwise worked like the most ductile metals. Instead of weld- ing the platinum, Deville and Debray fuse the metal in a lime crucible by means of the nxy-coal-gas blowpipe. Deville and Debray have also attempted to obtain platinum from its PLATINUM. 587 ores in the dry way. For this purpose the ores are fused with galena, glass, and borax. The iron present in the ore is thus converted into sulphide. Litharge is then gradually added. The litharge and galena react to yield metallic lead, the sulphur burning off as sulphurous anhy- dride (see Lead, Extraction of). The platinum and the other metals contained in the ore, with the exception of osmiridium, dissolve in the lead. The liquid portion is ladled off from the osmiridium, cupelled, and the resulting platinum fused in a lime crucible as above described. This process has been abandoned, as the platinum obtained by it is not sufficiently pure. Preparation of pure Platinum.—ln order to obtain pure platinum, commercial platinum is dissolved in aqua-regia, and from the solution, after expelling the excess of acid, the platinum and iridium are precipi- tated by caustic soda as platinic hydrate (PtHo4) and iridic hydrate (IrHo4). A little alcohol is now added, and the liquid with the pre- cipitate is boiled. Platinic hydrate is not altered by this treatment, but iridic hydrate is converted into a lower hydrate of the formula Ir2Ho6, and on reacidifying with hydrochloric acid, these hydrates go into solution, yielding the corresponding chlorides. Di-sridio hexa- chloride is not precipitated by ammonic chloride, so that on adding to the liquid an excess of this reagent the whole of the platinum is thrown down in the form of pure ammonic platinic chloride, which, after wash- ing, is converted by ignition into pure metallic platinum. Properties.—Platinum is a white metal with a tinge of gray, capable of taking a high polish. When obtained by heating to redness com- pounds of platinum—for example, ammonic platinic chloride—it forms a loosely coherent gray mass, known as spongy platinum. In the very finely divided condition in which it is deposited from the solutions of its chloride by the action of reducing agents, it forms platinum black, a black powder, quite devoid of metallic lustre. Platinum is very malle- able and ductile. Perfectly pure platinum has about the same hardness as copper, but the presence of a small quantity of iridium increases its hardness considerably. In the form of very thin wire it can be fused in the flame of a candle; * in larger masses it requires the heat of the oxyhydrogen flame for its fusion. The fusing-point has been estimated by Deville at 2000° C. (3632° F.). It does not combine directly with oxygen at any temperature, but possesses in the molten state the property of absorbing oxygen. The absorbed oxygen is expelled dur- ing the solidification of the metal, which thus exhibits the phenomenon of “spitting” (see Silver, p. 451). In like manner hydrogen passes through a diaphragm of red-hot platinum, owing to the property which the metal possesses of dissolving the gas. The red-hot metal is, how- ever, impermeable to oxygen, nitrogen, carbonic anhydride, and other gases. Cold platinum has the power of condensing various gases, espe- cially oxygen, upon its surface. This action is exhibited in a very high degree by platinum black, which, owing to its state of extremely fine subdivision and consequently increased surface, is capable of thus con- * It is possible that the fusion in this case is due to the formation of a fusible car- bide of platinum. 588 INORGANIC CHEMISTRY. densing eight hundred times its volume of oxygen. To this property is due the so-called catalytic action of platinum in bringing about the combination of various gases. Thus platinum black, when introduced into a mixture of oxygen and hydrogen, determines the explosion of the mixture. Sulphurous anhydride and oxygen, when passed over plati- num black, form sulphuric anhydride; hydrogen and iodine unite to yield hydriodic acid—the action in this and in the former case being aided by gently heating the finely divided metal, A heated spiral of platinum wire, when plunged into a mixture of ether vapor and air, or of alcohol vapor and air, continues to glow, and effects the oxidation of the organic substance. Indeed, the wire need only be warmed to 50° C. in order to glow when introduced into the vapor.—Platinum is not attacked by any single acid; but aqua-regia, or any other liquid in which chlorine is contained or is being evolved, dissolves it. It is oxi- dized by fusion with caustic alkalies or with nitre. Fused alkaline cyanides also attack it. It unites directly with silicon when heated with it, to form a brittle silicide; and with phosphorus and arsenic it yields fusible compounds. With many of the metals it forms fusible alloys. A knowledge of these facts is of importance in working with vessels of platinum ; thus phosphates ought never to be heated with carbon or with filter-paper in a platinum crucible, and the heating of compounds of easily reducible metals in such vessels is to be avoided altogether. Platinum vessels ought never to be heated over a smoky flame, as, owing to the alternate formation and oxidation of a carbide of platinum, the metal becomes blistered and porous. Contact with burning charcoal is also to be avoided, as the platinum combines with the silicon reduced from the ash.—Platinum may be sublimed at a red heat in a current of chlorine, and may thus be obtained in crystals. The sublimation of the platinum is only apparent and depends in reality upon the formation and decomposition, in rapid succession, of a chloride of platinum. Uses.—The high fusing-point of platinum, and its power of resisting chemical action, have caused it to be extensively employed in the man- ufacture of vessels for laboratory purposes. Thus platinum crucibles and evaporating basins, platinum foil and wire, are in constant requi- sition in the processes of chemical analysis. Large platinum stills are used for the concentration of sulphuric acid. The marked electro- negative character of platinum renders it capable of forming, with electropositive metals, such as zinc, voltaic combinations of high electro- motive force. Grove’s battery is a combination of this description. Platinum forms two series of compounds: platinous compounds, in which the metal is dyadic ; and platinic compounds, in which it is tetrad ic. COMPOUNDS OF PLATINUM WITH THE HALOGENS. a. Platinous Compounds. Platinous chloride, PtCl2, is obtained by heating platinic chloride to 225-230° C. (437-446° F.}. It forms a grayish-green powder, insolu- COMPOUNDS OF PLATINUM. 589 ble in water, soluble in hot hydrochloric acid, yielding a reddish-brown solution. It unites with other metallic chlorides to form double salts: thus the compounds PtCl2,’2KCI and PtCl2,2NH4CI are obtained in large red prisms by adding potassic and aramonic chloride to the solution of platinous chloride in hydrochloric acid, and evaporating the liquid.— When platinous chloride is heated in a current of carbonic oxide, the gas is absorbed with formation of the compounds PtCl2,G0, PtCl2,2C0, and PtCl2,3C0. It also unites directly with ethylene ("CgHJ and other unsaturated hydrocarbons. Platinous bromide, Ft Bra, is prepared by heating hydric platinic bromide to 200° C. It forms a brown mass. Platinous iodide, Pth, is obtained as a black powder by warming platinous chloride with a concentrated solution of potassic iodide. b. Platinic Compounds. Platinic chloride, PtCl4, is prepared by dissolving platinum in aqua-regia, destroying the nitric acid by repeated evaporation with hydrochloric acid, and heating to expel the excess of hydrochloric acid. It crystallizes from water in large red non-deliquescent crystals of the formula PtCl4,50H2.—Platinic chloride forms numerous double salts with other chlorides: thus with hydrochloric acid it forms the com- pound PtCl4,2HC1, which is deposited in brownish-red deliquescent prisms with 6 aq. from the solution of platinic chloride in hydrochloric acid. Potassic platinic chloride, PtCl4,2KC1, and ammonia platinic chloride, PtC!4,2NH4CI, are obtained as yellow crystalline precipitates, consisting of microscopic octahedra, when platinic chloride is added to solutions of potassic and amnionic chloride. These precipitates are almost insoluble in water and quite insoluble in alcohol. Sodio platinic chloride crystallizes in reddish-yellow prisms of the formula PtCI4,2i5raC1,601T2, readily soluble in water and in alcohol. The dif- ference in the solubility of these compounds is turned to account in the separation of the alkali metals. Platinic bromide, Ptßr4, has not been prepared, but hydric 'platinic bromide, FtRr4.2H8r,90H2, is known. Platinic iodide, Ptt4, separates as a black powder when potassic iodide is added to a solution of platinic chloride and the liquid warmed. COMPOUNDS OF PLATINUM WITH OXYGEN AND HYDROXYL. Platinous oxide, , . PtO. Pt=o. Platinous hydrate, . PtHo2. H—O—Pt—O—H. Platinic oxide, . . Pt02. o=Pt=o. H—O O—H \/ Platinic hydrate, . , PtHo4. Pt /\ H—O O—II 590 INORGANIC CHEMISTRY. Platinous oxide, PtO.—This compound is obtained as a grayish- black powder by gently heating the corresponding hydrate. Platinous hydrate, PtHo2, is a bulky black powder, obtained by digesting platinous chloride with warm caustic potash. Boiling caustic potash decomposes it with separation of metallic platinum and forma- tion of platinic oxide. It acts as a weak base and yields with the hydracids the corresponding haloid salts; but the oxy-acids, with the exception of sulphurous acid, decompose it. Platinic oxide, Pt()2, is a black powder obtained by gently heating platinic hydrate. Platinic hydrate, PtHcq.—A solution of platinic chloride is precipi- tated by boiling with caustic potash, and the precipitate is treated with acetic acid to remove the potash, when a white compound of the formula PtHo4,20H2 remains. This, on drying at 100° C., parts with 2 aq. and assumes an amber-brown color. Platinous hydrate acts both as a weak base and as a weak acid. The salts which it forms with bases are known as pfatinates. Baric platinate is a yellow powder of the formula PtHo28a0",30H2. OXY-SALTS OF PLATINUM. Very few of the simple oxy-salts of platinum have been prepared, but various double salts are known. Platinous sulphite is obtained as a gummy mass of unknown composition by evapo- rating the solution of platinous hydrate in sulphurous acid.—Potassic platinous sul- phite, SOPto//,3SOKo2,20H2 crystallizes in readily soluble needles. Sodic platinous sulphite, SOPto//,350Na0,70H2, is a sparingly soluble crystalline precipitate. Platinonitriies.—Platinum forms a series of remarkable compounds with the nitrites of other metals. These compounds do not behave like ordinary double salts: the platinum cannot be detected in their solutions by the ordinary reagents. They may be regarded as salts of platinonitrous acid, HaPtfJS'Ch q.— Potassicplatinonitrite, K2pt(No2)4, is deposited in small lustrous prismatic crystals when solutions of potassic nitrite and po- tassic platinous chloride are warmed together. Its solutions are not precipitated either by alkalies or by sulphuretted hydrogen.—Ammonieplatinonitrite, (N11412Pt(N02)4.20112, crystallizes in prisms. It decomposes with sudden incandescence when heated. COMPOUNDS OF PLATINUM WITH SULPHUR. Platinous sulphide, PtS//.—This compound may be obtained as a black amorphous powder by passing sulphuretted hydrogen over moistened platinous chloride, or in a crystalline form by fusing platinous chloride with sodic carbonate and sulphur, and lixiviating the mass with water. Platinic sulphide, PtS//2.—Sulphuretted hydrogen precipitates, from solutions of platinic salts, black platinic sulphide, and this compound then unites with a further quantity of the gas to form light-brown hydric platinic sulphide, an unstable com- pound which parts with sulphuretted hydrogen when exposed to the air.—By fusing a mixture of spongy platinum, potassic carbonate, and sulphur, and extracting the n f PtPts//Ks mass with water, an insoluble dipotassic diplatinous sulphodiplati ate, -< ptpts//j£s> *9 obtained in thin lead-gray six-sided tablets. When heated in a current of gaseous hydrochloric acid, this compound evolves sulphurretted hydrogen, and is converted into potassic chloride and a platinic sulphide of the formula (possibly, how- f PtPts" \ , ever, Pt4S//6= | )’ w^ remains as a steel-gray powder on extracting the mass with water. COMPOUNDS OF PLATINUM. 591 AMMONIUM COMPOUNDS OF PLATINUM {PL A TINA MINES). Platinum forms a remarkable class of ammonium bases, the salts of which may be empirically formulated as double compounds of platinum salts with two or more molecules of ammonia. In this respect these compounds resemble the cobaltamines {q.v.). They have been divided into no fewer than twelve distinct classes. The members of one class are sometimes isomeric with those of another class. A complete account of these compounds would go beyond the scope of the present work. The following will serve as examples: Platosotetrammonic chloride (chloride of “Reiset’s first base”) is obtained in colorless ( NH2(N*H4)CI prisms of the PC' ,OH2, when platinous chloride is dissolved in an ( NH2(NTH4)CI excess of boiling aqueous ammonia and the solution evaporated. It forms with plati- nous chloride an insoluble double salt, crystallizing in dark green needles of the t N1T2(NtH4)CI formula < Pt// ,PtCl2, also known as the green salt of Magnus. This compound, _ ( NH2(N;H4)CI which is interesting as the first discovered of the platinum ammonium compounds, may be obtained direct by supersaturating with ammonia a hot solution of platinous ( NH2(NtH4)Ho chloride in hydrochloric acid.—Platosotetrammonic hydrate, -I Pt// , is pre- ( NH2(NtH41Ho pared by precipitating a solution of the sulphate with baric hydrate and evaporating the filtrate. It crystallizes in deliquescent needles. It acts as a caustic, absorbs carbonic anhydride from the air, and precipitates the metals as hydrates from the solutions of iheir salts. ( NH3CI Platosodiammonic chloride (chloride of “Reiset’s second base”),-\ Vi" .—This cora- ( N FTgCI pound, which is isomeric with the green salt of Magnus, is formed when platosote- trammonic chloride is heated to between 220° and 270° C. (430-518° F.). It forms microscopic, yellow rhombohedra. It is sparingly soluble in water, and is formed as a precipitate when hydrochloric acid is added to the solutions of other salts of this base. Both the hydrate and the oxide are known. The latter compound, which has f ITH3 1 the formula t Pt// O, is obtained by heating platosotetrammonic hydrate to 110° C. Inh3- ( NH2CI Platinodiammonic chloride (chloride of ‘‘Oerhardt’s base”),\ PtCl2 ,is formed bv the ( NH2CI direct union of platosodiammonic chloride with chlorine, when the gas is passed through water in which this salt is suspended. It crystallizes in minute yellow octahedra. f NH2(NtH4)CI Platinotetrammonic chloride (“Oros’ chloride”) Ft'Jl2 .—This compound is k NH2(NMT4)CI formed in a similar manner by the union of platosotetrammonic chloride with chlo- rine, or by treating platinodiammonic chloride with ammonia. It crystallizes in yellow octahedra of the regular system. General Properties and Reactions of the Compounds of Platinum.—a .Platinous Compounds.—These are of subordinate inte- rest. The platinous salts are of a red, brown, or green color. b. Platinio Compounds.—The platinic salts have a yellow color. With caustic soda they give a yellow precipitate of platinic hydrate, soluble 592 INORGANIC CHEMISTRY. in an excess of the alkali. Sulphuretted hydrogen precipitates, slowly in the cold, more rapidly on heating, platinic sulphide, which is soluble in a large excess of amraonic sulphide. Potassic chloride and ammonia chloride produce yellow crystalline precipitates of potassic platinic chloride and amnionic platinic chloride. Stannous chloride in acid solu- tions produces a dark coloration, owing to the reduction of the platinic salt to the platinousstage, but no separation of metallic platinum occurs. Ferrous sulphate precipitates metallic platinum, but only after protracted boiling. Oxalic acid does not reduce the salts of platinum (separation from gold); but by boiling with soluble formates in alkaline solution, metallic platinum is precipitated. All platinum compounds, when ignited with access of air, are converted into metallic platinum. PALLADIUM, Pd. Atomic weight = 105.7. Molecular weight unknown. Sp. gr. 11.4. Atomicity " andlv. Evidence of atomicity : Fallacious chloride, . . . . . . . PcF'Cl, Fallacious oxide, .... . . . . Pd"0. Falladic chloride, . . . . . PdivCl4. Palladio oxide, History.—Palladium was discovered by Wollaston in 1803. Occurrence.—Granules of this metal, sometimes in the form of octa- hedra, occur in the platinum ore of Brazil. Alloyed with platinum and other metals, it occurs in all ores of platinum. Preparation.—One method of separation of palladium from the other metals of the platinum-group with which it occurs, depends upon the fact that palladium is precipitated as insoluble palladous iodide by the careful addition of potassic iodide to the solution of palladous chlo- ride. The other metals remain in solution. An excess of the precipitant is to be avoided, as it dissolves the palladous iodide. The iodide loses its iodine when strongly heated, and is converted into spongy palladium. —ln order to extract the palladium from platinum ore, the solution which is obtained after dissolving the ore in aqua-regia and removing the platinum by precipitation with amrnonic chloride, is treated with mercuric cyanide. In this way a precipitate of palladous cyanide is produced, which by ignition may be converted into the metal. Properties.—Palladium is a silver-white lustrous metal. It sometimes occurs crystallized, either in octahedra or in small hexagonal plates. Palladium is the most fusible of the platinum metals and can be welded at a red heat more readily than platinum. When heated to low redness it undergoes superficial oxidation, and assumes a blue color, but at a higher temperature regains its lustre. It is soluble in hot nitric acid and in hot concentrated sulphuric acid. Hydrochloric acid dissolves spongy palladium in presence of air. It is not altered by exposure to air or to sulphuretted hydrogen.—Spongy palladium, like spongy plat- inum, is capable of effecting the combination of oxygen and hydrogen when introduced into a mixture of these gases. If the two gases are present in the proportions necessary to form water, the palladium COMPOUNDS OF PALLADIUM. 593 becomes red-hot, causing explosion; but if a considerable excess of oxygen is present or if air be substituted for oxygen, the combination takes place slowly at ordinary temperatures without explosion. In the case of a mixture of hydrogen, marsh-gas and air, it is possible to effect the slow combustion of the hydrogen, leaving the marsh-gas untouched, and in this way the hydrogen present in a mixture of combustible gases may be determined.—lf a piece of palladium foil be heated in the flame of a spirit lamp, or in a coal-gas flame, the foil becomes covered with cauliflower-like excrescences of soot, and when these are burnt they leave a skeleton of filaments of metallic palladium, whilst the foil is found to have become porous. In like manner, when spongy palladium is heated in a current of ethylene, the gas is decomposed with separa- tion of carbon at a temperature at which ethylene alone is perfectly stable. These phenomena probably depend upon the affinity of palla- dium for hydrogen, palladium hydride (q.v.) being successively formed and decomposed. In the formation of this compound carbon is liberated from the gases present in the flame; in its decomposition the palladium disintegrates. Uses.—Palladium is used for the graduated scales of physical instru- ments and also for coating silver goods. COMPOUND OF PALLADIUM WITH HYDROGEN. Palladium hydride, Pd4H2.—This compound is formed by the direct union of its elements when palladium is heated in a current of hydrogen, or when this metal is employed as negative electrode in the electrolysis of dilute sulphuric acid.—Palladium hydride is a lustrous metallic mass with a specific gravity of 11.06. It conducts electricity. It parts with its hydrogen only very gradually at ordinary temperatures, but rapidly on heating. On exposure to the air in a finely divided state it becomes red hot, owing to the absorption of oxygen and oxidation of the hydrogen to water. It acts as a reducing agent; thus it precipitates metallic mercury from solutions of the salts of that metal. COMPOUNDS OF PALLADIUM WITH THE HALOGENS. Palladous chloride, PdCl2.—When a solution of palladium in aqua- regia is evaporated to dryness, the palladic chloride which is at first formed is decomposed and converted into palladous chloride, which remains as a brown deliquescent mass. This compound may also be obtained as a red crystalline sublimate by heating palladous sulphide (PdS") in a current of dry chlorine. In this form it dissolves only slowly in water.—Like the corresponding platinum compound it forms numerous double chlorides. Potassic palladous chloride has the*formula PdCl2,2KC1. a. Palladous Compounds. Palladous bromide is not known in the pure state. Palladous iodide, Pdl2.—This compound is precipitated as a black powder when potassic iodide is added to solutions of palladous chloride 594 INORGANIC CHEMISTRY. or nitrate. It is soluble in an excess of potassic iodide. lodine may be estimated as palladous iodide in presence of chlorine and bromine. h. Palladio Compounds. Of these only the chloride is known, and this has been obtained only in solution. It forms, however, well-characterized double salts, corre- sponding to those of platinum: thus potassic palladia chloride, PdCI4,- 2110!, which crystallizes in brownish-red octahedra; and amnionic pal- lodic chloride, PdCl4,2NH4CI, which forms a sparingly soluble red crystalline powder. COMPOUNDS OF PALLADIUM WLTH OXYGEN. f Pd Hypopalladous oxide, . 1 Pd°* I /a Palladoas oxide, . . . PdO. Pd=o. Palladio oxide, . . . Pd02. o=Pd=o. Hypopalladous oxide, is obtained as a black powder by heating palladous hydrate to low redness as long as oxygen is evolved. Acids decompose it with separation of metallic palladium and formation of palladous salts. When heated in a current of hydrogen it is reduced with sudden incandescence. Palladous oxide, PdO, is prepared by careful ignition of the nitrate. It forms a black powder which dissolves with difficulty in acids. When brought into hydrogen at ordinary temperatures it is instantaneously reduced with incandescence.—Alkaline carbonates precipitate from solutions of palladous salts a dark-brown hydrate, which dissolves readily in acids. Palladio oxide, Pd02, is a black powder obtained by boiling potassic palladic chloride with caustic potash and washing the precipitate with hot water. PALLADOUS OXY-SALTS. Palladous nitrate, N2OJPdo", is prepared by dissolving the metal or the oxide in nitric acid. On evaporation the solution deposits long brown deliquescent prisms. Palladous sulphate, S02Pdo//,20H2, is obtained by dissolving the metal in sulphuric acid, with the addition of nitric acid, and evaporating. It forms brown soluble crystals, which are decomposed by excess of water with separation of a basic salt. A series of ammonium compounds of palladium, corresponding with those of plati- num, is known. COMPOUNDS OF PALLADIUM WITH SULPHUR. These correspond with the oxides. Hypopalladous sulphide, /Pd/2S//, is formed when either palladous sulphide or pal- ladic sulphide is heated in a current of carbonic anhydride. It is most readily ob- tained by fusing together at a red heat a mixture of palladous sulphide, potassic car- bonate, sulphur, and arnmonic chloride. On dissolving the mass in water, hypopallad- ous sulphide remains as a brittle, green, metallic regulus. It is only slowly attacked by nitric acid. IRIDIUM. 595 Fallacious sulphide, PdS//, is obtained as a grayish-white metallic mass by heating the metal in the vapor of sulphur, when combination occurs with incandescence. The same compound is precipitated as a black amorphous powder when sulphuretted hy- drogen is passed into solutions of palladons salts. Palladio sulphide, PdS//2.—When palladons sulphide is fused with sulphur and sodic carbonate, sodic sulphnpalladate, PdS// Nas2, is formed. On decomposing this com- pound with hydrochloric acid, palladia sulphide is obtained as a dark-brown powder. It dissolves readily in aqua-regia. General Properties and Reactions of the Compounds of Palladium.—The palladons salts are for the most part soluble, yield- ing solutions which, when concentrated, are brown or reddish-brown, when dilute, yellow. Both sulphuretted hydrogen in acid solution and ammonia sulphide precipitate black palladous sulphide, insoluble in excess of amnionic sulphide, but soluble in boiling hydrochloric acid. Caustic alkalies precipitate brown basic salts of palladium, soluble in an excess of the alkali on heating. Ammonia gives a flesh-colored precipi- tate of a palladamraoniura compound, soluble in excess of ammonia. Potassic iodide precipitates black palladous iodide. Ferrous sulphate precipitates metallic palladium, the action being facilitated by heat. All palladium compounds yield on ignition in air metallic palladium. IRIDIUM, Ir. Atomic weight 192.5. Molecular weight unknown. Sp. gr. 22.38. Atomicity " and lv, also a pseudo-triad. Evidence of atomicity: Iridous sulphide, .... Ir"S" Di-iridic hexachloride, .... /Ir"/2C1B, Di-irldic trioxide, .... 'Ir"'203. Iridic chloride, Iridic oxide, .... IriT0„. History.—lridium was discovered in 1804 by Smithson Tennant. Occurrence.—lridium occurs in most ores of platinum in the form of granules of the alloys platiniridiura and osrairidium. Extraction.—For the preparation of iridium the residue which re- mains when the platinum ore is treated with aqua-regia is employed. This residue, which consists chiefly of iridium and osmium, but con- tains small quantities of all the other platinum metals, is fused with from 20 to 30 times its weight of zinc. On dissolving the zinc in hy- drochloric acid, the platinum metals remain as a fine powder. This powder is mixed with from 3 to 4 parts of anhydrous baric chloride, and the mixture is heated to low redness in a current of dry chlorine. On dissolving in water, ruthenium remains behind, whilst the other platinum metals dissolve as double chlorides of barium with the plati- num metal. Sulphuric acid is then added so as exactly to precipitate the barium. The liquid, which now contains the platinum metals as chlorides, is heated in an atmosphere of hydrogen in a flask on a water- bath. In this way the metals are reduced from their aqueous solution. During the whole of this operation air must be carefully excluded, as the finely divided metals would bring about the explosive combination of the hydrogen with the oxygen of the air. Platinum and palladium 596 INORGANIC CHEMISTRY. are first reduced, then rhodium. Before the iridium is precipitated if undergoes reduction to di-iridic hexachloride, the presence to which is manifested by an olive-green coloration of the liquid. At this point the operation is interrupted, and after filtering off the reduced metals, the iridium is precipitated from the filtrate by first oxidizing it with nitric acid to iridic chloride, IrCl4, and then adding a solution of potassic chloride, with which it forms a black, almost insoluble crystal- line precipitate of potassic iridic chloride, IrCl4,2KC1. This on igni- tion yields spongy iridium. A trace of ruthenium may be removed by fusing the spongy metal with nitre. On lixiviating the fused mass with water the ruthenium dissolves as potassic ruthenate, leaving the iridium. Properties.—lridium is a white metal, which when polished has a lustre resembling that of steel. It is harder than platinum, and much more brittle. It is also more refractory than platinum, but may be fused in the oxyhydrogen flame. Very finely divided iridium (iridium black) dissolves in aqua-regia and oxidizes when heated in air. Com- pact iridium is not attacked under any of these conditions, but may be oxidized by fusion with potassic hydrate to which nitre or potassic chlo- rate has been added. Iridium black is obtained as an impalpable pow- der by exposing an alcoholic solution of di-iridic sulphate to sunlight. It is more energetic in its catalytic action than platinum black. A small quantity brought upon paper moistened with alcohol causes ignition. Uses.—An alloy of 1 part of iridium with 9 parts of platinum is ex- tremely hard and elastic, capable of taking a high polish, and unal- terable in air. It has been employed in the preparation of standard measures of length. Gold pens are sometimes tipped with an alloy of iridium and osmium. COMPOUNDS OF IRIDIUM WITH THE HALOGENS. Di-iridic hexachloride, Tr'^Clg.—This compound is formed when the metal is heated in chlorine. It is most readily obtained by heating one of its alkaline double chlorides, such as potassic di-iridic chloride, with concentrated sulphuric acid and pouring the cooled liquid into water, when the chloride separates as a pale olive-green pre- cipitate, insoluble in water and in acids. It may be obtained in a solu- ble form by treating a solution of iridic chloride with sulphurous anhy- dride until the solution has become green.—The alkaline double chlo- rides are formed when the corresponding iridic double chlorides are reduced in aqueous solution with sulphurous anhydride or sulphuretted hydrogen. Potassic di-iridic chloride, Tr/;yil6,6KC1,601I2, sodic di- iridic chloride, /1r,"2C16,6NaC1,240H2, and ammonia di-iridic chloride, 3OH2, all form olive-green crystals, soluble in water, insoluble in alcohol. a. Di-i7'idiG Compounds. Di-iridic hexabromide, /Ir///28r6,80H2, is deposited in light olive-green six-sided crystals when a solution of iridic hydrate, Ir Ho4, in hydrobromic acid is evaporated. The iridic bromide does not appear to be capable of existing ; the solution evolves bromine and contains the lower bromide. Di-iridic hexabromide forms double bro- mides corresponding with the double chlorides. COMPOUNDS OF IRIDIUM. 597 h. Iridic Compounds. Iridic chloride, IrClt, is obtained as a black mass by dissolving irid- ium black, di-iridous trioxide, or di-iridic hexachloride in aqua-regia, and evaporating the solution at a temperature below 40° C. (104° F.). On heating to a higher temperature chlorine is evolved, and the solution contains the lower chloride.—It forms with the chlorides of the alkalies double chlorides, isoraorphous with those of platinum. Potassic iridic chloride, IrCl4,2KC1, and amnionic iridic chloride, IrCI4,2NH4CI, crys- tallize in minute dark-red octahedra, sparingly soluble in cold water. Sodic iridic chloride, Ir(Jl4,2NaCI, is readily soluble in water, and forms black tabular crystals or prisms. Iridic bromide, Irßr4, is not known; but numerous double bromides corresponding with the double chlorides have been prepared. Iridic iodide, Irl4, is obtained as a black powder by the action of potassic iodide upon the solution of the chloride in hydrochloric acid. COMPOUNDS OF IRIDIUM WITH OXYGEN. o /\ Di-iridic trioxide, Ir—lr. il II o o Iridic oxide, . . Ir02. o=lr=o. Di-iridic trioxide, 'lr'"203.—This compound is formed when finely divided iridium is heated in air. At a higher temperature it is again decomposed into oxygen and metal. It is most readily prepared by heating a mixture of potassic iridic chloride and sodic carbonate to low redness : 2IrCl4,(KCI)2 + 4CONao2 = 'lr"f203 + BNaCI + Potassic iridic Sodic Di-iridic Sodic chloride. carbonate. trioxide. chloride. 4KCI + 4G02 + O Potassic Carbonic chloride. anhydride. On extracting the mass with water the oxide remains behind as a black powder. Hydrogen, even at ordinary temperatures, reduces it to the metallic state.—When a solution of potassic di-iridic chloride is pre- cipitated by a small quantity of caustic potash with exclusion of air, yellowish-green di-iridic hexahydrate, is obtained. It is solu- ble in excess of alkali, and oxidizes on exposure to air. Iridic oxide, Ir02.—When moist di-iridic hexahydrate undergoes spontaneous oxidation by exposure to air, it is converted into iridic hy- drate, IrHo4. The same compound is obtained by precipitating iridic chloride with caustic alkali. It forms an indigo-blue powder, which is not atacked by dilute acids with the exception of hydrochloric. When carefully heated in a current of carbonic anhydride it is converted into iridic oxide, which is thus obtained as a black powder insoluble in acids. 598 INORGANIC CHEMISTRY. OXY-SALTS OF lEIDIUM. These are comparatively unimportant. Salts of the unknown iridous oxide, IrO, have been prepared ; thus a sodic iridous sulphite of the formula S4G4Nao6lro//,100H2 is known. An oxy-salt corresponding to di-iridic trioxide is di-iridic trisulphite, B303(/Ir///206)V 1,66 which is obtained as a crystalline powder by dissolving the hexyhydrate in sulphurous acid and evaporating. No iridic oxy-salts are known. Ammonium compounds of iridium corresponding with those of platinum have been prepared. COMPOUNDS OF IRIDIUM WITH SULPHUR. Iridous sulphide, IrS//, is obtained as a lustrous metallic mass when the metal is heated in the vapor of sulphur. Di-iridic trisulphide, /Ir///2S//3, is obtained as a brown precipitate when sulphuretted hydrogen is passed into the solution of a di-iridic salt. Iridic sulphide, IrS//2.—This compound is prepared by heating the finely divided metal with sodic carbonate and sulphur, extracting the mass with water. The iridic sulphide remains as a black powder. General Properties and Reactions of the Compounds of Iridium.—A not too dilute solution of an iridic salt yields with am- nionic chloride a dark-red crystalline precipitate of ammonic iridic chloride. From the solution of an iridic salt sulphuretted hydrogen precipitates brown di-iridic trisulphide (/Ir///2S//3) with separation of sulphur. Ferrous sulphate decolorizes the solution of an iridic salt; zinc precipitates black spongy iridium. RHODIUM, Rh. Atomic weight 104. Molecular weight unknown. Sp.gr. 12.1. Atom- icity " aridlv, also a pseudo-triad. Evidence of atomicity: Rhodens oxide, . . . Rh"G. Dirhodous hexaehloride,.... . . . 'Eh'"Cl, Dirhodous trioxide, . . . /Eh///2G3. Rhodic hydrate, . . . RhiTHo4. Rhodic oxide, . . . EhlvO„. History.—Rhodium was discovered by Wollaston in 1804, and afterwards investigated more thoroughly by Berzelius and Claus. Occurrence.—The metal occurs in small quantity in platinum ore. Extraction.—The only source of rhodium is the platinum residue already referred to. The mixture of platinum, palladium, and rhodium precipitated by hydrogen in the process of separating the platinum metals is redissolved in aqua-regia, and the platinum is precipitated by potassic chloride. After expelling the excess of acid, the rhodium may be precipitated as sodic dirhodous sulphite, S606Nao6('Rh/"206)vi, by boil- ing the dilute solution with hydric sodic sulphite. The metal may be precipitated by reducing agents from the solutions of its salts and fused into a coherent mass in the oxyhydrogen furnace. Properties,—Rhodium is a malleable metal, resembling aluminium in color and lustre. Its fusing-point lies between that of platinum and that of iridium. When heated in air it undergoes superficial oxida- tion. Pure rhodium is insoluble in all acids and in aqua-regia. If, however, it is alloyed with an excess of platinum, or with zinc, lead, and other oxidizable metals, aqua-regia dissolves it. 599 COMPOUNDS OF RHODIUM. COMPOUND OF RHODIUM WITH CHLORINE. Dirhodic hexachloride, 'Rh'^Clg.—This is the only halogen compound of rhodium which is known with certainty. The anhydrous chloride is formed when the finely divided metal is heated in chlorine. It is an insoluble rose-red powder. By dissolving dirhodic hexahydrate in hydrochloric acid and evaporating the solution, a dark-red hydrated chloride is obtained, which on heating is converted into the anhydrous chloride. Dirhodic hexachloride forms double salts with the alkaline chlorides. COMPOUNDS OF RHODIUM WITH OXYGEN. Rhodous oxide, . RhO. Rh=o. O /\ Dirhodic trioxide, /Rli///203. Eh—Eh. II II o o Rhodic oxide, . Rh02. o=Rh=o. Rhodous oxide, RhO.—This compound is formed with incandescence when the hexahydrate is heated. It is a dark-gray powder, insoluble in acids. Dirhodic trioxide, /Rh///2()3, is obtained as a gray spongy lustrous mass by heating the nitrate. It does not dissolve in acids.—Dirhodic hexa- hydrate is prepared by the action of hot caustic potash upon sodic di- rhodic chloride, lt is a brownish-black gela- tinous precipitate, difficultly soluble in acids. By the action of caustic soda upon the double chloride in the cold, yellow crystals of the hy- drate /Rh/"2H06,20H2 are obtained. These dissolve readily in acids. Rhodic oxide, Rh02, is obtained by repeatedly fusing finely divided rhodium with caustic potash and nitre. It is a brown powder, insoluble in acids. OXY-SALTS OF RHODIUM. Dirhodic nitrate, Ne012(/Kh///206)vl, is nncrystallizable. Dirhodic sulphate, S306(/Bli///206)vi,120H2, is obtained as a yellow soluble crystal- line mass by evaporating the solution of the yellow hydrate in sulphuric acid. Dirhodic sulphite, S303(/E.h///2<)6)v1,60H2, remains as a yellow, difficultly crystalliza- able mass when the solution of the yellow hydrate in sulphurous acid is evaporated. Ammonium compounds of rhodium have been prepared. These are derived from dirhodic t rioxide. Rhodous sulphide, Rh3//.—This compound is formed as a fused metallic mass when rhodium is heated in the vapor of sulphur. COMPOUND OF RHODIUM WITH SULPHUR. General Properties and Reactions of the Compounds of Rhodium.—The solutions of the dirhodic salts are sometimes rose- colored, sometimes yellow. Caustic alkalies give a yellow precipitate, which, on heating the liquid with the precipitate, becomes brownish- 600 INORGANIC CHEMISTRY. black, and then consists of dirhodic hexahydrate. Sulphuretted hydro- gen and ammonia sulphide give, after protracted action aided by heat, a brown precipitate, probably a dirhodic trisulphide (/Rh///2S//3). Potassic iodide precipitates sparingly soluble yellow dirhodic hexiodide. Zinc precipitates black metallic rhodium. OCTAD ELEMENTS. OSMIUM, Os. Atomic weight =198.6 ? Molecular weight unknown. Sp. gr. 22.477. Atomicity", i7,,vi, and Tiii, also a pseudo-triad. Evidence of atomi- city : Osmous oxide, .... 0s"O. Diosmic trioxide, .... '0s'"2O3. Osmic chloride, Potassic osmate, Osmic peroxide, History.—Osmium was discovered in 1804, by Smithson Tennant. Occurrence—lt occurs alloyed with iridium, in the ores of platinum. This alloy, known as osmiridium, remains behind when the ore is treated with aqua-regia. Extraction.—lf in the preparation of iridium (p. 595) the mixture of the finely divided platinum metals with baric chloride be heated in a current of moist chlorine, the greater part of the osmium is volatilized as osmic peroxide, and may be condensed in a cooled receiver. The rest of the osmium may be recovered if the solution containing the chlorides of the platinum metals, which remains after the precipitation of the barium in the above operation (p. 695), be mixed with excess of nitric acid and distilled. The aqueous distillate contains the osmium as per- oxide. On adding to the solution of the peroxide ammonia and ara- rnonic sulphide, the osmium is precipitated as osmic persulphide, OsS"*. This is mixed with sodic chloride and heated in a slow current of chlorine. On extracting with water, a solution of sodic osmic chloride, OsCl4,2NaCI, is obtained, from which on the addition of amnionic chloride the osmium is precipitated as amnionic osmic chloride, When this is ignited in a covered crucible, metallic osmium is obtained as a spongy mass. By fusing spongy osmium with tin, and dissolving the tin with hydrochloric acid, osmium is obtained in crystals. Properties.—Osmium is not fusible at the highest temperatures, though it is volatile when heated to the fusing-point of iridium. Heated in air it burns, forming osmic peroxide, and if a quantity of finely divided osmium be ignited at one point, the ignition is propagated throughout the mass. Aqua-regia also oxidizes the finely divided metal to peroxide. Crystallized osmium forms cubes. In this con- dition it has asp. gr. of 22.477, and is therefore the heaviest substance known. COMPOUNDS OP OSMIUM. 601 Diosmic hexachloride, 'Os"'2C\6, is known only in the form of its double chloride. Potassic diosmio chloride, /os///2C16,6KC1,60H2, forms dark-red crystals. Osmic chloride, OsCl4, is obtained as a red sublimate when the metal is heated in dry chlorine. It dissolves in water yielding a yellow so- lution, which gradually deposits lower oxides of osmium, and becomes colorless. The solution then contains osmic peroxide and hydrochloric acid. Osmic chloride forms double salts. COMPOUNDS OF OSMIUM WITH CHLORINE. COMPOUNDS OF OSMIUM WITH OXYGEN Osmons oxide, . OsO. Os=o. O /\ Diosmic trioxide, 'os'"2O3. Os—Os. II II O O Osmic oxide, . 0s02. o=os=o. O II Osmic peroxide,. 0s04. O—os=o. O Osmous oxide, OsO, is obtained as a grayish-black powder, insoluble in acids, by heating a mixture of osmous sulphite, BOOso", with sodic carbonate, in a current of carbonic anhydride. Diosmic trioxide, 'os'"203, is prepared by heating potassic diosmic chloride with sodic carbonate. It is a black powder, insoluble in acids. Osmic oxide, 0s02, is obtained in a similar way from potassic osmic chloride, OsC14,2KC1. Thus prepared it forms a grayish-black pow- der; but by heating osmic hydrate in a current of carbonic anhydride, it is obtained in copper-colored masses, possessing a metallic lustre.— Osmic hydrate, OsHo4, is formed as a black precipitate when reducing agents, such as alcohol, are added to the aqueous solution of osmic per- oxide. Osmic peroxide (Osmic anhydride, Osmic acid), 0s04. Molecular volume I I I.—This remarkable compound is formed when the finely divided metal, or any of the lower oxides of osmium, is heated in air or oxygen, or dissolved either in nitric acid or in aqua-regia. If the finely divided metal has been previously ignited with exclusion of air, these solvents are without action upon it, Osmic peroxide forms long colorless prisms or needles, with a powerful and irritating odor. They sublime even at ordinary temperatures, and when gently heated fuse to a colorless liquid, which boils without decomposition at 100° C. Osmic peroxide dissolves in water, yielding a neutral solution with a powerful odor and a burning taste. Alcohol and ether precipitate from the solu- tion osmic hydrate. Sulphurous anhydride colors the solution in turn 602 INORGANIC CHEMISTRY. yellow, brown, green, and finally blue, at which point the liquid con- tains osrnous sulphite. The vapor of osmic peroxide, even when largely diluted with air, attacks the lungs, producing dangerous inflammation of the mucous membrane. It also acts violently upon the eyes, and may even cause blindness, owing to the deposition of a film of metallic osmium upon the eye. Brought in contact with the skin, osmic perox- ide produces a painful eruption, which is very difficult to heal. OXY-SALTS OF OSMIUM. These are few in number, and unimportant. Osrnous sulphite, SOOso//, is obtained by passing sulphurous anhydride into a solu- tion of osmic peroxide until the solution assumes a blue color, and then adding sodic sulphate. The osmium salt, which is sparingly soluble in a solution of sodium sul- phate, is deposited as a dark-blue precipitate.—Hydric potassic osrnous sulphite, is obtained as a rose-red precipitate by heating a solution of potassic diosmic chloride (p. 601) with potassic sulphite. The Osmates. Neither osmic acid, 030,Ho2, nor its anhydride, 0s03, is known; but some of the salts of osmic acid have been prepared. Potassic osmate, Os02Ko2,20H, is obtained by adding alcohol or po- tassic nitrite to a sufficiently concentrated solution of the peroxide in potassic hydrate. The peroxide is reduced and unites with the alkali to form potassic osmate, which gradually separates as a dark-red crys- talline powder. Baric osmate, Os()2Bao//, forms black lustrous prismatic crystals. COMPOUNDS OF OSMIUM WITH SULPHUR. The sulphides of osmium have been but little studied. Osmium combines with sul- phur when heated in its vapor, and sulphuretted hydrogen precipitates osmium as sulphide from its solutions. From solutions containing osmium in its lower stages of oxidation a yellow sulphide is precipitated ; whilst solutions of the peroxide give a brown precipitate of osmic persulphide, OsS^. General Properties and Reactions of the Compounds op Osmium.—Osmium and its compounds are best characterized by the readiness with which they yield the volatile peroxide, recognizable by its powerful odor. All osmium compounds when boiled with nitric acid give off'vapors of the peroxide. RUTHENIUM, Ru Atomic weight = 104. Molecular weight unhnoivn. Sp. gr. 12.26. Atomicity ", iv, Ti, and viii, also a pseudo-triad and a pseudo-heptad. Evidence of atomicity: Rathenous oxide, . . . Ru"0. Diruthenic hexachloride, . . . . 'Rn'"Cle. Ruthenic chloride, Potassic ruthenate, . . . Ruvi02Ko2. Potassic perruthenate, . Ruthenic peroxide, . . . RuvlI104. COMPOUNDS OF RUTHENIUM. 603 History.—Ruthenium was first directly recognized as a new metal by Claus, in 1845. Occurrence.—Ruthenium is found alloyed with the other platinum metals in platinum ore. Combined with sulphur it occurs as the min- eral laurite, rRTI"'2S"^ Extraction.—The insoluble residue of ruthenium obtained in the preparation of iridium (p. 596) may be purified by fusion with a mix- ture of potassic hydrate and nitre. On treating the fused mass with water the ruthenium goes into solution as potassic ruthenate. The orange-red solution is boiled with an excess of nitric acid until the color has disappeared; in this way the ruthenium is precipitated as diru- thenic trioxide, which by ignition in a graphite crucible is converted into the metal. It may be fused into a coherent mass in a lime cruci- ble by means of the oxyhydrogen flame. Properties.—Ruthenium is a white metal, hard and brittle like irid- ium, and still more difficultly fusible than this metal. The finely divided metal is oxidized when heated in air. Aqua-regia attacks it only very slowly. COMPOUNDS OF RUTHENIUM WITH THE HALOGENS. Ruthenous chloride, RuC12, is prepared by gently heating the finely divided metal in a current of chlorine. It is a black crystalline pow- der, insoluble in acids. Diruthenic hexachloride, Ru///,Clfi, is obtained as a yellow crystalline deliquescent mass by dissolving diruthenic hexahydrate in hydrochloric acid and evaporating to dryness. It forms double chlorides with the chlorides of the alkalies : /Ru/"3C16,4KC1, and 'Ru'ACIgANHqCI. Diruthenic hexiodide, /Ru''//216, is obtained as a black powder when potassic iodide is added to a solution of the chloride. Ruthenic chloride, RllCl4, is obtained as a reddish-brown mass by dissolving ruthenic hydrate in hydrochloric acid and evaporating. It forms with the chlorides of the alkalies double chlorides, corresponding with those of platinic chloride. The potassium compound has the form- ula RuC14,2KC1, and crystallizes in red regular octahedra. COMPOUNDS OF RUTHENIUM WITH OXYGEN. Ruthenous oxide, . RuO. Ru=o. Diruthenic trioxide, Ru—Ru II I! o o Ruthenic oxide, . Ru02. o=Ru=o. o II Euthenic peroxide, RuOv o=Ru=o. II O 604 INORGANIC CHEMISTRY. Ruthenous oxide, RuO, is obtained by calcining ruthenous chloride with sodic carbonate and extracting the cooled mass with water, when the oxide remains as a dark-gray powder insoluble in acids. Diruthenic trioxide, /Eu///2( );1, is formed when finely divided ruthe- nium is heated for a considerable time in contact with air. It is a bluish-black powder, which does not part with its oxygen even at a white heat. Acids are without action upon it.—Diruthenic hexahydrate, is obtained as a dark-brown precipitate when a caustic alkali is added to a solution of diruthenic hexachloride. It dissolves in acids, yielding a yellow solution. Ruthenic oxide, Ru02, is prepared by heating ruthenic sulphide in air or by heating finely divided ruthenium very strongly in a current of air. In the latter case the oxide sublimes in green quadratic pyra- mids, isomorphous with those of tin-stone and rutile.—Ruthenic hydrate, RuHo4,30H2, is a dark-red powder obtained by precipitating solutions of ruthenic salts with caustic alkali. It deflagrates on heating. Ruthenic peroxide, Ru04.—ln order to prepare this compound a so- lution of potassic ruthenate {infra) is introduced into a retort and a rapid current of chlorine is passed through the liquid. In the oxida- tion which occurs considerable heat is evolved, and the ruthenic perox- ide which is formed volatilizes in the current of chlorine, and condenses in the neck of the retort and in the well-cooled receiver as a yellow crystalline mass consisting of rhombic prisms. It is purified by fusion under a small quantity of water. The crystals fuse at 40° C. (104° F.) to a liquid which boils a little above 100° C. yielding a golden-yellow vapor with an extremely irritating odor. The experiment of distilling the peroxide alone ought never to be performed, as the heated substance is apt to decompose with violent explosion. The compound ought to be volatilized as above at a lower temperature in a current of some gas. Moist ruthenic peroxide is rapidly decomposed with evolution of oxy- gen and formation of diruthenic hexahydrate; the dry substance is more stable. It is sparingly soluble in water. OXY-SALTS OF RUTHENIUM. These are unimportant and have been little studied. Ruthenic sulphate, S204Euoiv, is obtained by oxidizing ruthenic sulphide with nitric acid and evaporating the solution. It is a deliquescent powder resembling in appear- ance mosaic gold. R UTIIENA TES AND PERR UTHENA TES. Two oxides of ruthenium—ruthenic anhydride, RuOa, and perru- thenic anhydride, Ru207—intermediate between ruthenic oxide and ruthenic peroxide, are known only in the form of the salts of their acids. Potassic ruthenate, Ru02Ko2, is formed when finely divided ruthe- nium is fused with a mixture of caustic potash and nitre or potassic chlorate. It dissolves in water, yielding a reddish-yellow solution with an astringent taste. The solution colors organic substances black, Potassic perruthenate, 'Ruvii206Ko2, is formed when chlorine acts upon the preceding salt in aqueous solution : LEAD 605 2Ru02Ko2 + Cl2 = 'RuTii206K:o2 + 2KCI. Potassic Potassic Potassic ruthenate. perruthenate. chloride. The dark-green solution deposits small black crystals isornorphous with potassic permanganate. Ammonium compounds of ruthenium have been prepared. Hiruthenic trisulphide, /Ru///S//3.—This compound occurs as the mineral laurite in some platinum ores. It crystalizes in octahedra. A part of the ruthenium is generally replaced by osmium. The same compound is obtained as a dark metallic powder by precipitating solutions of ruthenium salts with sulphuretted hydrogen and drying the precipitate in a current of carbonic anhydride. COMPOUND OF RUTHENIUM WITH SULPHUR. General Properties and Reactions of the Compounds of Ruthenium.—Solutions of ruthenic salts yield with potassic chloride and ammonia chloride dark-red crystalline precipitates of the corre- sponding double chlorides. Sulphuretted hydrogen first changes the color of the liquid to blue, and afterwards precipitates brown dirutbenic trisulphide. Zinc also changes the color of the solution to blue, and afterwards decolorizes it with precipitation of black metallic ruthenium. The formation of a volatile peroxide (p. 604) is common to this metal and osmium. CHAPTER XXXIX. TETRAD ELEMENTS. Section V. LEAD, Pb. Atomic weight 206.5. Molecular weight unknown. Sp. gr. 11.37. Fuses at 326° C. (619° F.). Bails at a white heat. Atomicity" andiy. Sometimes also a pseudo-triad. Evidence of atomicity : Plumbic chloride, Plumbic oxide, Plumbic tetrethide, Plumbic peroxide, ..... Diplumbic hexetliide, .... .... Pb"'2Et6, History.—Lead has been known from the earliest historical times. The alchemists, who believed that a connection existed between the metals and the planets, designated lead Saturn, a name which is still preserved in the expression “ saturnine poisoning,” sometimes applied to poisoning by lead. Occurrence.—Lead occurs widely distributed in nature. Native lead has been found in small quantities in volcanic tufa. The chief ore of 606 INORGANIC CHEMISTRY. lead is the sulphide, or galena, PbS;/. Other lead minerals are the carbonate or cerussite, OOPbo", and the sulphate or anglesite, S02Pbo". It also occurs as phosphate, arsenate, chromate, and molybdate. Eng- land and Spain furnish the chief supply of lead. In England the most important mines are those of Cornwall and Cumberland. Extraction.—Lead is chiefly obtained from galena. This ore is first roasted in a reverberatory furnace, by which treatment a portion of the sulphide is converted into oxide or sulphate. The temperature of the furnace is then raised, when the oxide and sulphate react with the unal- tered sulphide, and a mutual reduction to metallic lead occurs, with evolution of sulphurous anhydride. PbS" + 2PbO = 3Pb + S02. Plumbic Plumbic Sulphurous sulphide. oxide, anhydride. PbS" + S02Pbo" = 2Pb + 2S02. Plumbic sulphate. The above process can be employed only with ores of lead which are free from other metallic sulphides. In the case of ores containing py- rites, zinc-blende and other impurities, the precipitation process is em- ployed. In this process the ore is reduced by fusion with cast iron, less of this metal being employed than is required to reduce the whole of the galena present. The iron combines with the sulphur to form ferrous sulphide, which rises to the surface with the other sulphides, whilst the molten lead sinks to the bottom of the furnace, and can be drawn off. The lead obtained by either of the above processes always contains silver. This is profitably extracted by Pattinson’s process of desilveri- zation (p. 448). The oxide obtained in cupelling the portions of lead rich in silver is reduced by heating with carbon in a low blast-furnace. Lead generally contains antimony, tin, and other impurities, the presence of which renders the metal hard. The process of removing these impurities, known as softening or improving the lead, consists in partially oxidizing it in a shallow cast-iron pan on the bed of a rever- beratory furnace. The impurities are oxidized more readily than the lead, and pass into the layer of oxide which forms on the surface of the metal. Properties.—Lead is a bluish-white metal, lustrous on the freshly cut surface. It is very soft and may be cut with a knife or scratched with the nail. It may be rolled into sheets of foil, but, owing to its want of tenacity, cannot be drawn into thin wire, though it may be formed into wire by pressing through a narrow opening. Lead contracts in solidi- fying, and objects cast in this metal frequently contain cavities. It may be obtained in regular octahedra by fusing a quantity of the metal, allowing it partially to solidify and then pouring off the liquid portion. It may also be obtained in the form of an aggregation of lustrous laminae (lead-tree) by the electrolysis of solutions of its salts, or by suspending a piece of zinc or iron in such a solution. A clean and bright surface of lead speedily tarnishes on exposure to air, owing to oxidation. The fused metal becomes covered with a black film of sub- COMPOUNDS OF LEAD. 607 oxide, which at a higher temperature is converted into yellow oxide. Pure water is without action upon lead as long as air is excluded, but in presence of air plumbic hydrate is formed, which is somewhat solu- ble in water. The presence of minute quantities of carbonates and phosphates in water greatly diminishes this solubility and prevents the corrosion of the lead. These facts are of great importance from a sani- tary point of view, owing to the universal employment of lead pipes for conveying a supply of water, and the poisonous character of the compounds of lead. Fortunately almost all natural waters contain carbonates or phosphates, and the lead is thus protected from corrosion. Dihydric calcic dicarbonate—the solution of calcic carbonate in carbonic acid—an impurity present in most natural waters, is especially effica- cious in this respect, causing a film of insoluble basic plumbic carbo- nate to be formed upon the surface of the lead. Basic plumbic carbonate, GO(OPb"Ho)2 dissolves in pure water only to the extent of a sixtieth of a grain to the gallon : when a solution of plumbic hydrate in dis- tilled water is exposed to the air carbonic anhydride is absorbed and the basic carbonate is deposited in silky crystals. Lead resists to a great extent the action of sulphuric and hydrochloric acids, but dissolves readily in nitric acid. Uses.—The ease with which lead may be worked and its power of resisting the action of air, moisture, and acids, have led to its employment for various purposes: thus it is used for water-pipes, for roofing houses, and in the construction of sulphuric acid chambers. Rifle bullets and small shot are also made of this material, about 0.5 per cent, of arsenic being added in the latter caye in order to aid the metal in assuming the spherical form. Various alloys of lead are also used in the arts. Type metal is an alloy of 2 parts of lead, lof anti- mony and lof tin. Plumber’s solder is an alloy of lead and tin (p. 323). COMPOUNDS OF LEAD WITH THE HALOGENS. Plumbic chloride, PbCl2. Molecular volume I I I.—This com- pound has been found in the crater of Vesuvius as the mineral cotun- nite. Hydrochloric acid attacks lead only very slowly, but hot aqua- regia dissolves it readily, depositing crystals of the chloride on cool- ing. It is best prepared by dissolving the oxide or the carbonate in hydrochloric acid. It is also precipitated as a crystalline powder when hydrochloric acid or a soluble chloride is added to a not too dilute solution of a lead salt.—Plumbic chloride crystallizes from water in long, colorless, lustrous prisms. It is soluble at ordinary temperatures in 130 parts, at 100° C. in less than 30 parts of water. When fused ■with exclusion of air, it solidifies on cooling to a white horn-like mass, but if air be admitted, it is converted into oxychloride. Oxychlorides of varying composition are obtained by fusing together plumbic oxide and plumbic chloride, or by precipitating a solution of plumbic chloride with an insufficiency of lime-water or ammonia. Those which are rich in chlorine are white; those which are rich in oxygen are yellow. Some of these compounds are employed as pigments. Cassel yellow is 608 INORGANIC CHEMISTRY. an oxychloride obtained by heating plumbic oxide with amnionic chlo- ride. • A white oxychloride, prepared by precipitating plumbic chloride with lime-water, is employed as a substitute for white lead. Plumbic perehloride, PbCl4, exists only in solution. When plumbic peroxide is dissolved in well-cooled concentrated hydrochloric acid, a strongly oxydizing liquid, which evolves chlorine on heating, is obtained. Plumbic bromide, Pbßr2, resembles the chloride. Plumbic iodide, Pbl2.—This compound is precipitated as a crystal- line yellow powder when a soluble iodide is added to a solution of a lead salt. It is almost insoluble in cold, but dissolves slightly in hot water, yielding a colorless solution, which on cooling deposits the iodide in yellow laminse. Plumbic iodide, when heated, becomes first red, then black, and finally fuses to a dark-colored liquid, which on cooling solidifies to a yellow crystalline mass. It dissolves in solutions of the alkaline iodides to form double salts. Plumbic fluoride, PbF2. is precipitated as a white almost insoluble powder, when hydrofluoric acid is added to the solution of a lead salt. COMPOUNDS OF LEAD WITH OXYGEN. r b* Plumbous oxide, . . .< O. ) yO. *-rD Pi/ Plumbic oxide, . . . PbO. Pb:=o. /? Diplumbic trioxide, . . PbOPbo". o==Pb ( Pb ( o>Pb I Triplumbic tetroxide, . Pb\ 1 X 0 Uo>Pb Plumbic peroxide, . . Pb02. o=Pb=o. Plumbous oxide, is best prepared by heating plumbic oxa- late to 300° C. with exclusion of air: 2{coPbo" = |pb° + 3CO* + co- Plumbic Plumbous Carbonic Carbonic oxalate. oxide. anhydride. oxide. It is a black powder. When lead is fused in air, avoiding too high a temperature, the same compound is formed as a gray film on the sur- face of the metal. When heated to redness with exclusion of air, plumbous oxide is decomposed into plumbic oxide and metallic lead; if air is admitted, it burns like tinder and is totally converted into plumbic oxide. It slowly undergoes the same conversion when exposed to the air in a moist state. With acids it yields plumbic salts with separation of metallic lead. COMPOUNDS OF LEAD. 609 Plumbic oxide {Litharge), PbO.—This compound is prepared by heating lead in air or by igniting plumbic carbonate or nitrate. It is obtained as a by-product in various metallurgical operations'—notably in Pattinson’s process for the desilverization of lead (p.44B)—Plumbic oxide is a yellow powder, which when strongly heated fuses, and on cooling solidifies to a yellow micaceous mass, sometimes with a shade of red. It is slightly soluble in water, to which it imparts an alkaline reaction. Acids dissolve it, forming the various salts of lead. Car- bonic oxide at 100° Cl, and hydrogen at 310° C., reduce it to metallic lead. Plumbic oxide absorbs carbonic anhydride from the air.—Litharge is employed in the preparation of various salts and pigments of lead, in the manufacture of flint-glass, and in glazing earthenware. Diplumbic oxydihydrate, Pb2OHo2, is precipitated when ammonia is added in excess to a solution of plumbic nitrate. Caustic alkalies may be substituted for ammonia, but in this case an excess of the precipitant must be avoided, as this would dissolve the plumbic hydrate. It is a white bulky precipitate, difficult to obtain free from basic salts. It is slightly soluble in water.—The hydrate FbHo2 has not been pre- pared. Diplumhio trioxide, PbOPbo//, is precipitated as a reddish-yellow powder, when sodic hypochlorite is carefully added to a solution of plumbic hydrate in caustic soda. It is decomposed at a red heat into plumbic oxide and oxygen. Hydrochloric acid dissolves it completely in the cold, yielding a yellow liquid, which speedily evolves chlorine and then contains plumbic chloride. Oxy-acids take up half the lead of this oxide to form plumbic salts, whilst the other half remains undissolved as plumbic peroxide. Triplumbic tetroxide, PbPbor/2.—This compound appears to be contained in red-lead or minium, which is, however, a substance of varying composition, intermediate between plumbic oxide and diplura- bic trioxide. When finely divided litharge or plumbic carbonate is heated in air for twenty-four hours to dull redness, it is converted into a heavy scarlet crystalline powder. It becomes dark when heated, but recovers its original color on cooling. At a red-heat it is decomposed like the trioxide into plumbic oxide and oxygen. In its behavior towards acids it also resembles that compound.—Red-lead is employed as a pigment, also for electrical storage batteries, and in the manufac- ture of the finer sorts of flint-glass. For the latter purpose the excess of oxygen which it contains serves to effect the combustion of organic matters, and thus to prevent the reduction of the lead which would cause the glass to blacken. Plumbic peroxide {Puce-colored oxide of lead), Pb02, is most readily obtained by treating diplurabic trioxide or red-lead with nitric acid, when the peroxide remains as a dark-brown amorphous powder. The same compound is formed when chlorine is passed into an alkaline solution in which plumbic hydrate is suspended. It is also deposited on the positive electrode when the solution of a lead salt is electrolyzed. It occurs native in black six-sided prisms as plattnerite. At a red heat it is decomposed like the other higher oxides of lead into plumbic oxide and oxygen. When introduced into an atmosphere of sulphur- 610 INORGANIC CHEMISTRY. ous anhydride, it is converted with incandescence into plumbic sul- phate ; Pb03 + S02 = S02Pbo" Plumbic Sulphurous Plumbic peroxide. anhydride. sulphate. Sulphuric acid dissolves it with evolution of oxygen and formation of plumbic sulphate; hydrochloric acid dissolves it with evolution of chlorine and formation of plumbic chloride; nitric acid is without ac- tion upon it.—A porous mass of plumbic peroxide, generated by elec- trolysis, forms the negative plate in the Plante secondary battery and other electrical storage batteries constructed on the same principle (see p. 106). OXY-SALTS OF LEAD. [NO, Plumbic nitrate, < Pbo", is best prepared by dissolving litharge In°2 in an excess of nitric acid and evaporating to the crystallizing point. The salt forms colorless octahedral crystals, soluble in twice their weight of cold water, much less soluble in water containing nitric acid. It is almost insoluble in alcohol. At a red heat it fuses and is decomposed into plumbic oxide, nitric peroxide, and oxygen. When thrown upon red-hot charcoal, it deflagrates. It is employed as a mordant in dyeing and calico-printing.—A boiling aqueous solution of plumbic nitrate dissolves plumbic oxide, and on cooling deposits acicular crystals of plumbic nitrate, hydrate, NG3(OPb//Ho). Other basic nitrates, of the formulae N303Pbo//(OPb//110)3 and KPbo/'3(OPb//IIo), are obtained by precipitating solutions of the normal nitrate with ammonia. i NO Plumbic nitrite, 1 PboVOIL,.—This compound is most readily obtained by accurately (no precipitating argentic nitrite with plumbic chloride and evaporating the solution in vacuo over sulphuric acid. It forms soluble yellow prisms or laminae. If the solution be boiled, nitrogen is evolved and a basic nitrite is formed. If a solution of plumbic nitrate in fifty times its weight of water be boiled with one and a half parts of lead for twelve hours, the liquid deposits on cooling flesh-colored needles of diplumbic nitrite hydrate, NPbo//(OPb//Ho). If carbonic anhydride be passed into the solution of this salt, three-fourths of the lead is precipitated as carbonate, and the liquid contains the normal nitrite. If a solution of plumbic nitrate be digested with metallic lead for a few hours at a temperature of 75° C. (167° F.), a yellow liquid is obtained, which on cooling deposits lustrous yellow tabular crystals ( NHojPbo'7 of dihydric diplumbic nitrate nitrite, -j Pbo// ,—a salt formerly termed “ basic hypo- ( NO nitrate of lead.” Various other basic nitrites of lead have been prepared. Plumbic carbonate, OOPbo", occurs native as the mineral cerus- site in lustrous transparent rhombic crystals, isomorphous with those of arragonite. The same salt is obtained as a white crystalline pre- cipitate by pouring a solution of plumbic nitrate into a solution of sesquicarbonate of ammonia. The carbonates of sodium and potassium cannot be employed for this purpose, as these precipitate mixtures COMPOUNDS OF LEAD. 611 of basic plumbic carbonates, the composition of which varies with the concentration and the temperature. White lead is a basic carbonate of lead—triplumhic dicarbovate dihydrate, 'm manufactured on a large scale as a pigment by one or other of the following processes: (1) Dutch Process.—This is the oldest process and yields the finest product, but the operations are somewhat tedious. Glazed earthen- ware pots are filled to a quarter of their depth with weak malt vinegar. In each pot, above the surface of the liquid and resting on a wooden support, a thin sheet of lead coiled into a spiral is placed vertically, or a series of cast gratings Is put into the pot, and the pot is covered with a plate of lead. The pots are then embedded in spent tan-bark or horse-dung on the floor of a shed. The first layer of pots is then cov- ered with boards, and a second layer, arranged like the first and also embedded in tan-bark or horse-dung, is built up over these, and so on till the shed is full. The pile generally reaches a height of from 18 to 20 feet, and contains about 12,000 pots with from 50 to 60 tons of lead. The action which takes place is as follows: The heat evolved by the fermentation of the bark or dung volatilizes the acetic acid in the vinegar, which gradually in presence of the oxygen of air, which for this purpose must have free access to the heap, converts the lead superficially into basic plumbic acetate : ,/CH, „ph , 0 _ 2/CHs 1 \ COHO ~ 1 ~ CO(OPb"Ho)' Acetic Plumbic acetate acid. hydrate. The carbonic anhydride which is given off during the fermentation then acts upon the basic acetate, converting it into basic carbonate (white lead) and normal acetate: rcri3 , 2CO _ CO(OPb"Ho)pb „ b\CO(OPb"Ho) + “ CO(OPb"HorDo Plumbic acetate Carbonic Tripiurnbic dicarbonate • hydrate. anhydride, dihydrate. J gh3 + + 20H2. I GH3 Plumbic Water, acetate. then reacts with a fresh portioi The normal acetate then reacts with a fresh portion of lead in presence of oxygen and water, and regenerates the basic acetate: JO H3 jOgpbo" + Pb + o + 0H2 = 2{^bpb„Ho). ICHS Plumbic Water. Plumbic acetate acetate. hydrate. 612 INORGANIC CHEMISTRY. The basic acetate is again acted upon by the carbonic anhydride as above. In this way the process is theoretically continuous, and a small quantity of acetic acid ought to suffice for the conversion of an un- limited quantity of lead. In practice 100 lbs. of acetic acid are required to convert 50 tons of lead into white lead. At the end of from four to five weeks the conversion is nearly complete; the pile is taken to pieces and, on uncoiling the spirals, the white lead peels off in flakes from the unaltered lead if any of the latter is left. The crude product is ground while moist, and well washed to free it from acetate. (2) Thenard’s Process.—A solution of basic plumbic acetate of lead is first prepared by boiling sugar of lead with litharge. The basic car- bonate is then precipitated from this solution by passing in carbonic anhydride. As a pigment, the product lacks opacity, and is conse- quently deficient in “ body ” or “ covering power.’’ (3) Milner's Process.—ln this process, which yields good results, an oxychloride of lead is converted into white lead by the action of gaseous carbonic anhydride. A mixture of litharge, common salt, and water is ground for some hours. Into the mixture of caustic soda and plumbic oxychloride thus obtained, carbonic anhydride is passed until the liquid is neutral. At this point the operation must be interrupted, otherwise the product will be spoiled. White lead is a white amorphous powder. Its chief drawbacks are its poisonous character, and the fact that it is blackened by sulphuretted hydrogen. Plumbic sulphate, S02Pbo", occurs native as anglesite in trans- parent rhombic crystals. It is obtained as a heavy white crystalline precipitate when sulphuric acid or a soluble sulphate is added to the solution of a lead salt. The precipitate is almost insoluble in water, and still less soluble in dilute sulphuric acid; but concentrated sulphuric acid dissolves about 6 per cent, of its weight of the sulphate. It is also slightly soluble in dilute hydrochloric and in dilute nitric acid, whilst sodic thiosulphate and many ammonia salts, particularly the acetate and the tartrate, dissolve it readily. When plumbic sulphate is boiled with a solution of ammonic sulphate, the liquid deposits on cooling minute lustrous crystals of plumbic diammonic disul- { S02Amo ■phate< Pbo// . Pure water decomposes this salt with separation of ( B02Amo insoluble plumbic sulphate. By treating the normal salt with ammonia, diplumhic sulphate, is obtained. Plumbic dithionate, | g|^2Pbo//,40112, or | Pbc/g is best prepared by neutral- izing a solution of dithionic acid with plumbic carbonate. It forms large colorless hexagonal crystals, readily soluble in water. Plumbic chromates.—See Chromates. Plumbic phosphates.—The normal orthophosphate, P./VPbc/t), is obtained as a white- amorphous precipitate when hydric disodic phosphate- is added to a solution of an excess of plumbic acetate. It is insoluble in water and acetic acid, readily COMPOUNDS OP DEAD. 613 soluble in nitric acid and caustic potash.—Hydric plumbic phosphate, POHoPbo//, is precipitated by free phosphoric acid from a solution of plumbic nitrate as a white crystalline powder.—A double phosphate and chloride of lead of the formula P303Pbo//4^QPb occurs in nature in hexagonal crystals as the mineral pyromor- phite. It is isomorphous with apatite (p. 357). Plumbic arsenates.—These resemble the phosphates. A native double arsenate and chloride corresponding to pyromorphite is the mineral, mimelesite, As303Pbo",(Opb"), which forms hexagonal crystals. Intermediate gradations between pyromorphite and mimelesite occur, in which the phosphorus and arsenic replace each other isomor- phously. Plumbic borates.—When the solution of a lead salt is precipitated with borax, octo- hydric diplumbic hexaborate, B603HoBPbo//'2, is formed. When this is warmed with ammonia it is converted into a white powder of dihydric plumbic diborate, BaOHosPtx/''. —By fusing together plumbic oxide and boric anhydride, a transparent vitreous mass f Farad ay’s heavy glass) is obtained, which possesses a much higher refractive power than flint-glass. Plumbic silicate—No definite silicate of lead has been prepared. When silica is fused with plumbic oxide a vitreous mass is obtained. Plumbic silicate is one of the constituents of flint-glass. COMPOUND OF LEAD WITH SULPHUR, Plumbic sulphide, PbS".—As the mineral galena this compound forms the principal ore of lead. It occurs in regular cubes with a bluish-gray color and a brilliant metallic lustre ; also in crystalline masses. It possesses a very perfect cubical cleavage. The same com- pound is formed as a leaden-gray crystalline mass when lead is fused with sulphur, and as an amorphous black powder by precipitating a solution of lead salt with sulphuretted hydrogen. It fuses without decomposition at a bright red heat when air is excluded, and may even be sublimed in a current of hydrogen or carbonic anhydride. In this way it is obtained in small cubical crystals. When fused with access of air it is converted into plumbic sulphate. It dissolves in hot con- centrated hydrochloric acid with evolution of sulphuretted hydrogen. Dilute nitric acid converts it into nitrate with separation of sulphur; the concentrated acid oxidizes it to sulphate.—When sulphuretted hydrogen, in quantity insufficient for complete precipitation, is passed into a solution of plumbic chloride, red and yellow sulpho-chlorides of varying composition separate out : Cl—Pb—S—Pb—Cl, and Cl—Pb—S—Pb—Pb—S—Pb—Cl. General Properties and Reactions of the Compounds of Lead.—The salts of lead are mostly colorless. They have a sweet, astringent, metallic taste, and are poisonous. When continually intro- duced in minute quantities into the system, the salts of lead act as a cumulative poison. The soluble normal salts with strong acid redden litmus; the basic salts, on the other hand, have an alkaline reaction. Caustic alkalies and ammonia precipitate white basic salts of lead, solu- ble in excess of caustic alkali, insoluble in ammonia. Sulphuretted hydrogen and ammonic sulphide produce a black precipitate of plumbic 614 INORGANIC CHEMISTRY. sulphide, which is converted by fuming nitric acid into white insoluble plumbic sulphate, whilst dilute nitric acid converts it into soluble plum- bic nitrate with separation of sulphur. Sulphuric acid and soluble sul- phates precipitate plumbic sulphate, very sparingly soluble in water, still less soluble in dilute sulphuric acid, insoluble in alcohol, but solu- ble in solutions of various ammonium salts, such as the acetate and the tartrate. Hydrochloric acid and soluble chlorides yield with not too dilute solutions a white precipitate of plumbic chloride, soluble in hot water. Potassic chromate precipitates yellow plumbic chromate; po- tassic iodide yellow plumbic iodide. All compounds of lead, when heated with sodic carbonate or potassic cyanide upon charcoal in the reducing flame, yield a malleable bead of metallic lead. The lead compounds give a faint flame spectrum, containing lines in the green and a characteristic spark spectrum. CHAPTER XL. HEXAD ELEMENTS. Section 11. URANIUM, U. Atomic weight = 238.5. Molecular weight unknown. Sp.gr. 18.7. At- omicity iT, vi, and Tiii ?* ; also a pseudo-triad and a pseudo-pentad. Evidence of atomicity: Uranous chloride, ...... , . . . UivCl4. Diuranoua hexachloride, . . . . / • • • 1 U;//C1 • Uranic oxide, . . . . UTiOs. I)iuranic decachloride, . . . . j UVC15 • • • lira; History.—Klaproth first pointed out in 1789 the existence of a new metal in the mineral pitchblende, and to this metal he gave the name uranium. The metal was isolated by Peligot in 1842. Occurrence.—Uranium is of rare occurrence, and is never found native. Its chief ore is pitchblende, an impure uranous diuranate, gg’TO’. It also occurs as phosphate in uranium mica, and as carbon- ate in liebigite. Preparation.—Metallic uranium is obtained by the action of sodium upon uranous chloride, UCJ4. The two substances are heated together in * Uranium and molybdenum, which have been included in the hexadic group, ap- pear to be capable of exercising octadic functions: thus inperuranicanhydride (UO4) and molybdic persulphide (MaSJ. COMPOUNDS OF URANIUM. 615 a porcelain crucible with the addition of potassic chloride as a flux. The porcelain crucible is packed in powdered charcoal within a larger cru- cible. The whole is heated, at first to redness, afterwards to a higher temperature so as to fuse the uranium, which is thus obtained as a black regulus. Properties.—Metallic uranium has a silvery lustre, but tarnishes by exposure to the air, becomingin course of time steel-blue, and ultimately black. It is hard and somewhat malleable. When heated in air it burns with scintillations, forming uranous diuranate. It does not de- compose water even at its boiling-point. Acids readily dissolve it. COMPOUNDS OF URANIUM WITH THE HALOGENS. Diuranous hexachloride, is obtained in dark-brown needles by heating uranous chloride to redness in a current of hydrogen. It dissolves in water, yielding a purple solution, which rapidly absorbs oxygen from the air. \ Uranous chloride, UC14, is prepared by heating a mixture of charcoal and any of the oxides of uranium in a current of dry chlorine. It is volatile at a red heat, and may be obtained by sublimation in dark- green octahedra, possessing a metallic lustre. It is very deliquescent, and hisses when thrown into water. Its solutions absorb oxygen from the air, and turn yellow. Uranous bromide, UBr4) and uranous fluoride, UF4, have also been prepared. Diuranic decachloride (Uranic pentachloride), 'U^Cl^.—This com- pound is formed along with uranous chloride in the preparation of the latter compound, especially when the temperature is not permitted to rise too high. As it is more volatile than uranous chloride, it collects in a part of the tube further removed from the source of heat. If the current of chlorine be sufficiently slow, the decachloride forms black, needle-shaped crystals. The compound rapidly deliquesces on exposure to air. It begins to decompose at 120° C. into uranous chloride and free chlorine. COMPOUNDS OF URANIUM WITH OXYGEN. Uranous oxide, . » . . U02. o=ll=o. O Uranic oxide (uranic an- ] IT/ s ~ J| hydride), | U0*- O=U=G. o II Peruranic anhydride, . . U04. o=U=;o. ii O The remaining oxides of uranium—U205 = UOUolv, uranous ura- nate, and U3Os uo!u oiv, uranous diuranate—are regarded as com- binations of the two first oxides with each other. 616 INORGANIC CHEMISTRY. Uranous oxide, U02.—Tin’s oxide remains when any of the higher oxides of uranium, or uranic oxalate, is heated in a current of hydrogen. It forms a brown powder, which when heated in air burns with form- ation of uranous diuranate. Strong acids dissolve it, yielding green solutions of uranous salts, from which alkalies precipitate dark-brown flocculent uranous hydrate, Uflo4. Urania oxide (Urania anhydride), U03, is obtained as abrownish- yellow powder when uranic nitrate is heated in an oil bath to 250° C. until nitrous fumes cease to be evolved. At higher temperatures it parts with oxygen, and is converted into uranous diuranate. Uranic oxide acts both as a basic oxide and as the anhydride of an acid: thus, on the one hand, it combines with acids to form salts in which the dyad radical uranyl (Uvio2)// plays the part of a dyad metal, and, on the other, it unites with alkalies to form the uranates {q.v.).—A uranic hydrate is also known, but is very difficult to obtain of constant compo- sition. Uranous diuranate (Green oxide of uranium), |Jq2UoIv, occurs native in an impure state as pitchblende. It is obtained as a green pow- der when uranous or uranic oxide, or ammonic uranate is gently heated in air. It is difficultly soluble in hydrochloric and sulphuric acids, readily soluble in nitric acid. Uranous uranate (Black oxide of uranium), UOUolv, or |jq20, is obtained as a black powder when any of the other oxides of uranium, or ammonic uranate, is strongly ignited in air. It is used in painting on porcelain. OXY-HALOGEN COMPOUNDS OF URANIUM. Uranylic chloride, U02C12, is formed when uranous oxide is heated in a current of chlorine. It is a yellow, deliquescent, and very soluble mass, which is readily fusible, but volatilizes with some difficulty. It unites with the alkaline chlorides to form well-crystallized double salts : thus U02CI,2KC12,20H2, and U02C12,2NH4C1,20H2. Uranylic bromide, and uranylic fluoride, UO2F2, have also been prepared. OXY-SALTS OF URANIUM. a. Uranous Salts. SO Uranous sulphate, gQ2Uoiv, occurs native, but partially oxidized to uranic sulphate as uranium vitriol or johannite. It is formed when uranous oxide is dissolved in sulphuric acid. The most convenient mode of preparing the salt consists in dissolving the green oxide in sulphuric acid, adding alcohol, and exposing the whole to sunlight. The liquid at first contains a mixture of a uranous and a uranic salt, but under the above conditions the uranic salt is reduced to the uranous stage, and the uranous sulphate, which is insoluble in dilute alcohol, COMPOUNDS OF UEANIUM. 617 separates in crystals containing 4 aq. From aqueous solutions it crys- tallizes in green prismatic crystals with 8 aq. Excess of water decom- poses it with separation of a green basic salt. Uranous phosphate.—A hydric uranous phosphate, P202H02Uoiv,2OH2, is formed as a green gelatinous precipitate when hydric disodic phosphate is added to a solution of uranous chloride. b. Uranic (Uranylic) Salts. In the salts the dyad radical uranyl (Uv,02)" plays the part of a dyad metal. They are characterized by possessing a yellow color with a magnificent green fluorescence. Uranylic nitrate, S_O>UTIO2>60H2> is obtained by dissolving any of the oxides in nitric acid and evaporating the solution. It crys- tallizes in large greenish-yellow rhombic prisms. Uranylic sulphates.—The normal salt, S()2<| |> U vi02,30H2, is de- posited in small lemon-yellow crystals when a solution of the nitrate is mixed with sulphuric acid and evaporated. A hot solution of this salt in moderately concentrated sulphuric acid deposits on cooling deliquescent, yellowish-green, fluorescent crystals of hydricuranylic sul- f SOjlo O phate, U02. If, on the other hand, the normal salt be dissolved in I o S02110 fuming sulphuric acid, small yellow crystals of uranylic pyrosulphate, rso2-o < O b>UT!02, are obtained. These attract moisture with great [so-o7 avidity, and dissolve with a hissing noise when thrown into water. Uranylic sulphate forms double salts with the sulphates of the alkali metals; thuspotassic uranylic sulphate, f S02Ko I o U02 ,20112, forms yellow monoclinic crystals. O S02Ko Phosphates and arsenates of uranyl occur native as rare minerals. Besides behaving as a base towards acids, uranic oxide behaves towards strong bases as the anhydride of an acid, forming salts called uranates, in which the group uranyl (U02)" plays the part of an acid radical. These salts are, however, not derived from a normal uranic acid of the formula U0211o2, corresponding to sulphuric acid, but from an anhydro-acid or diuranic acid of the ruo2Ho formula < O , corresponding to disulphuric or dichromic acid. Free ( U02110 diuranic acid has not been obtained. THE UR AN A TES. 618 INORGANIC CHEMISTRY. ruo2Ko Potassic uranate, < O , is formed when uranic oxide is fused (UO2Ko with an excess of potassic carbonate, and remains behind as a yellow powder when the mass is extracted with water. ( U02J^ao Sodic uranate, < O , is obtained in a similar manner by fusing ( U( )2Nao uranic oxide with sodic carbonate. It is prepared on a large scale from pitchblende, and is employed under the name of uranium yellow in painting on porcelain and in the preparation of a beautiful greenish- yellow fluorescent glass. r U02(NvHt0) Ammonic uranate, 1 O , is formed as a yellow precipitate when ammonia [ U02(NvH40) is added to the solution of a uranyl salt. On heating, it is converted into pure uranous diuranate. f U02(OBi///Ho2) Bismuthous uranate hydrate, 1 O ,OH2, or TJ02Ho(OBi///Ho2), occurs (U02(oBi///Ho2) native as uranospherite in brick-red hemispherical aggregations. A series of peruranates has recently been obtained by the action of hydroxyl upon uranylic salts in alkaline solution. Sadie peruranate, U02iSrao4,80H2, forms golden-yellow needles. The peruranates are very unstable, and have not yet been thoroughly examined. Uranous sulphide, XJS'^.—This compound is obtained as a grayish- black amorphous powder by passing sulphuretted hydrogen over uranous chloride heated to redness. At a white heat a crystalline product is obtained. It is slowly decomposed in moist air with evo- lution of sulphuretted hydrogen. It is insoluble in dilute hydrochloric acid, but concentrated acids dissolve it readily. Uranylic sulphide, U()2S//, is a dark-brown precipitate obtained by adding ammonic sulphide to a solution of uranylic nitrate. General Properties and Reactions of the Compounds of Uranium: COMPOUNDS OF URANIUM WITH SULPHUR. a, Uranous Salts.—The uranous salts are green. In solution they absorb oxygen from the air and are converted into uranic salts, whilst their color changes from green to yellow. Caustic alkalies and ammonia produce in their solutions a dark-brown flocculent precipitate of uranous hydrate. This absorbs oxygen and. is converted into uranic hydrate, which at the same time combines with the base to form an insoluble uranate. Sulphuretted hydrogen gives no precipitate in acid solutions; ammonic sulphide precipitates a black sulphide. b. Uranic (Uranylic) Salts.—The uranic salts are yellow. From their solutions caustic alkalies or ammonia precipitate a yellow insoluble ura- nate of the base. The hydric carbonates of the alkalies and ammonic carbonate precipitate yellow double carbonates of uranium with alkali or ammonium, which are readily soluble in an excess of the precipitant. COMPOUNDS OF MOLYBDENUM. 619 Sulphuretted hydrogen gives no precipitate in acid solution; ammonia sulphide precipitates dark-brown uranylic sulphide, readily soluble in dilute acids, even in acetic acid. Potassic ferrocyanide gives a reddish- brown precipitate. The uranium compounds yield with borax and microcosmic salt beads which in the reducing flame are green, in the oxidizing flame yellow. The uranium compounds do not color the non-luminous flame. Atomic weight = 95.5. Molecular weight unknown. Sp. gr. 8.6. Atomicity ", IT, vl, and VUI ?* Evidence of atomicity: MOLYBDENUM, Mo. Hypomolybdous chloride, . . . . . . Mo "Cl.,. Molybdous chloride, . . . MoivCl4. Molybdic anhydride, . . . MoTiOs. History.—Metallic molybdenum was first obtained by Hjehn in 1782. Occurrence.—Molybdenum is of rare occurrence. It is found in combination with sulphur as molybdenite, with oxygen in molybdenum ochre or native molybdic anhydride, MoOa; and as plum- bic molybdate, Mo()2Pbo// in wulfenite. Many iron ores contain traces of molybdenum, which thus finds its way into the pig-iron. Preparation.—Metallic molybdenum is obtained by heating molybdic anhydride or one of the chlorides to redness in a current of hydrogen. In the case of the oxide the reduction is always incomplete, and it is necessary to purify the product by heating in a current of dry gaseous hydrochloric acid, when the unattacked oxide volatilizes as molybdic hydroxy-chloride, MoOHo2C12. Properties.—Pure molybdenum is a silver-white metal. It appears to be infusible at the highest temperature that can be artificially pro- duced, but if it contains carbon it may be fused by the oxy-hydrogen flame. It is permanent in air at ordinary temperature, but, when heated in air, undergoes oxidation and is ultimately converted into molybdic anhydride. It is not attacked by dilute hydrochloric or sul- phuric acid, but hot concentrated sulphuric acid dissolves it with a brown color. It is readily soluble in nitric acid and aqua-regia. COMPOUNDS OF .MOLYBDENUM WITH THE HALOGENS. Hypomolybdous chloride, MoC12, is formed when dimolybdous hexa- chloride is heated in a current of dry carbonic anhydride: 'M.o'"2C\6 = MoC12 + MoC14. Dimolybdous Hypomolybdous Molybdous liexachloride. chloride. chloride. The tetrachloride volatilizes, whilst hypomolybdous chloride remains as a yellow amorphous powder. Hypomolybdous chloride is stable * See note, p. 614. 620 INORGANIC CHEMISTRY. when exposed to air at ordinary temperatures, but is decomposed when heated in air. It is insoluble in water, but soluble in hydrochloric acid. A hypomolybdous bromide, Moßr2, has also been prepared. Dimolybdous hexachloride, /Mo///2Clfi, is obtained as a reddish-brown powder, resembling in appearance amorphous phosphorus, when mo- lybdic pentachloride is heated to 250° C. in a current of hydrogen. It is insoluble in water and in hydrochloric acid. When strongly heated, it yields a mixture of hypomolybdous chloride and molybdous chlo- ride. Dimolybdous hexabromide, /Mo///28r6, is also known. Molybdous chloride, MoC!4, is formed as above by heating the dirao- lybdous hexachloride. It is a brown crystalline powder, which when exposed to air deliquesces to a brown liquid. It may be volatilized with partial decomposition in a current of carbonic anhydride. Molybdous iodide, MoI(, is obtained by dissolving molybdous hydrate, Mollo4, in hydriodic acid and evaporating the solution. Molybdic pentachloride, MoC15. 3lolecular volume I i I.—This com- pound is obtained by heating molybdenum or molybdous sulphide in a current of chlorine. It forms a lustrous, radio-crystalline mass, which fuses at 194° C. (481° F.) and boils at 268° C. (514° F.). It fumes on exposure to air,and gradually deliquesces. The molecular formula, MOC15, as deduced from the vapor density of this compound, is abnor- mal, as this formula would necessitate the assumption either of pentadic molybdenum or of the presence of an odd number of free affinities in the molecule (see p. 179, footnote). COMPOUNDS OF MOLYBDENUM WITH OXYGEN. Hypomolybdous oxide, . . . MoO. O Dimolybdous trioxide, . f MoOn • • \ MoO A o II t> c' II p Molybdous oxide, . . . . . Mo02. d II o U-, t° O Molybdic anhydride, . . MoOa. O- 11 =Mo=0. Hypomolybdous oxide, MoO, appears to be formed as a black powder by the action of hot caustic potash upon hypomolybdous chloride. Dimolybdous trioxide, /Mo///2Os.—When dimolybdous hexachloride is decomposed with a caustic alkali, dimolybdous hexahydrate, 'Mo//72Ho6, is obtained as a dark-brown powder, and this, when heated with exclu- sion of air, parts with water and is converted into dimolybdous triox- ide. It forms a gray metallic powder, insoluble in acids. Molybdous oxide, Mo02.—This oxide is obtained, like the preceding, by heating the corresponding hydrate in absence of air. Thus pre- pared it forms a brown powder. When sodic trimolybdate, Mo3OgNao2, is fused with a third of its weight of zinc, and the mass extracted with THE MOLYBDATES. 621 water, molybdous oxide remains in the form of dark-blue prisms which appear violet-red by transmitted light. It is insoluble in water, hydro- chloric acid, and caustic potash. Hot nitric acid oxidizes it to raolyb- dic acid.—Molybdous hydrate, MoHo4, is obtained as a reddish-brown precipitate by treating molybdous chloride with ammonia. Molybdic anhydride, Mo03,—This compound is most readily prepared by roasting the native sulphide, MoS2, in air. After the sul- phur has burnt off’, the impure molybdic anhydride is extracted with ammonia, and the ammonium salt thus obtained is purified by crystal- lization. The ammonium salt may be converted into the anhydride either by heating it in small portions with free access of air, or by de- composing it with nitric acid, evaporating to dryness, and washing the residue thoroughly with water, when the anhydride remains undis- solved. It forms a while powder which turns yellow on heating, but becomes white again on cooling. It fuses at a red heat, and may be sublimed in lustrous laminae. It is insoluble in water and acids, but dissolves readily in caustic alkalies and ammonia. Molybdic acid, M 602Ho2, separates as a white crystalline j owder from the solution of a molybdate to which hydrochloric or nitric acid has been added. The compound is insoluble in water, but dissolves in an excess of acid. From hot solutions a molybdic acid of the formula Mo5OuHo2is deposited. A soluble colloidal modification of molybdic acid is obtained by dissolving sodic molybdate in hydrochloric acid and subjecting the solution to dialysis ; a yellow acid liquid remains, which yields on evaporation a gummy deliquescent mass. When a solution of molybdic acid in hydrochloric acid is treated with zinc the liquid be- comes first blue, then green, and finally brown, owing to the formation of various molybdous and hypomolybdous molybdates. Numerous oxy-lmlogen compounds of molybdenum have been pre- pared. They are generally volatile, and are mostly decomposed by water. The following list contains some of the compounds of this class : Molybdic oxytetrachloride, MoOClr Molybdic dioxydichloride, Mo02C12. Molybdic dioxydibromide, Mo02Br2. ( MoOC13 Dimolybdic trioxy-hexachloride, . . . < O ( MOOCI3 THE MOLYBDATES. The salts of molybdic acid may be divided into the following classes : Normal molybdates, Mo02Ro2. Hi molybdates, Mo2G6Ro2. Tri molybdates, Mo208Ro2. Tetramolybdates, Mo4OuRo2. Heptarnolybdates, Mo7G12Ro6. Octomolybdates, Mo8023R<>2. Decamolybdates, Mo10029,Ro2. in which R stands for a monad metal. 622 INORGANIC CHEMISTRY. All these salts, with the exception of the heptamolybdates, are de- rived from dibasic acids. Polemic molybdates.—The normal salt, Mo03Ko2, is obtained by fusing together equal molecular proportions of potassic carbonate and molybdic anhydride, dissolving the mass in water, and evaporating the filtered solution over sulphuric acid. It forms small soluble deliques- cent crystals.—The dimolybdate has not been obtained.—The trimolyb- date, Mo308Ko2,301T2, is prepared like the normal salt, employing the requisite proportions of anhydride and carbonate. It crystallizes in flexible silky needles.—Other potassic molybdates have been obtained. Sodic molybdates.—These are prepared like the potassium salts. Normal sodic molybdate, Mo02Nao2,20H2, forms nacreous lamina; or acute rhombohedra; sodic dimolybdate, Mo205Nao2, small silky needles; sodic trimolybdate, Mo;10BNao2,70H2, very tine, sparingly soluble needles. Sodic molybdates corresponding to all the various classes in the above list have been prepared. Of the other molybdates, those of barium, strontium, and calcium are either only sparingly soluble or insoluble in water, the magnesium and zinc salts are soluble and crystallize well. Normal plumbic molybdate, Mo02Pbo//, occurs native in yellow quadratic crystals as wulfenite. Phosphomolybdic Acid. Molybdic acid forms with phosphoric acid a remarkable compound hexabasic acid,.which maybe regarded as a combination of 2 molecules of phosphoric acid with 22 molecules of molybdic anhydride. Both this acid and its salts contain large and varying proportions of so-called water of crystallization, which is very possibly present as water of con- stitution. Owing to the complexity of these salts and the absence of ail certain knowledge with regard to their constitution, it will be sim- plest to formulate them as molecular combinations. Phosphomolybdic acid, 2POHo3,22M003.—This compound is ob- tained by boiling amnionic phosphomolybdate with aqua-regia, and allowing the solution to evaporate spontaneously. From this solution it crystallizes in yellow triclinic prisms with 20 aq., from pure water in cubes with 50 aq., and from concentrated nitric acid in rhombic crys- tals with 40 aq. Ammonia phosphomolybdate, 2POAmo3,22M003,120H2, is precipi- tated as a yellow crystalline powder when orthophosphoric acid or a soluble orthophosphate is added to an excess of a solution of ammonic molybdate in nitric acid. It is insoluble in water and in dilute acids. In solutions containing an excess of phosphoric acid no precipitate is formed. Potassic phosphomolybdate, 2POKo3,22M003,T201T2, is obtained in minute four-sided prisms by boiling the ammonium salt with caustic potash, or by precipitating a potash salt with a solution of phosphomo- lybdic acid. A second series of phosphomolybdates derived from an acid of the formula 2POHo3,-r)Mo03, is obtained by spontaneous evaporation of a solution of the above salts in excess of alkali or ammonia. Thus from an ammoniacal solution of the yel- low precipitate of ammonic phosphomolybdate in ammonia, lustrous prisms of a salt, 2P0Am05,5M003,7 OH3, are deposited. COMPOUNDS OF MOLYBDENUM. 623 COMPOUNDS OF MOLYBDENUM WITH SULPHUR. Molybdous sulphide, MoS"2, occurs native as molybdenite in lead- gray hexagonal crystals, or in masses closely resembling graphite in appearance, with which it was formerly confounded. It is obtained as a lustrous powder when molybdic anhydride is heated in a current of sulphuretted hydrogen : Mo03 + 3SH2 = MoS"2 + S + 30H2. Molybdic Sulphuretted Molybdous Water, anhydride. hydrogen. sulphide. The trisulphide, when heated with exclusion of air, is also converted with evolution of sulphur into the disulphide. When heated in air, molybdous sulphide is oxidized to molybdic anhydride and sulphurous an hydride. Molybdic sulphide (Molybdic sulphanhydride), is precipitated when hydrochloric acid is added to the solution of a molybdate pre- viously saturated with sulphuretted hydrogen. It is a dark-brown powder which dissolves in solutions of alkaline sulphides, forming sulphomolybdates. Potassic sulphomolybdate, MoS//2Ks2, forms pris- matic crystals, which by reflected light appear green with a metallic lustre, and by transmitted light ruby-red. Molybdic persidphide, MoS"4.—When a solution of potassic molyb- date is saturated with sulphuretted hydrogen and then boiled, a mixture of molybdous sulphide with molybdic sulphide is precipitated, and the solution contains potassic persulphornolybdate, which crys- tallizes in small, transparent, red scales. On adding hydrochloric acid to the solution of this salt molybdic persulphide, MoS"4, is precipitated as a reddish-brown powder. General Properties and Reactions of the Compounds of Molybdenum.—The hypomolybdous and molybdous salts are of relatively slight importance. The molybdates and molybdic acid give characteristic reactions with reducing agents. Thus, if metallic zinc be added to a dilute hydrochloric acid solution of a molybdate, the liquid becomes blue, then green, and finally dark-brown. Sulphuretted hydrogen first colors the acid solution blue, and then precipitates molybdic sulphide; but the whole of the molybdenum can be precipitated only by repeated treatment with sulphuretted hydrogen, allowing the solution to stand in a warm place. Potassic ferr ocyanide gives a reddish-brown precipi- tate. The compounds of molybdenum yield, with borax and with microcosmic salt, beads which in the oxidizing flame are colorless or pale yellow; in the reducing flame the borax bead is brown, and the bead of microcosmic salt green. TUNGSTEN, W. Atomic weight = 184. Molecular weight unknown. Sp. gr. 19.129. Atomicity ", IV, and Tl. Evidence of atomicity: Hypotungstous chloride, . . . . W"C12. Tungstous chloride, . . . . WiTCl*. Tungstic hexachloride, . . . . . . . . WTiCJ6. 624 INORGANIC CHEMISTRY. History.—Tungstic acid was first obtained by Scheele from the mineral scheelite in 1781. Occurrence.—Tungsten occurs only in combination, and almost invariably in the form of tungstates. Wolfram is a tungstate of iron and manganese; scheelite is a calcic tungstate, WCXjCao" ; and scheelitine is a plumbic tungstate, W02Pbo". Tungstic anhydride, WOa, occurs as the rare mineral wolfram ochre. Preparation.—Metallic tungsten is prepared by the reduction of the oxides or chlorides in a current of hydrogen. The reduction of the chlorides may also be effected by means of sodium, and that of the oxides by carbon. The metal has not been obtained in the coherent state. Properties.—Tungsten forms a lustrous metallic powder, which, when the reduction has been effected at a white heat, consists of minute quadratic plates. It is unalterable in air at ordinary temperatures, but when heated to redness in air is converted into tungstic anhydride. Nitric acid oxidizes it slowly, aqua-regia rapidly, to tungstic acid. The quality of steel is stated to be improved by the addition of tungsten. COMPOUNDS OF TUNGSTEN WITH THE HALOGENS. Hypotungstous chloride, WC12, is most readily obtained by heating the tetrachloride in a current of carbonic anhydride (see Tungstous chloride). It forms a gray non-volatile powder, which is decomposed by exposure to the air. In contact with water it slowly evolves hydro- gen, and is converted into brown hydrated dioxide, whilst the liquid contains hydrochloric acid. Hypotungstous bromide, WBr2, and hypotungstous iodide, Wl2, have also been prepared. Tungstous chloride, WCI2, is formed during the preparation of the pentachloride from the hexachloride. As it is non-volatile, it remains behind in the process of sublimation. It forms a yellowish-brown infusible crystalline mass. When strongly heated with exclusion of air it splits up into tungstic pentachloride, which volatilizes, and hypo- tungstous chloride, which remains : 3WC14 = 2WC15 + WCI2. Tungstous Tungstic Hypotungstous chloride. pentachloride. chloride. It is hygroscopic, and is decomposed by water into hydrochloric acid and brown hydrated tungstous oxide. Tungstic pentachloride, WC15. Molecular volume 1 1 I.—This compound may be obtained by careful distillation of the hexachloride in a current of hydrogen. It is best, however, to carry the reduction as far as the formation of the tetrachloride, which may be done by employing a higher temperature, and then to decompose the tetrachloride by heating still more strongly in a current of carbonic anhydride, when it breaks up into pentachloride and dichloride (see Tungstous chloride). It forms black lustrous needles, fusing at 248° C. (478° F.) and boiling COMPOUNDS OF TUNGSTEN. 625 at 275.6° C. (528° F,). The vapor is yellowish-green. (As regards the anomalous molecular weight of this compound, as deduced from the vapor density, see p. 179, footnote.) It is very hygroscopic, and is decomposed by water with separation of a blue compound supposed to be a tungstous tungstate. Tungstic itexachloiude, WC16. Molecular volume I I L—When tungsten is heated in a current of chlorine, combination occurs with incandescence, and the hexachloride is formed. The metal employed must be perfectly free from oxide,and the chlorine must contain neither air nor moisture, otherwise the product will be contaminated with oxytetrachloride, WOCl+. It forms a violet-black crystalline mass, fusing at 275° C. (527° F.) and boiling at 346.7° C, (654° F.). In the* neighborhood of its boiling-point the vapor possesses a density corre- sponding with the formula WCI6; at higher temperatures the density is less, owing to dissociation. Pure tungstic hexachloride is not altered by exposure to air, but when it contains oxychloride it undergoes decomposition, evolving fumes of hydrochloric acid. In like manner the pure hexachloride is not decomposed by water until heated with it,- but that which contains oxychloride is at once decomposed in the cold with formation of a greenish oxide. It is soluble in carbonic disulphide, yielding a reddish-brown solution, from which it is deposited in brown six-sided plates. Tungstous oxide, W02, is obtained when tungstic anhydride is heated to low redness in a current of hydrogen. Too high a temperature must be avoided, as otherwise metallic tungsten will be formed. On the other hand, if too low a temperature be employed, tungstous tungstate, W02(02W1V0)", is obtained as a blue powder. Tungstous oxide is a brown powder, which is scarcely attacked by acids. When freshly prepared it is pyrophoric, and must be allowed to remain for some time in an atmosphere of hydrogen before exposing it to the air. Tungstic anhydride, WO,, occurs native as the rare mineral wolfram ochre. It is best obtained from wolfram, a tungstate of manganese and- iron. The finely powdered mineral is digested for several days on the water-bath with hydrochloric acid, and finally with aqua-regia. The insoluble portion, which consists of tungstic acid along with unattacked wolfram and gangue, is washed with water and extracted with ammonia, which dissolves the tungstic acid. The am- m’onic tungstate is converted into the ;fnhydride by ignition.—Tungstic anhydride is a yellow powder, which is fusible at a very high tempera- ture, and may be volatilized at a white heat. Exposure to light colors it green. It may be obtained in greenish crystals by fusion with borax, or by igniting a mixture of sodic carbonate and sodic tungstate in a current of gaseous hydrochloric acid. It dissolves in caustic soda and caustic potash, but is insoluble in ammonia and in acids. Tungstic acid.—This acid exists in several modifications. When an acid is added to a cold solution of a tungstate, a white precipitate is obtained, which when dried by exposure to air possesses the composition COMPOUNDS OF TUNGSTEN WITH OXYGEN 626 INORGANIC CHEMISTRY. WOHo4. When this compound is dried over sulphuric acid it parts with water, and is converted without change of color into the dibasic acid, W02Ho2. The latter compound may also be obtained as a yellow precipitate by pouring the hot solution of a tungstate into hot nitric acid, or by boiling an insoluble tungstate with nitric acid. These acids are insoluble in water. In contact with zinc and hydrochloric acid, tungstic acid is colored first blue and afterwards brown, owing to the formation of tungstous tungstate and of a lower oxide—probably the hydrated dioxide.—A soluble metatungstic acid, W4OnHo2,70H2, is obtained by decomposing baric raetatungstate (see Tungstates) with sulphuric acid, or plumbic metatungstate with sulphuretted hydrogen, and evaporating at ordinary temperature. It forms soluble yellow octahedra. The solution has an acid reaction, and may be concentrated to a syrup, but on boiling the concentrated solution a separation of ordinary insoluble tungstic acid occurs.—A second soluble modification, colloidal tungstic acid, is obtained like the corresponding modification of molybdic acid (p. 621) by adding to a 5 per cent, solution of sodic tungstate a quantity of hydrochloric acid sufficient to combine with the sodium, and subjecting the liquid to dialysis. The solution may be boiled either alone or with acids without depositing ordinary tungstic acid. The colloidal acid may be obtained by evaporation as a vitreous mass, which may be heated to 20J° C. (392° F.) without being converted into the insoluble modification. The vitreous acid dissolves slowly but completely in one-fourth of its weight of water. When strongly heated, all the modifications of tungstic acid are converted into the anhydride. As in the case of molybdenum, oxy-halogen compounds of tungsten have been prepared : Tungstic oxytetrachloride, . . . WOCl4. Tungstic dioxydichloride, . . W02C12. Tungstic dioxydibromide, . . . . . W02Br2. THE TUNGSTATES. Tungstic acid forms a series of very complex salts.. These resemble in many respects the salts of molybdic acid, especially in the case of the poly tungstates, which correspond with the polymolybdates, and are formed by the combination of the normal salt with the anhydride in varying proportions. The complexity is further increased by the exist- ence of a separate class of saljts, the metatungstates, which are distin- guished by not yielding a precipitate on the addition of an acid, except after prolonged boiling. Potassic tungstates.—The normal salt is obtained by adding tungstic anhydride in small quantities at a time to fused potassic carbonate, dis- solving the cooled mass in hot water. The solution deposits on cooling prismatic crystals of the formula WO2Ko22,0H2. When a solution of the normal salt is boiled with tungstic anhydride as long as the latter dissolves, a duodecatung.state of the formula Wl2031Ko1(), 1101I2 separates in lustrous scales as the liquid cools. Sodic tungstates.—The normal salt, W02Nao2,20H2, is obtained like THE TUNGSTATES. 627 the potassium salt, and crystallizes in thin rhombic prisms. The following is a list of the sodic tungstates which have been prepared: Disodic ditungstate, . . . W02Nao220H2. Disodic ditungstate, . . . W2O5Nao2,20H2. Tetrasodic tritungstate, . . W3O7Nao4,70H2. Tetrasodic pentatungstate, . W5Ol3Nao4,110112. Hexasodic heptatungstate, . W7O18Nao6160H2, or 210H2. Decasodic dodecatungstate, . Wl2O31Nao10,210H2, or 250H2,0r 280H2. The dodecatungstate, also known as sodic paratungstate, is manufac- tured by roasting the mineral wolfram with soda ash and extracting the fused mass with water. The solution is almost neutralized with hydro- chloric acid and then left to crystallize. At ordinary temperatures the aquate with 28 aq. is deposited in large triclinic crystals ; at higher temperatures the other aquates given in the above list are obtained. This salt is sometimes employed as a mordant, and also in rendering cotton and linen fabrics uninflammable.—Sodic metatungstate, W4OnNao2,- 100H2, is obtained by boiling normal sodic tungstate with tungstic anhy- dride. It crystallizes in efflorescent octahedra, which are soluble in less than one-tenth of their weight of cold water. Ammonia tungstates.—The normal salt has not been prepared, but various polytung- states and a metatungstate are known. The following tungstates occur as minerals : Calcic tungstate (scheelite1, . . . . . . . W02Cao//. Plumbic tungstate (slolzite), . . . . Ferrous tungstate (farberite), . . . . . . WOzFeo" Manganous tungstate (hiibnerite) . . . , . . WOsMno". An isomorphous mixture of the last two compounds constitutes the mineral wolfram. A class ofphospho-tungstates is known, corresponding with the phospho- molybdates. Silica tungstic Acids.—Some of the polytungstic acids combine with silicic acid to form peculiar compound acids. When sodic or potassic heptatungstate is boiled with gelatinous silicic acid, salts of silico-dodecatungstic acid, SiWj2034[1o8, are formed. In order to obtain the free acid, mercurous nitrate is added to the solution of the salts, and the precipitate of mercurous silicotungstate, after washing, is decomposed with hydro- chloric acid. The filtrate from the mercurous chloride deposits on spontaneous evapo- ration large, colorless, lustrous, quadratic octahedra of the above acid with 29 aq. When heated it fuses in its water of crystallization and deposits at a temperature of 53° C. rhombohedra containing 22 aq. It forms both normal and acid salts : thus the three potassic silicotungstates have had the following formulae assigned to them : SiW120,4Ko8.14OH2, Si W12039 Ho4 Ko4, 150 H2i 2SiW12034Ho5Ko3,250H2. If gelatinous silicic acid be boiled with an amnionic polytungstate, the ammonium salts of two other silicotungstic acids are formed: of a silicodecatungstic acid, Si WIOO.2SHo8, 30112, and of a silico-dodecatungstic acid isomeric with that above described. This second dodeca-acid is known as tungsto-silicic acid. It crystallizes in triclinic prisms with 20 aq., and its salts are distinguished from those of ordinary silico-tungstic acid by greater solubility, by crystallizing in different forms, and by containing a different number of molecules of water of crystallization. The Tungsto-txjngstates. These compounds, which may be regarded as combinations of the tungstates with 628 INORGANIC CHEMISTRY. tungstens oxide, are obtained by the reduction of the polytungstates. They are charac- terized by metallic lustre, and have been employed as bronze powders, f W02Ko O Potnssic tungsto-tungstate, -{ WO .—Tungstic anhydride is added to fused potassic O _ WOjKo tungstate as long as it dissolves. The mass thus obtained is then reduced by gently heating in a current of hydrogen, and is then extracted successively with water, hydro- chloric acid, caustic potash, and again with water. It is thus obtained in dark-blue needles, with a coppery lustre. Sodic tungsto-tungstate, W307Nao2, may be obtained either by a method similar to the above, or by fusing a polytungstate of sodium with tin, and extracting the mass with caustic soda and hydrochloric acid. It crystallizes in golden cubes, with a fine metallic lustre. COMPOUNDS OF TUNGSTEN WITH SULPHUR. Tungstous sulphide, WS2", is formed when the trisulphide is heated with exclusion of air, or when tungstic anhydride is heated in a current of sulphuretted hydrogen : W03 + 3SH2 = WS"2 + S + 30H2. Tungstic Sulphuretted Tungstous Water, anhydride. hydrogen. sulphide. It forms a blue-black crystalline powder. Tungstic sulphide (Tungstic sulphanhydride), WS"3, is obtained like the corresponding molybdenum compound by saturating the solution of a tungstate with sulphuretted hydrogen and then adding an acid. It is a dark-brown powder, which dissolves in alkaline sulphides with form- ation of sulpho-tungstates. Potassic sulphotungstate, WS"2Ks2, forms yellow prismatic crystals. General Properties and Reactions of the Compounds of Tungsten.—The insoluble compounds of tungsten can be obtained in a soluble form as alkaline tungstates by fusion with an alkali, prefer- ably with the addition of nitre. When metallic zinc or tin is introduced into the hydrochloric acid solution of a tungstate, the liquid assumes a deep-blue color. Ammonic sulphide produces in the solution of a tung- state no precipitate, but if hydrochloric acid be added to the liquid thus obtained, tungstic sulphide is precipitated as a dark-brown powder. The tungsten compounds yield with microcosmic salt a bead which, in the oxidizing flame, is colorless or pale yellow, in the reducing flame blue. On the addition of ferrous sulphate, the bead assumes a blood- red color in the reducing flame. CHROMIUM, 629 CHAPTER XLI. HEX AD ELEMENTS. Section 111. CHROMIUM, Cr. Atomic weight = 52. Molecular weight unknown. Sp. gr. 7.3 (Bunsen). Atomicity ", 1V,.V1, and possibly viii; also a pseudo-triad. Evidence of atomicity: Chromous chloride, Chromic chloride, .... , Cr"Cl2. / Cr'^Clg \ Cr'"Cls- Chromic perfluoride, . . . Chromic anhydride, . . . Perchromic acid, .... . CrTiF6. . CrviOs. / Crvii03Ho(?). \ 0rviiO3Ho History.—Chromium was discovered by Van quel in in 1797, and in- dependently by Klaproth about the same time. Occurrence.—Chromium does not occur abundantly, and is never found in the free state. Its chief natural compounds are those which it forms with other metals, together with oxygen. Of these the most abundant is chrome iron ore, ,Cr202Feo//. It also occurs as plumbic chromate, Cr02Pbor/, crocoisite. The color of various minerals and gems, such as serpentine, chromic mica, and the emerald, is due to the presence of small quantities of chromium. Preparation.—Chromium may be reduced from its chloride by means of zinc. For this purpose the chloride is heated with zinc in a Hessian crucible, employing a mixture of potassic chloride and sodic chloride as a flux. The zinc regulus is treated first with cold and afterwards with warm dilute nitric acid, as long as anything dissolves. The metallic chromium remains as a gray powder. For the above reaction, it is not necessary to prepare anhydrous chromic chloride : the mixture of chro- mic chloride and potassic chloride obtained by the reduction of potassic dichromate with hydrochloric acid and alcohol is evaporated with the addition of sodic chloride, and the mass thus obtained is carefully dried and employed as above.—Chromium may also be obtained by heating chromic oxide to a very high temperature with sugar in a lime crucible. —Bunsen has prepared the metal by the electrolysis of a solution of chromous chloride containing chromic chloride. Properties.—Metallic chromium, reduced from the chloride by zinc, is a light-gray crystalline powder, in which, by the aid of the micro- scope, tin-white lustrous octahedra may be perceived. Prepared by elec- trolysis, it is deposited on a platinum electrode as a coherent plate. It is more difficultly fusible than platinum, and as hard as corundum. It is only slowly oxidized when heated in air with a Bunsen or hydrogen flame, but burns with brilliant scintillations in the oxy-hydrogen flame. 630 INORGANIC CHEMISTRY. When thrown on potassic chlorate which has been heated to incipient fusion, it is oxidized with dazzling incandescence, yielding potassic chro- mate. Hydrochloric acid dissolves it readily with evolution of hydro- gen ; dilute sulphuric acid scarcely attacks it in the cold, but when hot dissolves it, also evolving hydrogen ; nitric acid, even when hot and concentrated, does not act upon it. The hardness of steel is greatly increased by the addition of 0.5 to 0.75 percent, of chromium. COMPOUNDS OF CHROMIUM WITH THE HALOGENS. a. Chromous Compounds. Chromous chlorides, CrCl2.—A solution of this compound is obtained when the metal is dissolved in hydrochloric acid. The anhydrous chlo- ride is best prepared by gently heating chromic chloride in a current of dry hydrogen. It forms a white crystalline mass, and dissolves in water, yielding a blue solution, which rapidly absorbs oxygen from the air and possesses powerful reducing properties. Chromous bromide, Crßr2, is prepared in a similar manner from chromic bromide. It resembles the chloride in its properties. 6. Chromic Compounds. Chromic chloride, /Ci’2Cl6, is prepared by heating a mixture of chromic oxide and carbon in a current of dry chlorine. It forms lus- trous scales, of the color of peach-blossom, which may be sublimed in a current of chlorine. When heated in air, it evolves chlorine and is converted into chromic oxide. Pure chromic chloride is almost insolu- ble in water at ordinary temperatures, and dissolves only slowly when boiled with water for a considerable time, but in presence of a very- minute quantity of chromous chloride, it dissolves readily in cold water, yielding a green liquid. Stannous chloride and other reducing agents produce the same effect. The green solution, which may also be obtained by dissolving chromic hydrate in hydrochloric acid, yields, by evapora- tion over sulphuric acid, green, very soluble needles of the compound These, when heated, part with water and hydrochloric acid, and are converted into an oxychloride. By heating in a current of gaseous hydrochloric acid, they may, however, be converted into the anhydrous violet chloride. Chromic bromide, /Cr2Br6, is prepared like the chloride. It forms black hexagonal scales, with a submetaliic lustre. The crystals exhibit, by transmitted light, olive-green and red dichroisra. Chromic fluoride, /Cr2F6, is obtained by dissolving chromic hydrate in hydrofluoric acid. On evapol'hting the solution a green crystalline mass is obtained, which fuses at a red heat, and at a very high temperature sublimes in lustrous regular octahedra. c. Perchromic Compounds. Only one of these is known—the perjiuoride. In all circumstances where the formation of a perchloride or perbroraide might be expected, chlorine or bromine is evolved, and the corresponding chromic com- pound is formed. 631 COMPOUNDS OF CHROMIUM. Chromic perfluoride, CrF6, is prepared by heating a mixture of calcic fluoride and ignited plumbic chromate with concentrated sulphuric acid in a retort of lead or platinum : 0rO2Pbo" + 3CaF2 + 4S02Ho2 = CrF6 + Plumbic Calcic Sulphuric Chromic chromate. fluoride. acid. fluoride. S02Pbo" + SS02Cao" + 40H2. Plumbic Calcic Water, sulphate. sulphate. A red gas passes over, which condenses to a red fuming liquid. In con- tact with water it is decomposed, yielding chromic and hydrofluoric acids. COMPOUNDS OF CHROMIUM WITH OXYGEN. o r f!rO /n Chromic oxide, O=Cr—Cr=o. O Chromic anhydride, . . CrOs. O—Cr=o. Chromous oxide, CrO, is not known, but its hydrate and several of its salts have been prepared.— Chromous hydrate, CrIIo2, is obtained as a brownish-yellow precipitate by the addition of caustic potash to a so- lution of chromous chloride. It readily absorbs oxygen, and must be dried with exclusion of air. When heated in absence of air, it parts with water and hydrogen, being converted into chromic oxide : 2CrHo2 = Cr203 + 0H2 + H2. Chromous Chromic Water, hydrate. oxide. Chromic oxide, 'CrO^.—This oxide occurs native as the mineral chrome-ochre, contaminated with earthy impurities. It is formed when chromic hydrate, chromic anhydride, or diammonic dichromate, is heated : fCrO2(NII40) fCQ o = \cloO + N2 + 40H, (OrO2(NH4O) Diammonic Chromic Water, dichromate. oxide. It is most readily obtained by heating a mixture of equal parts of dipo- tassic dichromate and sulphur, or ofdipotassic dichromate and ammonic chloride. On extracting the residue with water, the chromic oxide re- mains undissolved. It is a dark-green amorphous powder, which fuses in the oxy-hydrogen flame, and solidifies to an almost black, crystalline mass. It may be obtained in lustrous, dark-green, hexagonal crystals by passing the vapor of chromic oxydichloride, Cr02Cl2, through a red- hot tube. The strongly ignited oxide is almost insoluble in acids. 632 INORGANIC CHEMISTRY. Chromic oxide is used as a pigment under the name of chrome green, and in the preparation of green glass and enamel. Chromic hydrate, 'Ci^Hog.-—Ammonia produces in solutions of chromic salts free from alkali a pale-blue precipitate of a hydrate which, after drying over sulphuric acid, has the formula a vacuum it slowly parts with 3 aq., and when heated to 220° C. in a current of hydrogen is converted into the hydrate 'Gr20211o2. Another hydrate of the formula /Cr2OHo4, employed as a pigment under the name of Guignefs green, is prepared by fusing dipotassic dichromate with boric acid, and extracting the mass with water. These hydrates are difficultly soluble in acids. Freshly precipitated chromic hydrate dissolves slightly in aqueous ammonia,yielding a peach-blossom-colored solution. This solubility depends upon the formation of a chromamine corresponding with the cobaltamines (q.v.). The freshly precipitated hydrate also dissolves in a solution of chromic chloride, and from the solution thus obtained the greater part of the hydrochloric acid may be removed by dialysis, leaving a soluble colloidal modification of chromic hydrate. (Graham found in the liquid remaining in thedialyser 1 mol. of hydrochloric acid to 33 mols. of chromic hydrate). The dark-green solution is not precipitated by dilution or by boiling, but the addition of the slightest trace of a salt causes it to coagulate.—The precipitate produced in solutions of chromic salts by caustic alkalies, which dis- solves in an excess of the precipitant, and is reprecipitated by boiling, always contains alkali; and this cannot be removed by washing. Chromic anhydride, 0rO3 —ln order to prepare this compound, T| volumes of concentrated sulphuric acid are added to one volume of a cold saturated solution of dipotassic dichromate. On cooling, the chromic anhydride crystallizes out in long red needles. It may be freed from the excess of acid by allowing it to drain upon a porous tile, in which condition it is sufficiently pure for most purposes. In order to obtain it quite pure, the crystals must be filtered off, employing a filter of asbestos or spun glass, as organic substances instantly reduce the anhydride, and the substance must be washed upon the filter with pure nitric acid free from nitrous anhydride, and finally freed from the nitric acid by warming in a current of dry air.—Chromic anhydride forms either a woolly mass of fine red needles, or red prisms. It is very soluble in water, yielding a reddish-brown solution, which becomes yellow on dilution. It is also soluble both in concentrated and in dilute sulphuric acid, but it is almost insoluble in a sulphuric acid containing from 16 to 17 per cent, of water—a property which is utilized in its preparation. It fuses without decomposition when heated, but at 250° C. (482° F.) is resolved into chromic oxide and oxygen. It is very readily re- duced to chromic oxide, and therefore acts- as a powerful oxidizing agent. Sulphurous anhydride, sulphuretted hydrogen, nitrous anhy- dride, and most organic substances effect its reduction. Alcohol poured upon the dry anhydride inflames. Glacial acetic acid, however, dissolves it without decomposition. Both the aqueous and the acetic acid solutions of chromic anhydride are employed in organic chemistry as oxidizing agents, the latter solution being particularly efficacious, owing to the fact that the acetic acid generally also acts as a solvent for the organic substance which is to be oxidized. Sometimes, instead of OXY-SALTS OF CHROMIUM. 633 aqueous chromic anhydride, a solution of dipotassic dichromate in dilute sulphuric acid is employed as an oxidizing agent. When heated with hydrochloric acid, chromic anhydride evolves chlorine, and is converted into chromic chloride; heated with concentrated sulphuric acid it gives off oxygen, yielding chromic sulphate. Chromic acid, 0rO2Ho2. See Chromates. Perchromic add, | Cr() H()^'^’—When hydroxyl is added to a solution of chromic anhydride or of a chromate acidified with sulphuric acid, a deep-blue coloration is produced. The compound thus formed, which is possibly a perchromic acid of the above composition, speedily decomposes with evolution of oxygen, and the solution contains only chromic acid. On agitating the blue solution with ether, this solvent extracts the blue compound from the water, and rises to the surface as a dark-blue layer. The ethereal solution, though somewhat more stable than the aqueous solution, leaves only chromic anhydride on evapora- tion. The formation of this blue compound is a very delicate and char- acteristic test, both for chromic anhydride and for hydroxyl—indeed, for the latter substance it is the only thoroughly characteristic test. The other oxides of chromium generally enumerated are difficult to obtain of con- stant composition. A chromous dichromic tetroxide, /Cr2o/Cro//, is probably formed in the process of preparing the metal by electrolysis, but appears to be mixed with metallic chromium. The substance known as chromic dioxide, Cr02, is probably a compound of chromic anhydride with chromic oxide ; by washing with water it is decomposed into these two substances. OXY-SALTS OF CHROMIUM. a. Chromous Salts. The chromous salts are of slight importance. They are readily oxidizable, and absorb oxygen from the air. Chromous sulphate, SOjCro77, is known only in solution. It is formed when metallic chromium is dissolved in dilute sulphuric acid. Dipotassic chromous sul- I S02Ko phate, 1 Cro// , (iOH2, is prepared by dissolving potassic sulphate in a solution of ( SOJCo chromous chloride, adding alcohol, and then allowing the mixture to stand for some time with exclusion ol air. It crystallizes in blue monoclinic prisms, which on expo- sure to air quickly become green from oxidation. Chromous phosphate, F202Cro//3, and chromous carbonate, COCro// have also been prepared. b. Chromous Salts. Chromic oxide forms with acids the chromic salts, in which the hexadic group (/Cr///2)vi displaces six atoms of hydrogen in the acid. The aqueous solutions, prepared by dissolving the salts in cold water, are violet colored ; on heating, the color changes to green, and on cool- ing, the violet color returns only after a considerable time. Crystal- lized salts can be obtained only from the violet solutions : the green solutions yield, by evaporation or on the addition of alcohol, green amorphous masses. The violet solutions alone contain a pure chromic salt; this, on warming, is decomposed into basic salt and free acid, the chemical change being accompanied by the above change of color. 634 INORGAXIC CHEMISTRY. Chromic nitrate, N6012(/Cr///206)r1,180H2, is prepared by dissolving chromic hydrate in nitric acid. It forms red, soluble, monoclinic crystals. Chromic Sulphate, S3O6(/Cr//'2O6)vi,150H2, is prepared by dissolv- ing chromic hydrate in its own weight of concentrated sulphuric acid. The solution, which is green at first, becomes blue on standing, and deposits a violet-blue crystalline mass of the above salt. This may be purified by dissolving in cold water and precipitating with alcohol. From its solution in cold dilute alcohol it is deposited in blue regular octahedra. The aqueous solution prepared in the cold has a violet color, which changes to green on boiling. Dipotassic Chromic Sulphate (Chrome alum), SO,Ko-, SO I gQ2 ('Cr'//206)V1,2-10H2.—This compound is best prepared by dis- so22K0-1 solving equal molecular proportions of dipotassic dichromate and sul- phuric acid in water and passing sulphurous anhydride into the solu- tion : C Cr02Ko |'),aKol - O + S03Ho2 + 3S02 = s ?“('Cr'"A)" + 0H2. S§K„J Dipotassic Sulphuric Sulphurous Chrome Water* dichromate. acid. anhydride. alum. Other reducing agents, such as alcohol, may be employed instead of sulphurous anhydride, but in this case it is necessary to add a larger quantity of sulphuric acid. Chrome alum crystallizes in deep ruby-red octahedra, which by reflected light appear almost black. It dissolves in cold water with a reddish-violet color, which becomes green on boil- ing. After standing for a long time it recovers its original color. Chrome alum is employed in dyeing and calico-printing, and in tanning. —Ammonia chrome alum is prepared like the potassium compound, which it closely resembles in its properties. THE CHROMITES. Chromic oxide possesses the property of combining with other oxides—especially with the oxides of the dyad metals—to form com- pounds which may be regarded as salts of the acid /Cr202H02. To this particular hydrate of chromium the name chromous acid may therefore be applied, and these compounds would then be termed chromites. It lias already been mentioned that when chromic hydrate is precipitated by caustic alkalies, the precipitate contains alkali which cannot be removed by washing. This is due to the formation of a chromite of the alkali. Only the chromites of the dyad metals, however, have been obtained as well characterized compounds. These crystallize in regular octahedra, and belong to the same class as the alurninates of the dyad metals (p. 568), or as magnetic oxide of iron (q.v.), all of which also crystallize in regular octahedra, and may be regarded as formed by the combina- tion of a monoxide with a sesquioxide. THE CHROMATES. 635 Zlncic chromite/Cr,o.2Zno//, is obtained in lustrous blackisb-green octahedra by fusing a mixture of zincic oxide and chromic oxide with boric anhydride. Manganous-chromite, /Cr202Mno//, is obtained in a similar manner, substituting man- ganous oxide for zincic oxide. It forms very hard iron-gray octahedra. Ferrous chromite, '0r2O2Feo//, occurs in nature as the mineral chrome iron ore. It generally occurs in crystalline masses; but distinct octahedral crystals are also found. THE CHROMATES. These are the salts of the unknown chromic acid, Cr02Ho2. This acid possibly exists in the aqueous solution of chromic anhydride, but on evaporating this solution only chromic anhydride is obtained. Hydroxyl does not appear to be capable of entering into stable combi- nation with the radical chromyl (Cro2)//. Not only does chromic acid part spontaneously with the elements of water to yield an anhydride ; but not even the acid chromates arecapable of existing. Thus in all cases in which the formation of hydric potassic chromate might be expected, two molecules of this salt combine, with elimination of one molecule of water, ( 0rO2Ko an hydro-salt dipotassic dichromate, < O \cro2Ko When chromic oxide, a chromic salt, or any substance containing chromium is fused with nitre, the chromium is oxidized by the oxygen of the nitre, and a yellow mass is obtained which contains potassic chromate, Cr02Ko2. Formerly this compound was prepared by heat- ing chrome iron ore with nitre, but at the present day potashes are sub- stituted for the more costly nitre, and the oxidation is effected by means of the oxygen of the air. Chrome iron ore is first roasted and then finely ground. A mixture of roasted and powdered ore, crude pot- ashes, and lime is first dried at 150° C. (302° F.) and then heated to bright redness in the oxidizing flame of a reverberatory furnace. The addition of the lime prevents the fusion of the mass, which is thus kept in a pasty condition. During the operation the heated mass is constantly stirred, so as to expose fresh surfaces to the oxidizing action of the flame. As soon as the oxidation is complete the charge is withdrawn, and, after cooling, is lixiviated with the smallest possible quantity of boiling water. If the solution should contain calcic chromate, potassic sulphate is added in quantity sufficient to convert it into the potassium salt, the calcium being at the same time precipitated as sulphate. The solution now contains potassic chromate, but it would be impossible to separate this salt by crystallization from the other salts present, owing to its ready solubility. It is therefore necessary to convert it into the less soluble dichromate. For this purpose a quantity of sulphuric acid sufficient to saturate half the potassium present as chromate is diluted with twice its volume of water and added to the solution : f Cr02Ko 20rO2Ko2 4 S02Ho9 = { O + S02Ko2 + 0H2. [ 0rO2Ko Potassic Sulphuric Dipotassic Potassic Water, chromate. acid. dichromate. sulphate. 636 INORGANIC CHEMISTRY. The normal chromate is soluble in twice its weight of cold water, whilst the dichromate requires ten times its weight of water for solution; the greater part of the dichromate therefore crystallizes from the above liquid on cooling. The mother liquor, which contains potassic sul- phate, is employed in the extraction of another roasted charge. The potassic dichromate is purified by crystallization. (For the properties of this salt see below.) Potassic chromate, Cr(),Ko2.—(For the mode of formation, see preceding paragraph.) In order to obtain this salt in a state of purity an excess of caustic potash is added to a solution of the dichromate. The color of the solution changes from orange-red to yellow, and on evaporation yellow rhombic crystals of the normal chromate are deposited. The crystals are isomorphous with those of potassic sulphate, with which salt it is capable of crystallizing in all propor- tions. It is soluble in twice its weight of cold water, yielding a yellow solution. It has an alkaline reaction. The pure salt undergoes decomposition when its solutions are evaporated: crystals of the di- chromate are first deposited; afterwards when the solution begins to contain more free alkali, the normal salt crystallizes out. Acids, even carbonic, decompose it with formation of dichromate. On heating, it turns red and fuses at a high temperature without decomposition, recovering its original color on cooling. f Cr02Ko Dipotassic dichromate, < O .—(For the mode of preparation, ( Cr02Ko see p. 635.) This salt crystallizes in large garnet-red triclinic prisms or tabular crystals. It is soluble in 10 parts of water at ordinary temperatures, more readily soluble in boiling water. The solution has an acid reaction. The salt fuses below a red heat without decom- position, but at a white heat is decomposed into normal chromate, chromic oxide and oxygen. When heated with concentrated sulphuric acid it evolves oxygen and yields a green solution which, after dilution with water, deposits on standing crystals of chrome alum. It is a vio- lent poison.—Dipotassic dichromate is the starting point in the prepa- ration of the other chromium compounds. It is employed as a laboratory reagent, as an oxidizing agent, in dyeing and calico-printing, and in the process of producing permanent carbon photographs. f CrG2Ko 1 ° Dipotassic trichromate, -{ 0rO2 , and Dipotassic tetrachromaie, O ' CrG2Ko Cr4OuKo2, are deposited from solutions of the foregoing salt in nitric acid. They form deep-red crystals, which are decomposed by water into dichromate and chromic anhydride. Sodic chromate, CrCTNao,, is obtained when a solution of potassic chromate with an excess of caustic soda is evaporated at o°. It crystallizes at a low temperature in large yellow transparent deliquescent prisms of the formula CrO2Noa2,100H2, isomorphous with crystallized sodic sulphate, from warm solutions in anhydrous crystals.—Disodic dichromate, Cr200Na02,20H2, forms deliquescent red prisms. THE CHROMATES. 637 Ammonia chromate, 0rO2(Nvir.G)2, and diammonia dichromate, 0r2O5(NvH2O)2, are obtained by adding the requisite quantity of am- monia to an aqueous solution of chromic anhydride. They resemble in almost every respect the corresponding potash salts. When heated they are decomposed into nitrogen, water, and chromic oxide—the normal salt also evolving ammonia. In the case of the dichromate, this decomposition is attended with incandescence, and the chromic oxide swells up to a bulky mass resembling green tea in appearance. Baric chromate, CrOgßao", is obtained as a yellow crystalline precipitate when the solution of a chromate or dichromate is added to the solution of a barium salt. It is insoluble in water and in acetic acid, soluble in hydrochloric and in nitric acid. It also dissolves in a hot aqueous solution of chromic anhydride, and the liquid deposits on cooling red crystals of baric dichromate, Cr2O5Bao",20II2. These are decomposed by water into chromic anhydride and normal chromate.— Baric chromate constitutes the pigment yellow ultramarine. Strontie chromate, Cr02Sro//, closely resembles the barium salt, but is much more readily soluble in water and in acetic acid. Calcic chromate, Cr02Cao//,20 II.2, is obtained in large yellow prismatic crystals by digesting marble with a solution of chromic anhydride and evaporating the liquid over sulphuric acid. Magnesic chromate, Cr02Mgo//,701T2, forms soluble lemon-yellow rhombic crystals, and is isomorphous with magnesic sulphate. C Cr()2Ko Dipoiassic magnesic chromate, 1 Mgo// ,20H2, is deposited in yellow tabular crystals _ ( Cr02Ko when a solution of dipotassic dichroraate is neutralized with magnesia and then evapo- f Cr02{NvH40) rated. Diammonic magnesic chromate, 1 Mgo ,60H2, is isomorphous with diam- ( Cr02(NvH40) xnonic magnesic sulphate (p. 511). Zincic chromates.—The normal salt is not known, but various basic chromates of zinc have been prepared. Dizincic chromate dihydrate, Cr02(OZn//Ho)2,0H2. is obtained as a yellow precipitate when normal potassic chromate is added to a solution of an excess of zincic sulphate. Plumbic chromate, Cr02Pbo//, occurs native as crocoisife in red monoclinic crystals. The same substance is obtained as a bright yellow precipitate when potassic chromate or dichroraate is added to the solu- tion of a lead salt. This precipitate is employed as a pigment under the name of chrome yellow. It is insoluble in water and acetic acid, but soluble in nitric acid and in caustic potash. When heated it fuses without decomposition, and solidifies to a crystalline mass. Organic compounds, when heated with it, undergo complete oxidation: it is therefore employed in the ultimate analysis of such compounds, particu- larly of those which contain sulphur and chlorine or the metals of the alkalies and alkaline earths.—Chrome yellow is employed in calico- printing. The cloth is first mordanted with the solution of a lead salt. On afterward immersing it in a solution of potassic chromate, the chrome yellow is developed on the fibre of the mordanted parts.— Diplumbic chromate, OrOPbo//2, a basic salt, constitutes the chrome red of commerce. It is formed as a red powder by boiling chrome yellow with normal potassic chromate, or by digesting it with cold caustic soda. It is also obtained as a vermilion-colored crystalline powder by 638 INORGANIC CHEMISTRY. fusing chrome yellow with nitre. Chrome orange is a mixture of chrome red and chrome yellow. It is prepared by precipitating the solution of a lead salt with a weak alkaline solution of potassic chro- mate. Argentic chromate, Cr02Ago2, is formed as a red crystalline pre- cipitate when a dilute solution of normal potassic chromate is added to a concentrated solution of argentic nitrate. It may be obtained in dark-green crystals by boiling diargentic dichroraate with water, or by allowing a solution of the dichromate in ammonia to evaporate. The green crystals yield a red powder. It is insoluble in water, but dissolves in nitric acid, in ammonia, and in solutions of the alkaline chromates.—Diargentic dichromate, Cr305Ago2, is obtained as a scarlet precipitate when a solution of potassic dichromate is gradually added to a solution of argentic nitrate. When hot dilute solutions are employed the salt gradually separates in red triclinic crystals. Mercuric chromate, CrOzHgo'A—The normal salt is obtained in garnet-red rhombic prisms by boiling mercuric oxide with a solution of chromic anhydride. Excess of water decomposes it with separation of the red amorphous basic salt, trimercuric chromate, CrTlgo//3,—a decomposition which exactly corresponds with that which occurs when normal mercuric sulphate is treated with water (p. 636). Potassic chro- mate produces in solutions of mercuric and mercurous salts precipitates of basic chro- mates of mercury. COMPOUNDS OF CHROMIUM WITH OXYGEN AND CHLORINE. Chromic oxydichloride f Chromylic chloride, Cr02Cl2), Molecular volume I I I.—This compound may be theoretically derived from chromic acid by the substitution of chlorine for hydroxyl. It may therefore be regarded as the chloride of the acid radical chrorayl (Cro2)", and bears the same relation to chromic acid that sulphurylic chloride, S02CI2, does to sulphuric acid. In order to prepare this compound, a fused mixture of 10 parts of common salt and 12 parts of dipotassic dichromate is broken into small pieces and introduced into a retort, after which 30 parts of faintly fuming sulphuric acid are intro- duced. The reaction begins of its own accord. The dark reddish- brown vapors are condensed in a cooled receiver. In order to free the product from dissolved chlorine, it must be repeatedly rectified in a current of dry carbonic anhydride. The same compound is formed when a dry mixture of chromic anhydride and ferric chloride is dis- tilled.—Chromic oxydichloride is a mobile liquid, which appears almost black by reflected light, but exhibits a blood-red color by transmitted light. It boils at 128° C. (244° F.). It possesses a specific gravity of 1.92 at 25° C. (77° F.). In contact with moist air it fumes strongly, and when dropped into water is decomposed with violent ebullition, yielding chromic and hydrochloric acids. It has a most energetic action upon oxidizable substances: thus it acts upon phosphorus with explosive violence, whilst sulphur, sulphuretted hydrogen, ammonia, and many organic bodies, such as alcohol, inflame when brought in contact with it. Chromic oxychlorhydrate (Chromylic chlorhydrate, Chlorochromic acid), 0rO2CIHo, a compound corresponding with sulphuric oxychlorhydrate COMPOUNDS OF CHROMIUM. 639 (SO201Ho), has, ]ike chromic acid itself, not been isolated. The non- existence of this compound is a further instance of the inability of the semi molecule of hydroxyl to attach itself to the radical chromyl (see p. 638). Salts of chromic oxychlorhydrate, known as chlorochromates, have, however, been prepared. Potassic chlorochr ornate is obtained by gently warming 3 parts of dipotassic dichromate with 4 parts of con- centrated hydrochloric acid and a little water : f Cr02Ko J O + 2HCI = 20rO2CIKo + OII2. [ 0rO2Ko Di potass! c Hydrochloric Potassic Water, dicliromate. acid. chlorochromate. It crystallizes in large red prisms or tables having a specific gravity of 2.497. An excess of pure water decomposes it into potassic chloride and chromic anhydride; but it may be recrystallized from dilute hydrochloric acid. When heated at 100° C. it evolves chlorine. COMPOUND OF CHROMIUM WITH SULPHUR. ( CrV/ Chromic sulphide, j org//S//, is obtained by the direct union of its elements under the influence of heat. It is also formed when chromic oxide is heated to whiteness in the vapor of carbonic disulphide, or when chromic chloride is heated in a cur- rent of sulphuretted hydrogen.—Chromic sulphide is a gray-black powder with a metallic lustre. It possesses a specific gravity of 3.77. Concentrated nitric acid is without action upon it. When heated in air it is converted into chromic oxide.—Sul- phuretted hydrogen produces no precipitate in solutions of chromic salts, and alkaline sulphides precipitate chromic hydrate with liberation of sulphuretted hydrogen. COMPOUND OF CHROMIUM WITH NITROGEN. < OrN/// Chromic nitride, t is formed by the direct union of its elements when nitro- gen is passed over metallic chromium at a red heat; also by the action of gase- ous ammonia upon heated chromide chloride.—It forms a heavy black powder Which inflames when heated to 200° C. (392° F.) in contact with air. Heated with exclusion of air to a temperature higher than that at which it is formed, it is decomposed into its elements. Chlorine is without action upon it at ordinary temperatures, but when the substance is heated in a current of chlorine it is converted with a series of slight explosions into chromic chloride and free nitrogen. The explosions are due to the formation and immediate decomposition of nitrous chloride. It may be ignited without change in hydrogen and in steam. It is not attacked by hydrochloric or nitric acid, or by aqueous caustic potash. Concentrated sulphuric acid dissolves it, yielding a green liquid which, when diluted with water and allowed to stand, deposits crystals of ammonia chrome alum: S02(N’H40)1 {cJn/// + 4S02Ho2 = |g» (^W. sol(N’H4o>j General Properties and Reactions op the Compounds op Chromium.—a. Chromous compounds..—These are of subordinate im- portance. They are distinguished by their powerful reducing proper- ties. They rapidly absorb oxygen from the air, and are thus converted into chromic compounds. b. Chromic salts.—These are derived from chromic oxide. Their solutions are violet-colored or green; they have a sweetish astringent 640 INORGANIC CHEMISTRY. taste, an acid reaction towards litmus, and are poisonous. Ammonia produces a bulky precipitate of chromic hydrate, which is slightly solu- ble in a large excess of ammonia, yielding a peach-colored solution. Caustic alkalies precipitate green chromic hydrate, soluble in an excess of an alkali in the cold, but precipitated on boiling. Sulphu- retted hydrogen gives no precipitate; ammonia sulphide precipitates chromic hydrate with evolution of sulphuretted hydrogen. When a chromium compound is fused with a mixture of sodic carbonate and nitre, an alkaline chromate is formed which dissolves in water, yielding a yellow solution. c. Chromates.—The soluble chromates yield with lead salts a yellow precipitate of plumbic chromate; with argentic nitrate, red argentic chromate. When heated with concentrated hydrochloric acid they evolve chlorine, and the color of the liquid changes to green. Sulphu- retted hydrogen reduces the chromates in acid solution to chromic salts with separation of sulphur; alcohol and sulphurous acid effect the same reduction. Chromium compounds yield, with borax and with microcosmic salt, beads which are emerald-green, both in the oxidizing and in the reduc- ing flame. Chromium compounds do not color flame, but yield a characteristic spark-spectrum containing bright lines in the green and in the blue. MANGANESE, Mn. Atomic weight 55. Molecular weight unknown. Sp.gr. 7.99. Atom- icity ", lv, Tl, and possibly TUI; also a pseudo-triad and a pseudo- heptad. Evidence of atomicity: Manganous chloride, .... Manganic peroxide, . . . Mn"Cl2. . . . MnivO„. Potassic manganate, .... Potassic permanganate, / Mnvii03Ko ' ' - \ Mnvll03Ko* History.—The black oxide of manganese was known to the ancients, who were acquainted with its use in removing impurities from glass. They confounded it, however, with magnetic oxide of iron. Occurrence.—Manganese is widely distributed in nature. Tt is never found native. The chief ores of. manganese are the oxides, and of these the most important is manganic peroxide or pyrolusite, Mn02. Others are dimanganic trioxide or bramiite, Mn2Os; man- ganous diraanganic tetroxfde or haussmannite, -f It also occurs as manganous sulphide in manganese blende, MnS//, and as manganous carbonate, COM no", in manganese spar. It is present in small quantity in a number of other minerals, particularly silicates, so that in almost all rocks and soils traces of manganese are to be found. It occurs in minute quantities in the bodies of plants and animals. Preparation.—Manganese cannot be reduced from its oxides by means of hydrogen ; but the reduction may be effected by heating the oxide with carbon to intense whiteness. A mixture of 10 parts of COMPOUNDS OP MANGANESE. 641 manganous diraanganic tetroxide, Mn202Mno" (obtained by the igni- tion of the native peroxide), with 1 part of charcoal and 1 part of an- hydrous borax, is heated to whiteness in a carbon crucible. The regu- lus of manganese thus obtained contains carbon, together with silicon derived from the ash of the charcoal. Pure manganese may be obtained by heating manganous manganic oxide (prepared from the artificial dioxide) in a lime crucible with a quantity of carbonized sugar insuffi- cient for its total reduction. The lime crucible is placed inside a Hes- sian crucible, the intervening space is filled with charcoal, and the wffiole is heated in a wind-furnace. Properties.—Manganese is a grayish-white metal with a reddish tinge. It is very hard and brittle. It fuses at a white lieat. It oxidizes rapidly in moist air, and must therefore be preserved under rock-oil. Manganese is rapidly dissolved by dilute acids, and the finely divided metal decomposes water with evolution of hydrogen when gently warmed with it. COMPOUNDS OF MANGANESE WITH THE HALOGENS. Manganous chloeide, MnCl2.—The anhydrous chloride is formed when the metal is burnt in chlorine, or when any of the oxides or the carbonate is heated in a current of dry hydrochloric acid. The residues from the preparation of chlorine by the action of hydrochloric acid upon manganic peroxide may be employed as a source of manganous chloride. This solution contains manganous chloride contaminated with ferric chloride, and sometimes with the chlorides of copper, barium, and calcium, together with an excess of hydrochloric acid. The solution is evaporated to expel the acid, diluted, and about an eighth of the solution precipitated with sodic carbonate. The precipi- tate, consisting of manganous carbonate and ferric hydrate, is wTell washed, added to the rest of the solution, and boiled with it. In this way the iron is precipitated by the manganous carbonate, whilst an equivalent quantity of manganese goes into solution as chloride: 'Fe.fi], + 3COMno" + 30H2 = 'Fe2Ho6 + Ferric Manganous Water, Ferric chloride, carbonate. hydrate. 3MnCl2 + 30O2. Manganous Carbonic chloride. anhydride, The complete precipitation of the iron is ascertained by filtering a sample of the liquid and testing with potassic ferrocyanide. Should copper be present it is best removed with sulphuretted hydrogen. Cal- cium and barium are got rid of by precipitating the manganese with amnionic sulphide, washing the precipitate, and redissolving in hydro- chloric acid. The concentrated solution deposits pink-colored mono- clinic tabular crystals of the aquate, MnCl2,40H2, which on heating are decomposed with evolution of hydrochloric acid. If, however, a solu- tion of this compound be mixed with amnionic chloride, pink regular crystals of the double chloride, MnGI2, are deposited, from 642 INORGANIC CHEMISTRY. which, by careful heating, the water of crystallization may be expelled without further decomposition of the salt; and the anhydrous double chloride, when heated to a higher temperature, parts with ammonic chloride, leaving anhydrous manganous chloride. The anhydrous chlo- ride forms a pink, micaceous, easily fusible mass, which is gradually decomposed by exposure to moist air. The other chlorides of manganese—manganic perchloride, MnCl4, and dimanganic hexachloride, /Mn2CI6—are known only in solution. When the corresponding oxides —manganic peroxide, Mn02, and dimanganic trioxide, /Mn2o3—are dissolved in cold hydrochloric acid, these chlorides are formed ; but on heating they are decomposed with evolution of chlorine, and the solutions contain manganous chloride. Manganous bromide, Mnßr2, is obtained like the chloride, which it closely resembles in properties. It also forms an aquate, Mnßr2,40fI2. Manganous iodide. MnT2, is a white deliquescent mass. Manganous fluoride, MIIF2, is obtained by dissolving manganous carbonate in hydro- fluoric acid. It forms pale-red crystals, insoluble in pure water, soluble in aqueous hydrofluoric acid. Manganic perflvoride, MnF4, is known only in solution. It is formed when manganic peroxide is dissolved in concentrated hydrofluoric acid. Water precipitates from the solution manganic peroxide, but on the addition of potassic fluoride a rose-red precipi- tate of the double fluoride, MnF4,2KF, is formed. COMPOUNDS OF MANGANESE WITH OXYGEN. Manganese forms a large number of oxides, some of which are of great complexity. The following are the most important and best characterized : Manganous oxide, . . . MnO. O-Mn-O “tetroxWe dim“ganfc| MnO1*1110"' o=sk lln=o. or MnMno/;2. NMii/ xx xOx Dimanganic trioxide, . . MnOMuo". i MnOn 51 or 1 MnOo, / \ (ivmu o_Mn—Mn=o. Manganic peroxide, . . Mn02. o=Mn=o. O O r MnO II II Permanganic anhydride, .-j jy[nQ3C). ()—Mn—Mn—O. o o o Manganous oxide, MnO, is formed when the carbonate or any of the higher oxides is heated in a current of hydrogen. It may be pre- pared by fusing anhydrous manganous chloride with sodic carbonate to which a little ammonic chloride has been added. It is a grayish- green powder, which, if it has been prepared at a low temperature, ab- COMPOUNDS OF MANGANESE. 643 sorbs oxygen from the air and turns brown, but if it has been more strongly heated is permanent in air at ordinary temperatures. When heated to whiteness with exclusion of air, it fuses without loss of oxy- gen ; if air be admitted, it is converted on heating into manganous dimanganic tetroxide. It cannot be reduced to metal by heating in a current of hydrogen. By heating in a current of hydrogen containing a trace of hydrochloric acid, it is obtained in the form of small green transparent octahedra with an adamantine lustre. Manganous oxide is the chief salifiable oxide of manganese. Manganese hydrate, MnHo2, is obtained as a white precipitate when a caustic alkali is added to the solution of a manganous salt from which the air has been previously expelled by boiling. When exposed to the air it speedily turns brown from oxidation. It dissolves in solutions of ammonia salts. Manganous dimanganic tetroxide (Dimanganous manganite) 'Mn3G2Mno", or occurs as hausmannite in brownish-black acute quadratic pyramids. This compound represents the most stable stage of oxidation of manganese : thus when the higher oxides are intensely heated, they evolve oxygen and are reduced to this stage, whilst, on the other hand, when manganous oxide or manganous car- bonate is heated in air, oxygen is absorbed and the same compound is produced. The artificial oxide is a reddish-brown powder which, by gentle heating in a slow current of hydrochloric acid, is converted into crystals identical with those of the natural compound. Warm aqueous hydrochloric acid dissolves it with evolution of chlorine and forma- tion of manganous chloride : MnMno"2 + BHCI = 3MnCl2 + Cl2 + 40H2. Dimanganous Hydrochloric Manganous Water, manganite. acid. chloride. Dilute oxy-acids—sulphuric or nitric—dissolve two-thirds of the man- ganese to form a manganous salt, whilst one-third remains as manganic peroxide: MnMno" + 4NG2110 = 2N204Mno" + MnOs + 20H2. Dimanganous Nitric acid. Manganous Manganic Water, manganite. nitrate. peroxide. There are no salts corresponding to this oxide. Its reactions are most readily accounted for on the assumption that it is a dimanganous manganite, as formulated in the two foregoing equations. Dimanganic trioxide, MnOMno" (or /MnX)3).—This compound occurs as the mineral braunite in brownish-black lustrous quadratic pyramids. It may be obtained as a black powder by heating any of the other oxides of manganese in oxygen.—A dimanganic dioxydihy- drate, MnHo2Mno" (or 'Mn202Ho2), occurs as manganite in dark-gray rhombic crystals. The same compound is formed by the spontaneous oxidation of moist manganous hydrate in air. The constitution of the above oxide and hydrate cannot be fixed with certainty. On the one hand, they both yield, with hot nitric acid, manganous nitrate with separation of manganic peroxide: 644 INORGANIC CHEMISTRY. Mn2Os -f 2N02Ho = N204Mno" + MnQ2 + OH2. Dimanganic Nitric Manganous Manganic Water, trioxide. acid. nitrate. peroxide. This reaction would be best accounted for by the first of the alternative formulae above given, in which one atom of manganese is in the dyadic, the other in the tetradic condition. On the other hand, dimanganic trioxide occasionally acts as a basic oxide—in the formation of diman- ganic hexachloride, for example—yielding salts in which the manganese is apparently a pseudo-triad. This behavior would be better explained by the formula 'Mn303. Manganic peroxide (Manganic dioxide, black oxide of manganese), Mn02.—This is, as regards its usefulness, by far the most important of the ores of manganese. It occurs in large quantities as pyrolusite— sometimes in black or dark-gray rhombic prisms, more generally in fibrous or amorphous masses. It may be obtained artificially by care- fully igniting manganous nitrate: N204Mno" = Mn02 + 'N204. Manganous Manganic Nitric nitrate. peroxide. peroxide. The ignited mass is extracted with boiling nitric acid, and the undis- solved residue of manganic- peroxide well washed and then moderately heated. It is also formed by the action of nitric acid upon manganous dimanganic tetroxide or dimanganic trioxide (p. 643). The same oxide is obtained in a hydrated state when a manganous salt is precipitated with an alkaline solution of a hypochlorite. When heated to low redness, manganic peroxide parts with one quarter of its oxygen, yield- ing dimanganic trioxide; at bright redness it parts with one-third of its oxygen, and is converted into manganous dimanganic tetroxide. It dissolves in cold hydrochloric acid with formation of manganic per- chloride; on heating, chlorine is evolved and manganous chloride remains in solution. Nitric acid and dilute sulphuric acid are without action upon it; concentrated sulphuric acid dissolves it on heating with evolution of oxygen and formation of manganous sulphate. In pres- ence of oxalic acid and other oxidizable substances it dissolves in dilute acids in the cold.—Manganic peroxide forms, with basic oxides, com- pounds which may be regarded as salts of a manganous acid of the formula Mn509Ho2. Dipotassic pentamanganite, Mn5O0Ko2, is a yellow powder which separates out when carbonic anhydride is passed into a solution of potassic manganate : 15Mn02Ko2 + 180O2 + 90H2 = Mn509Ko2 + Potassic Carbonic Water. Di potassic manganate. anhydride. pentamanganite. 5'Mn206Ko2 + IBCOHOKO. Potassic Hydric potassic permanganate. carbonate. Manganic peroxide is used in the production of colorless glass (p. 481). It also serves as a cheap source of oxygen, when this gas is required in COMPOUNDS OP MANGANESE. 645 large quantities; but its chief employment is in the preparation of chlorine for the manufacture of bleaching-powder. Regeneration of Manganic Peroxide. Weldon’s Process.—Formerly the residues of manganous chloride obtained in the manufacture of chlorine were allowed to run to waste. At the present day, by means of a process devised by Weldon, the greater part of the manganese is reconverted into manganic peroxide and recovered in this form. For this purpose the chlorine residues (see Preparation of Chlorine, p. 151), which contain, along with manganous chloride, ferric chloride and other impurities, are first treated with calcic carbonate in order to neu- tralize the excess of acid and to precipitate the iron. To the clear solu- tion of manganous chloride and calcic chloride thus obtained milk of lime is added in the proportion of 1J molecules of calcic hydrate to each molecule of manganous chloride. The mixture of manganous hy- drate, calcic hydrate and calcic chloride is then heated by means of a current of steam to a temperature of from 55° to 75° C. (131°-167° F.), after which air is blown through the liquid. Manganous hydrate alone is oxidized only to hydrated dimanganic trioxide, but in presence of excess of lime a rapid oxidation of the manganous hydrate to man- ganic peroxide occurs. The manganic peroxide is obtained in combi- nation with calcic oxide, as calcic raanganite, MnOCao", and it is upon the formation of this compound that the greater readiness of oxidation depends. The oxidation is continued until about three-fourths of the manganese has been converted into peroxide. About 2 cubic metres of air are blown in for every pound of manganic peroxide regenerated, and the time required for the regeneration of a ton of the peroxide is five hours. The “ manganese-mud ”is allowed to settle and, after run- ning off the liquid, is pressed into a solid cake. In this form it is em- ployed in the preparation of chlorine. It usually contains about 33 per cent, of manganic peroxide in combination with lime. Permanganic anhydride,—This compound is obtained by the action of sulphuric acid upon potassic permanganate. The finely powdered pure salt (the absence of chlorine is especially essential, as, otherwise, dangerous explosions may occur, owing to the forma- tion of oxides of chlorine) is gradually added to well-cooled concen- trated sulphuric acid. From the olive-green solution thus obtained reddish-brown oily drops of the anhydride gradually separate—the more readily if the solution be allowed to absorb moisture from the air—and sink to the bottom. Permanganic anhydride is a very unstable com- pound : when rapidly heated it decomposes with a violent explosion. It undergoes slow decomposition at ordinary temperatures, evolving bubbles of oxygen which carry with them violet fumes of the anhy- dride. It is a powerful oxidizing agent: when brought in contact with paper, alcohol, or other organic substances, it causes their ignition. It rapidly absorbs moisture from the air, and dissolves in water with great rise of temperature, yielding a violet-colored solution of perman- ganic acid, a portion of the substance being at the same time decom- posed by the heat evolved. The acid cannot be isolated. 646 INORGANIC CHEMISTRY. OXY-SALTS OF MANGANESE. a. Manganous Salts. Manganous nitrate, N204Mno//,60H2, is prepared by dissolving the carbonate in ni- tric acid. It is difficultly crystallizable and very deliquescent. When heated it fuses, and is converted into manganic peroxide. Manganous carbonate, COMno", occurs native as manganese spar in pink hexagonal crystals. The native compound generally contains iron, calcium, and magnesium. It is precipitated as a white powder when an alkaline carbonate is added to the solution of a manganous salt. When exposed to the air in a moist state it speedily becomes brown from oxidation. Manganous sulphate, S02Mno".—Commercial black oxide of manganese is made into a paste with concentrated sulphuric acid, and the mixture is heated in a crucible, first gently, and afterwards to redness, in order to convert the ferric sulphate into insoluble ferric oxide. The mass is lixiviated, and the solution is digested with a small quantity of manganous carbonate, in order to precipitate the last traces of iron. At a temperature below 6° C. pink rhombic crystals of the formula SOHo2Mno//,60H2, isomorphous with ferrous sulphate, are deposited. From 7° to 20° C. triclinic crystals of the formula SOHo2Mn0",40H2, isomorphous with cupric sulphate, are obtained. Several other aqnates are known. All these salts become anhydrous at 200° C. (392° F.).—With the sulphates of the alkalies manganous sulphate forms double salts, isomorphous with the corresponding double sulphates of the other metals of the dyadic group with the alkalies. Dl- ( SG2Ko potassic manganous sulphate,< Mm/',GOH2, forms monoclinic crystals. ( SOJvo S02—, (so 2 i | Aluminio manganous tetrasulphate, < Mno//('Al'"206)T1,240H,. I so,- i S02—1 This double sulphate occurs as the mineral apjohnite. It has the composition of an alum, and is frequently termed manganese aluminium alum, but inasmuch as it possesses, in common with the other salts in which two atoms of a monad metal in alum are displaced by one atom of a dyad metal, a crystalline form differing from that of the ordinary alums, many chemists refer it to a separate class—that of the pseudo- alums. Other pseudo-alums are known containing iron, zinc, and mag- nesium, as dyad metals. f SO Manganous dithionate, < gQ2 — Finely powdered man- ganic peroxide is suspended in water, and sulphurous anhydride is passed into the liquid, avoiding any rise of temperature. The salt crys- tallizes in pale-red soluble rhombohedra. It forms the starting-point for the preparation of the other dithionates. THE MANGANATES. 647 Manganous silicate, SiOMno//, occurs native as rhodonite in brownish-red crystals.— Dimanganous silicate, SiMno//2, forms the mineral tephroite, which crystallizes in quad- ratic forms. S°;P Manganic sulphate (.Dimanganic trisulphate) S03—(/Mn///206)Vl, is . . ' S03^ obtained by the action of sulphuric acid upon hydrated manganic per- oxide. It is a green powder which deliquesces on exposure to air, and is decomposed at 160° C. (320° F.) with evolution of oxygen. Dipotassic dimanganic tetrasulphate (Manganese alum), h. Manganic Salts. S02Ko~] lo’~ SoJko_J is formed when potassic sulphate is added to a solution of manganic sulphate in dilute sulphuric acid. It crystallizes from very concentrated solutions in violet-colored regular octahedra. Excess of water decom- poses it, manganic hydrate being deposited. With ammonic sulphate a corresponding ammonia manganese alum is obtained. THE MANGANATES. Neither manganic anhydride, MnOs, nor manganic acid, Mn02Ho2, have been prepared ; but salts of this acid, called manganates, are known. These are isomorphous with the corresponding sulphates. Potassic manganate, Mn02Ko2.—When manganic peroxide is fused with caustic potash a deep-green mass is obtained, which contains potas- sic manganate. When the fusion is performed out of contact with air, the reaction takes place according to the equation— 3M02 + 2KHo = Mn02Ko2 + 'Mn203 + 0H2; Manganic Potassic Potassic Di manganic Water, peroxide. hydrate. manganate. trioxide. but if air be admitted, or if nitre or potassic chlorate be added to the mixture, the whole of the manganic peroxide is converted into manga- nate. The mass dissolves without decomposition in a small quantity of water, and the dark-green solution deposits, on evaporation in vacuo, rhombic crystals of potassic manganate, which, when first prepared, are almost black, and display metallic lustre, but become dull and green- colored by exposure to the air. Potassic manganate is stable only in solutions which contain an excess of free caustic alkali; when these solutions are diluted with a large quantity of water, the manganate is decomposed with separation of manganic peroxide and formation of potassic permanganate: SMnOjKoj + 20H2 = {Mnojlo + Mn°’ + 4KHo- Potassic Water. Potassic Manganic Potassic manganate. permanganate. peroxide. hydrate. 648 INORGANIC CHEMISTRY. The color of the solution changes at the same time from green to violet. The same decomposition occurs when carbonic anhydride is passed into the alkaline solution of a manganate. Sodic manganate, Mn02Nao2, is prepared in a similar manner by fusing manganic peroxide with sodic nitrate. It is deposited from its alkaline solutions in crystals of the formula MnO2Nao2.100H2, isomorphous with Glauber’s salt. Baric manganate, is obtained by fusing manganic peroxide with baric nitrate and extracting the mass with water. It is a green powder, insoluble in water. Acids decompose it. Permanganic acid, | *s know only in solution. The solu- tion is obtained, as already described (p. 645), by dissolving per- manganic anhydride in water, or, more readily, by decomposing a solution of baric permanganate with the equivalent quantity of sul- phuric acid. The solution is deep-red by transmitted and blue by reflected light. When heated, or even when exposed to light, it evolves oxygen with separation of hydrated manganic peroxide. Potassic permanganate, j^O^Ko' £reen mass obtained in the preparation of potassic manganate (p. 647) is extracted with boiling water. In this way the manganate is decomposed with for- mation of permanganate (p. 647). If an excess of alkali is present carbonic anhydride ought to be passed into the liquid. The precipitate of hydrated manganic peroxide is removed by filtration through asbestos or glass-wool (filters of paper, calico, or other organic substance would be attacked). The clear liquid, when allowed to evaporate, deposits crystals of potassic permanganate. These are isomorphous with potassic perchlorate. For this reason, if potassic chlorate has been employed in the preparation of the permanganate, the latter salt is apt to be con- taminated with perchlorate, from which it cannot be freed by crystalli- zation, as the two salts crystallize together in all proportions. Potassic permanganate forms large rhombic prisms, which are deep-red by trans- mitted and almost black by reflected light, with a metallic lustre. The salt is soluble in 16 parts of water at ordinary temperatures, more readily soluble in boiling water, yielding a solution of a deep purple color. The solution is a powerful oxidizing agent and destroys most organic sub- stances. A solution of crude potassic permanganate, or more commonly of the sodium salt, is employed as a disinfectant under the name of “Condy’s Disinfecting Fluid.” By exposure to the air the solution of permanganate is gradually reduced by organic matter from the atmo- sphere, changing its color from purple to blue, and at last to green. Owing to these changes this substance was formerly known as mineral chameleon. Boiling with concentrated caustic alkali converts potassic permanganate into manganate with evolution of oxygen : PERMANGANIC ACID AND THE PERMANGANA TES. [ MnOyKo + 2KHo = 2Mn +O 4 OH,. Potassic Potassic Potassic Water Potassic Water, manganate. permanganate, hydrate. COMPOUNDS OF MANGANESE. 649 the chemical change being accompanied by a change in the color of the liquid from purple to green. When the dry salt is heated to 240° C. (404° F.) it evolves oxygen and is converted into a mixture of manga- nate and manganic peroxide : {MnOgKo = Ma0*K“* + Mn°2 +Or Potassic Potassic Manganic permanganate. manganate. peroxide. Sodic 'permanganate, | prepared like the potassium salt. It is manu- factured on a large scale as a disinfectant by fusing black oxide of manganese with crude caustic soda in shallow iron vessels. Ammonic permanganate, | 4mo’ *s °^ained by decomposing the barium salt with ammonic sulphate. It is isornorphous with the potassium salt, which it closely resembles, but is more soluble. It is readily decomposed on heating. Baric permanganate, j^n()3®ao//>—Carbonic anhydride is passed through water in which baric manganate is suspended, and, after filtering from the baric carbonate, the red solution thus obtained is rapidly evaporated. It forms large deep-red rhombic crystals, readily soluble in water. Argentic permanganate, | seParates *n large red crystals, when warm solu- tions of argentic nitrate and potassic permanganate are mixed and allowed to stand. It is sparingly soluble in cold water. COMPOUND OF MANGANESE WITH OXYGEN AND CHLORINE. Permanganic hexoxy-dichloride, | compound sodic chloride is added to a solution of potassic permangan- ate in concentrated sulphuric acid. A yellow gas is evolved, which condenses in a freezing mixture, yielding a greenish-brown liquid. In contact with moist air it emits red fumes. Water decomposes it with formation of permanganic and hydrochloric acids; but these substances at once react upon each other, yielding chlorine and manganic peroxide. It explodes violently on heating. COMPOUND OF MANGANESE WITH SULPHUR. Manganous sulphide, MnS", occurs native as manganese blende in steel-gray granular masses, and occasionally in black cubical crys- tals. The same compound is obtained as a greenish-gray powder by heating any of the oxides of manganese in a current of sulphuretted hydrogen. Alkaline sulphides produce in solutions of manganous salts a flesh-colored amorphous precipitate of hydrated manganous sulphide, which is readily soluble in dilute acids, even in acetic, with evolution of sulphuretted hydrogen, and when exposed to the air be- comes brown from oxidation. By prolonged contact, or by heating, with an excess of the alkaline sulphide, the precipitate is transformed into a green crystalline powder of the formula 3MnS,OH2.—Manganous sulphide unites with the sulphides of the alkali metals to form double 650 INORGANIC CHEMISTRY. compounds. A double sulphide of this description is disulphopotassic trimanganous disulphide, s//. Manganic disulphide, MnS'A, occurs in nature as the mineral hauerite in dark red- dish-brown regular crystals. Characteristic Properties and Reactions of the Com- pounds of Manganese.—The manganous salts are of a pale rose color. Caustic alkalies precipitate white manganous hydrate, which speedily oxidizes and becomes brown. Ammonia only partially precip- itates the manganese as hydrate; in presence of an excess of amnionic chloride ammonia does not produce any precipitate, but the solution on standing absorbs oxygen from the air, and deposits hydrated trimanganic tetroxide. Alkaline carbonates precipitate basic manganous carbonate ; baric carbonate does not precipitate manganous salts in the cold. Am- monia sulphide precipitates flesh-colored hydrated manganous sulphide, soluble in dilute acids, even in acetic acid. All manganous compounds, when fused with sodic carbonate and nitre, yield a green mass containing an alkaline manganate. With borax or microcosmic salt, they give a bead which is amethyst-colored in the oxidizing flame, and colorless in the reducing flame. Manganous chloride colors the non-luminous flame green : the spectrum of the flame exhibits lines in the green and yellow. The spark-spectrum of manga- nese contains a large number of lines. IRON, Fe. Atomic weight 56. Molecular weight unknown. Sp.gr. 7.8. Atom- icity", IV, and. Vl. Evidence of atomicity : Ferrous chloride, .... .... Fe"Cl2. Ferric disulphide, .... .... FewS",. Ferric chloride, .... 'Fe'",Cl#. Potassic ferrate, .... Fevi02Ko2< History.—The process of obtaining iron from its ores has been known from very early times. Owing to its abundance, to the ease with which it can be reduced to the metallic state, and to its valuable properties, it is by far the most important of the metals. Occurrence.—lron is the most abundant and widely diffused of the metals, with the exception of aluminium. Native iron, which is of rare occurrence, may be divided into two kinds—meteoric iron, of extra- terrestrial origin, and telluric iron. Meteoric iron sometimes occurs in considerable masses ; the largest have been found on the island of Disko, off" the coast of Greenland, where there are fifteen of these blocks, the two,largest weighing 21,000 and 8,000 kilos. Weapons and imple- ments of meteoric iron have been found among the Eskimos, and also among tribes in Central Africa. Meteoric iron is never pure: it con- tains varying quantities of other metals, notably nickel and cobalt, the proportion of the first of these sometimes ranging as high as 30 per cent. On the snow-fields of Northern Europe and Asia the snow is IRON, 651 found to inclose minute magnetic particles possessing the composition of meteoric iron. It is probable that this meteoric dust is continually fall- ing upon the earth; but its presence can be detected with certainty only in localities which, like the above, are sufficiently remote from all sources of terrestrial dust. Telluric iron occurs in small spiculse dissem- inated through various basalts and lavas. Masses of terrestrial iron have also been observed in cases in which the fire of burning coal-mines has acted upon ores of iron. This variety is known as natural steel. Iron most frequently occurs in combination with oxygen or sulphur. In combination with oxygen it is found as ferric oxide, 'Fe203, in red haematite, or specular iron ore; as ferrous diferric tetroxide, •| pgQ in magnetic iron ore; as tetraferric trioxyhexahydrate, Fe403Ho6, in brown haematite ; and as ferrous carbonate, COFeo" in spathose iron ore. The disulphide, is of very common occur- rence as iron pyrites. Iron is also found in the form of a sulphide in copper pyrites, | Silicates of iron are contained in nearly all rocks, and by the disintegration and decomposition of these rocks the oxide of iron is produced which imparts to the soil its red color. From the soil plants extract the iron which is a necessary con- stituent of the chlorophyll, or green coloring matter of their leaves. Iron is also a necessary constituent of the haemoglobin, or red coloring matter of the blood. The chlorophyll of plants enables them, with the aid of sunlight, to decompose the carbonic anhydride and aqueous vapor of the atmosphere : a portion of the oxygen resulting from this decomposition is evolved, whilst the other products of decomposition are used in building up the tissues and principles of the plant. The haemoglobin of the blood acts as a carrier of the oxygen which is ab- sorbed during respiration, and which serves for the oxidation of the animal tissues. In this way the respiratory functions both of plants and of animals are dependent upon the presence of iron. The presence of iron in extra-terrestrial space is proved by its occur- rence in meteorites, and, further, by the results of spectrum analysis, which show that this metal is present in the sun and in many of the fixed stars. Extraction.—The important and complex subject of the metallurgy of iron can only be briefly sketched here. The compounds of iron from which the metal is extracted are the oxides, the hydrates, and the carbonate. The chief ores are: magnetic iron ore, red haematite, brown haematite, spathose iron ore, and clay iron- stone or argillaceous iron ore, which is a spathose iron mixed with clay or sand. JBlack band is a variety of clay iron-stone containing from 20 to 25 per cent, of coal. The ores are first calcined or roasted. In this process water and carbonic anhydride are expelled, whilst most of the sulphur, which may be present, is oxidized and burnt off as sul- phurous anhydride. At the same time the ore is rendered more friable and porous. The ore is then reduced by heating with coal, limestone, and occasionally silicates, in a hot-blast furnace. This furnace consists of a lofty shaft of strong masonry lined with fire-brick. The internal 652 INORGANIC CHEMISTRY. space is narrower towards the bottom, where the molten metal collects. The furnace is first lighted or bloim in, after which alternate layers of a mixture of calcined ore and limestone on the one hand, and of coal on the other, are thrown in at the top until the furnace is full. A powerful blast of air, previously heated to from 350° to 700° C. (662-1292° F.), is forced in through pipes or tuyeres placed at the bottom of the furnace. The chemical changes which occur in the furnace are as follows: The oxygen of the air on entering the furnace unites with the carbon to form carbonic anhydride, which in turn is converted into carbonic oxide by contact with the heated carbon. The carbonic oxide in passing upwards over the heated ferric oxide reduces it to finely-divided iron. The part of the furnace in which this change occurs is termed the “ zone of reduction,” At the same time the fusible flux of silicate of lime coats the particles of metal and protects them from oxidation. As the reduced iron sinks into the hotter parts of the furnace it begins to combine with carbon; this part of the furnace is therefore known as the “ zone of carburation.” At this point the iron also takes up phosphorus derived by reduction from phosphates con- tained in the ore. The metal gradually sinks till it reaches the hottest part of the furnace—the “ zone of fusion ”—when it melts and runs down to the hearth or lowest part of the furnace. Here it would be exposed to the danger of oxidation from the blast; but the fusible slag floats on the surface of the molten metal and protects it. The excess of slag runs off regularly through an opening. From time to time the molten iron is tapped and cast into bars known as pigs. As fast as the charge in the furnace sinks, fresh charges of ore, limestone, and coal are introduced. In this way a blast-furnace may be kept con- tinuously at work for many years. The crude iron thus obtained, known as pig iron or east iron, contains from 3 to 6 per cent, of carbon, together with varying quantities of manganese, silicon, sulphur, phosphorus, arsenic, and antimony. The carbon is present in two forms: partly in chemical combination, and partly as particles of graphite mechanically disseminated throughout the mass of the metal. When cast iron is dissolved in acids, the carbon displays a different behavior according to the form in which it is present: the mechanically disseminated carbon is left behind unchanged, whilst the chemically combined carbon enters into combination with hydrogen to form complex hydrocarbons, gaseous and liquid. According to color and other properties, the following varieties of cast iron are distin- guished : White cast iron, which contains the whole of its carbon in the combined condition ; and gray cast iron, which, in addition to the combined carbon, contains graphite disseminated throughout its mass. Various intermediate stages are classed as mottled cast irons. Spiegeleisen, spiegel, or specular pig iron is a white iron containing the highest per- centage (3.5 to 6 per cent.) of combined carbon. White iron is formed when the temperature of the blast furnace is low. It contracts on solidification, and therefore cannot be used for castings. Gray iron is formed when the temperature is high. It expands on solidifying, and is suitable for foundry work. Cast iron is brittle and cannot, as a rule, be forged. In order to IRON 653 impart to it the property of malleability, the greater portion of the carbon and the other foreign substances must be removed by a process of oxidation. In this way the cast iron is converted into wrought iron. The process most commonly employed in the production of wrought iron is that of puddling: the wrought iron is fused along with powdered haematite on the hearth of a reverberatory furnace, employing a flux of blast-furnace slag. During the process, the metal is stirred to promote oxidation. The silicon is first converted into silicic anhydride, which is taken up by the bases of the slag; afterwards, the carbon is burnt off as carbonic anhydride. A comparatively low temperature is essential to the effectual removal of the phosphorus, since at a high temperature the iron reduces the phosphates contained in the slag and takes up phosphorus. Wrought iron contains from 0.15 to 0.5 per cent, of carbon. The lower the proportion of carbon the more malleable and the less readily fusible is the iron. Rolled and hammered wrought iron, containing 0.3 per cent, of carbon, has a fibrous structure; if the percentage rises to 0.5, the structure becomes granular and crystalline. The hardness of the metal also increases with the percentage of carbon. Wrought iron is of a clear gray color, and capable of taking a high polish. At a red heat it softens and may be welded. The physical properties of iron are powerfully modified by the presence of minute quantities of various impurities: thus sulphur renders the metal “ red-short ”—that is, brittle at high temperatures; phosphorus renders it “cold-short,” or brittle at ordinary temperatures. If the proportion of chemically combined carbon in iron lies between 0.6 and 2 per cent., the product is known as steel. In chemical compo- sition, steel therefore stands midway between wrought iron and cast iron, and it may in fact be produced from the former of these by in- creasing, and from the latter by diminishing, the proportion of carbon present. Steel was formerly exclusively prepared from wrought iron by the cementation process. In this process bars of wrought iron are packed in powdered charcoal or soot, and heated to bright redness for from seven to ten days, according to the nature of the product required. In this way the iron takes up the carbon necessary for its conversion into steel. The exact mode in which this is-accomplished is not perfectly understood, though various hypotheses have been made with regard to this process. The bars of steel, after their conversion, exhibit a peculiar blistered appearance due to the production of gas within the mass of the metal. This imperfection is removed by hammering and rolling, or by melting the steel. Puddled steel is an inferior quality of steel pre- pared from cast iron by arresting the process of puddling at a point short of the production of wrought iron. In the Bessemer process of steel making, cast iron is melted, and then transferred to a vessel known as a converter, through the bottom of which a powerful blast of air is blown. The silicon, manganese, and carbon are thus oxidized, and so great is the heat evolved that the temperature of the molten metal rises considerably. Formerly the process was interrupted at the point of formation of steel, but at the present day the oxidation is carried on until the whole of the carbon is removed—a point much more readily INORGANIC CHEMISTRY. ascertained—after which the molten spiegel is added in quantity exactly sufficient to convert the whole into steel. Steel is of a clear gray color, and possesses a granular structure. It may be forged and welded like wrought iron, and fuses at a lower temperature than the latter. It possesses the property of becoming intensely hard and brittle when heated to redness and then suddenly cooled—for example, by plunging into water. This hardness and brittleness can be removed in any required degree by heating the hard- ened steel to temperatures between 200° and 300° C. (392-572° F.) and then allowing it to cool. This process is known as tempering. The lower the temperature employed, the harder will be the resulting steel. If the surface of the object to be tempered be first polished, it will ex- hibit shades of color on heating, due to the formation of films of oxide of varying thickness. By observing these colors the workman is enabled to judge with sufficient accuracy of the temperature which he is em- ploying. The specific gravity of hardened steel is somewhat lower than that of wrought steel. In hardened steel the whole of the carbon is present in the combined state, whereas wrought steel also contains graphitic carbon. Preparation of Pure Iron.—The purest iron of commerce is piano- forte wire, which contains only about 0.3 per cent, of impurities—for the most part carbon. Chemically pure iron is prepared by heating the pure oxalate or oxide in a current of hydrogen. It is thus obtained in the form of a black powder, which, when the reduction has been effected at a sufficiently low temperature, is pyrophoric, spontaneously oxidizing with incandescence when exposed to the air. If heated to a higher temperature during reduction, the product is denser and no longer spontaneously oxidizable. It may be fused into a regulus in a lime crucible by means of the oxyhydrogen flame. Very pure iron may also be obtained by fusing wrought iron with ferric oxide under a layer of melted glass free from lead. Properties.—Pure iron is almost silver-white, and is capable of tak- ing a high polish. It has a specific gravity of 7.84. It is softer, more malleable, and less tenacious than wrought iron. It is fusible only at the very highest temperatures. It does not undergo any change in dry air at ordinary temperatures ; but in moist air containing carbonic an- hydride it becomes coated with ferric hydrate or iron rust. The pro- cess of rusting takes place very slowly at first, but goes on rapidly as soon as a thin coating of rust has been formed upon the surface of the metal. When heated in air, iron becomes coated with ferrous diferric tetroxide, | which, on hammering, flies off in scales (smithy scales). It burns brilliantly when heated in oxygen, emitting showers of dazzling sparks, and yielding the foregoing oxide. It burns also in sulphur vapor. It combines directly with the halogens, and at a high temperature with carbon. At a red heat it decomposes water, with evo- lution of hydrogen, and formation of ferrous diferric tetroxide. Dilute hydrochloric or sulphuric acid dissolves it with evolution of hydrogen, and when the metal contains chemically combined carbon, hydrocar- bons are mixed with the hydrogen, imparting to it a peculiar and disa- COMPOUNDS OF IRON. 655 greeable odor. Nitric acid of sp. gr. 1.35, or lower, dissolves iron with violent evolution of nitrous fumes and formation of ferric nitrate; but under certain circumstances iron may be kept immersed for any length of time in nitric acid without the slightest action, or diminution of its brightness of surface. This condition, which is known as the passive state of iron, is produced when the metal is immersed in nitric acid of sp. gr. 1.45 or higher. The iron which has been thus rendered passive is not acted upon by dilute nitric acid. The same condition is induced when iron is made the positive plate of a voltaic couple in nitric acid : for example, when it is introduced into nitric acid of sp. gr. 1.35 in contact with a piece of platinum. The platinum may then be removed, and the iron remains passive. Passive iron does not pre- cipitate copper from its solutions, but if a piece of passive iron which has been dipped into the solution of a copper salt be scratched, the cop- per is instantly deposited on the whole surface of the iron. Passive iron is powerfully electronegative towards ordinary iron, and a voltaic couple may be constructed consisting of passive iron in concentrated nitric acid and ordinary iron in a solution of sodic sulphate, the two liquids being separated by a porous diaphragm. The phenomenon of passivity in iron depends upon the formation of a thin film of ferrous diferric tetroxide upon the surface of the metal. Thus iron may be rendered passive by moderately heating it. The deposition of copper in the case above described depends upon the fact that by scratching the passive metal the film of oxide is removed at that part and a sur- face of iron exposed ; a voltaic action then sets up between the electro- positive iron and the electro-negative oxide, and the hydrogen which is liberated on the surface of the latter reduces it, converting it into iron, which in its turn reduces the copper. The voltaic action between iron and ferrous diferric tetroxide may be employed in rendering the metal passive : thus if orffe end of a bright iron wire be heated so as to oxidize it, and then the wire be dipped, with the oxidized end first, into nitric acid of sp. gr. 1.35, the whole wire is rendered passive.—lron is at- tracted by the magnet, and may also be magnetized, but parts with its magnetism almost instantaneously, whilst steel is capable of permanently assuming the polar state. COMPOUNDS OF IRON WITH THE HALOGENS. a. Ferrous Compounds. Ferrous chloride, Fed,, is prepared by heating iron in gaseous hydrochloric acid. A solution of this compound is obtained by dis- solving iron in aqueous hydrochloric acid. The anhydrous chloride sublimes in colorless fusible six-sided scales. When volatilized in an atmosphere of gaseous hydrochloric acid, it possesses a vapor density lying between the densities required for the molecular fornmlte FeCl2 and respectively. It is therefore probable that the iron in this compound is at lower temperatures tetradic and at higher temperatures dyadic. When heated in air ferrous chloride is converted into ferric chloride, which volatilizes, and ferric oxide: 656 INORGANIC CHEMISTRY. 6FeCl2 + 30 = Te2C]6 + Te203. Ferrous Ferric Ferric chloride. chloride. oxide. It is deliquescent, and soluble both in water and in alcohol. The aque- ous solution, when concentrated out of contact with air, deposits pale- green deliquescent crystals of the formula FeCl2,40H2. The crystals absorb oxygen from the air and undergo decomposition. Ferrous chlo- ride forms double compounds with the chlorides of the alkalies. Potas- sic ferrous chloride, FeCl2,2KC1,20H2, is deposited from mixed solu- tions of its component chlorides in bluish-green monoclinic crystals. Ferrous bromide, Feßr2, is obtained as a yellowish crystalline mass when bromine vapor is passed over iron filings heated to low redness. The aqueous solution, pre- pared by dissolving iron in hydrobroraic acid, deposits on concentration the aquate, Fe8r2,60H2, in green tabular crystals. Ferrous iodide, Fel2, is obtained as a gray laminated mass by heating iron filings in a closed crucible and adding small quantities of iodine. An excess of iodine is then added, and the heating is continued until vapors of iodine cease to escape. The aque- ous solution, which is readily obtained by digesting iron filings with iodine and water, deposits on evaporation green crystals of the formula FeT2,40H2. Ferrous fluoride, FeF2.—When iron is dissolved in hydrofluoric acid, sparingly solu- ble green crystals of the compound FeF2,BOH2 are deposited, which, when heated with exclusion of air, become anhydrous. b. Ferric Compounds. Ferric chloride, 'Fe2C'6. Molecular volume I I I.—This compound is obtained in the anhydrous state by gently heating iron wire in a cur- rent of chlorine, and in solution by dissolving ferric oxide in hydro- chloric acid or iron in aqua-regia. The anhydrous compound forms dark-brown hexagonal plates, which possess a green metallic lustre, and appear red by transmitted light. It is fusible, and volatilizes more readily than the ferrous compound. It deliquesces in moist air, and is readily soluble in water; the dilute solution is yellow, the concentrated solution is dark-brown, and of an oily consistency. It is also soluble in alcohol and in ether: the latter solvent extracts the compound from the aqueous solution when agitated with it. The aqueous solution when concentrated over sulphuric acid deposits yellow prismatic crystals of the compound 'Fe2Cl6,120H2, and at a still higher degree of concen- tration brownish-red crystals having the formula 'Fe2Cl6,601I2. When the hydrated chloride is heated, it parts with water and hydrochloric acid, yielding an oxychloride, which at a higher temperature decomposes into volatile anhydrous ferric chloride and ferric oxide. A dilute aque- ous solution, containing less than 4 per cent, of ferric chloride, is de- composed on heating into soluble colloidal ferric hydrate (p. 658) and free hydrochloric acid, this chemical change being accompanied by a change in the color of the liquid from yellow to red. When a concen- trated aqueous solution is evaporated by heat it parts with hydrochloric acid and an insoluble oxychloride of varying composition separates out. —Ferric chloride forms numerous double compounds, Potassic ferric chloride, Fe2Cl6,4KC1,20H2, is deposited in garnet-red crystals from mixed solutions, of ferric and potassic chlorides. Anhydrous ferric chloride absorbs gaseous ammonia, yielding thecompound which in appearance is indistinguishable from ferric chloride. 657 COMPOUNDS OF IRON. Ferric bromide, /Fe2Brfi, is prepared by heating iron in an excess of bromine vapor. In its properties it closely resembles the chloride. Ferric iodide has not been obtained. It appears, however, to be capable of existing at higher temperatures. When the heated mass which is obtained in the preparation of ferrous iodide (p. 656), and which remains after all the vapors of iodine have been ex- pelled, is allowed to cool, it suddenly evolves, at a temperature somewhat below redness, large quantities of iodine vapor, a phenomenon which is probably due to the decompo- sition of ferric iodide contained in the mass. Ferric fluoride, /Fe2Fg, is formed when ferric oxide is dissolved in hydrofluoric acid. It forms colorless sparingly soluble crystals of the formula /Fe2F6,90H2. By heating these in a platinum crucible over the blowpipe, the water of crystallization is expelled, and the anhydrous fluoride is obtained as a fused mass. It sublimes in small trans- parent almost colorless cubes, isomorphous with aluminic fluoride. COMPOUNDS OF IRON WITH OXYGEN. Ferrous oxide, .... FeO O—Fe—O Ferrous diferric tetroxide f FeOT7, ~ r\ t/ r\ (Magnetic oxide),. .tFeOFeo" o=Fe-Fe=o. o r FeO /X Ferric oxide, . . . . < O=Fe—Fe=o. Ferrous oxide, FeO, is difficult to prepare in a state of purity. It is obtained as a black powder by heating ferric oxide to redness in a mixture of equal volumes of carbonic anhydride and carbonic oxide, or by heating ferric oxide to 300° C. (572° F.) in a current of hydrogen. The product obtained by the latter method undergoes oxidation with incandescence if exposed to air when freshly prepared, but loses this pyrophoric property after remaining for twelve hours at ordinary tem- peratures in an atmosphere of hydrogen. Ferrous hydrate, FeHo2, is formed when caustic alkali is added to the solution of a ferrous salt. The precipitation, washing and drying must be performed in an atmosphere free from oxygen. When pure it forms a white powder, but generally has a greenish tint, owing to the difficulty of entirely excluding oxygen. When exposed to the air it rapidly absorbs oxygen, and is converted into ferric oxide, sometimes with incandescence. Ferrous diferric tetroxide (Magnetic oxide), /Fe2o2Feo//.— This compound occurs native in black lustrous octahedra and other forms belonging to the regular system, more frequently, however, in granular masses, constituting the mineral magnetic iron ore. It is formed when iron is heated in steam or carbonic anhydride, with libera- tion of hydrogen and formation of carbonic oxide respectively. On the other hand, by precisely the reverse reactions, hydrogen and carbonic oxide reduce heated oxides of iron to the metallic state. When iron is heated in air it becomes coated with magnetic oxide in the form of so- called iron scale or smithy scales. This is not, however, a pure com- pound : the outer portions approximate more in composition to ferric oxide, ;Fe203, the inner portions, which are next the metal, to that of ferrous oxide. Ferrous diferric tetroxide is attracted by the magnet, 658 INORGANIC CHEMISTRY. and the native variety frequently possesses the property of attracting iron. This naturally magnetic variety of the mineral is known as loadstone, and its magnetism is derived from that of the earth. Ferric oxide, occurs as specular iron ore in lustrous steel- gray hexagonal crystals, also massive, as the important iron ore red haematite. It may be obtained artificially in reddish-brown lustrous scales by carefully heating a mixture of ferrous sulphate and common salt, extracting the mass with water : 2S02Feo" = 'Fe203 + S02 + S03. Ferrous Ferric oxide. Sulphurous Sulphuric sulphate. anhydride. anhydride. The same compound is obtained in the amorphous condition as a reddish powder by heating ferric hydrate or ferrous sulphate alone. The native oxide and the strongly ignited amorphous oxide dissolve with great diffi- culty in acids. Amorphous ferric oxide, obtained as a by-product in the manufacture of fuming sulphuric acid (p. 274), is employed as a red pigment under the name of rouge. It is also used in polishing jewellers’ goods and metallic surfaces generally. Ferric hydrate, /Fe2Ho6, is obtained as a bulky reddish-brown precipitate by adding ammonia to a solution of ferric chloride. When dried at ordinary temperatures it has the composition represented by the above formula, but wdien heated to 100° C., or when boiled with water, or even when left for a long time in contact with water, it undergoes partial dehydration, r Fello and is converted into the compounds of the formulae | jj 2C* and Hydrates of this composition occur in nature as needle iron ore or hrown iron ore. Ordinary iron rust has the composition f FeOHo FeIIo2 . 7 O , and this compound also occurs in nature as hrown haematite. FeHo2 v FeOHo Various other complex hydrates occur as wrell-characterized minerals. A soluble ferric hydrate is also known. Thus a solution of ferric chlo- ride dissolves large quantities of freshly precipitated ferric hydrate, yielding a dark-red liquid. The same solution may be obtained by adding ammonic carbonate to a solution of ferric chloride until a point is reached at which the precipitate of ferric hydrate no longer redis- solves. If either of these solutions be subjected to dialysis, ferric chlo- ride passes through the dialyser and a dark-red liquid remains, con- taining only 1.5 per cent, of hydrochloric acid to 98.5 of ferric oxide. Traces of alkalies and salts cause the solution to coagulate. All the ferric hydrates are converted on heating into ferric oxide. 659 COMPOUNDS OF IRON, OXY-SALTS OF IRON. a. Ferrous Salts. Ferrous nitrate, is best prepared by decomposing ferrous sulphate with baric nitrate. Crystals can be obtained only from well-cooled solutions. The crystals are very unstable, and by exposure to air are speedily converted into a red- dish-brown powder. Ferrous carbonate, COFeo", occurs native as spathose iron ore in rhombohedral crystals, which, however, generally contain varying quantities of the isomorphous carbonates of calcium, magnesium, and manganese. This compound may be obtained artificially in microscopic rhombohedra by heating a solution of ferrous sulphate with an excess of hydric sodic carbonate in sealed tubes to 150° C, (302° F.). Alka- line carbonates produce in solutions of ferrous salts a white precipitate of ferrous carbonate, which speedily becomes dark-colored from oxida- tion, and when exposed to air is eventually transformed into ferric hydrate with evolution of carbonic anhydride. Ferrous carbonate is soluble in water containing carbonic anhydride. It is in this form that iron usually occurs in chalybeate springs. Ferrous sulphate (Green vitriol), SOHo2Fe0",60H2.—This salt is prepared on a large scale by exposing moistened iron pyrites, FeS"2, to the air. The soluble ferrous sulphate, together with the excess of sulphuric acid, thus formed, runs off into tanks, where the excess of acid is also converted into ferrous sulphate by the addition of scrap iron. It is best prepared in a state of purity by dissolving pure iron wire in sulphuric acid, employing an excess of the metal. It forms large pale- green monocliuic crystals, which effloresce in dry air. These are solu- ble in 1|- times their weight of water at ordinary temperatures, and in of their weight of boiling water. The salt loses its 6 molecules of water of crystallization at 100° C.; at 300° C. (572° F.) it parts with its water of eonstitution, leaving white anhydrous SOgFeo'k The anhydrous salt is decomposed when heated to redness, yielding ferric oxide, together with sulphurous and sulphuric anhydrides (p. 658). The moist salt absorbs oxygen from the air and turns brown. Ferrous sulphate also crystallizes in the rhombic forms of zincic sulphate. Crystals of this form may be obtained by introducing a small crystal of zincic sulphate into a supersaturated solution of ferrous sulphate. If, on the other hand, a crystal of cupric sulphate be employed to start the crystallization, trinclinic crystals of the formula SOHoaFeo", 40H2, isomorphous with those of cupric sulphate, are obtained. Fer- rous sulphate crystallizes in all proportions with sulphates of copper, zinc, manganese, and the other metals of the isomorphous dyadic group, and cannot be purified by crystallization if any of these are present. Ferrous sulphate is employed in the preparation of inks, iron mor- dants, etc.—Ferrous sulphate forms, with the sulphates of the alkalies, double sulphates isomorphous with the double sulphates of the metals of the magnesium group with the alkalies. Amnionic ferrous sulphate, t S02Amo < Feo" ,60H2, is obtained by dissolving equivalent quantities of fer- SC2Amo 660 INORGANIC CHEMISTRY. rous sulphate and ammonic sulphate in a small quantity of hot water and allowing the solution to crystallize. It forms transparent bluish- green raonoclinic crystals. It is much more permanent in air than fer- rous sulphate, and for this reason is largely used instead of this salt in volumetric analysis. Ferrous phosphate, PFAFec/^BOII.^.—This compound occurs as the mineral vivianite in thin monoclinic prisms, generally of bluish-green tint. It is precipitated on the addition of hydric disodic phosphate to a solution of ferrous sulphate as a white amorphous powder which rapidly becomes blue from oxidation. Ferrous silicate, Siireo//.2, occurs native as the mineral fayalite. It also forms the chief constituent of refinery-slag, obtained in the process of refining iron previous to puddling. It also occui’S in combination with other silicates in a great variety of min- erals. Ferric nitrate, N6012 (/Fe///206)vi, is obtained by dissolving iron in an excess of nitric acid, and carefully evaporating the solution. On adding nitric acid, crystals of the nitrate are deposited, sometimes with 12, sometimes with 18 aq., according to the concentration of the solu- tion and the quantity of nitric acid employed. The crystals are deli- quescent and readily soluble in water, but only sparingly soluble in nitric acid. The brown aqueous solution is decomposed on boiling, with separation of brown insoluble basic nitrates. Ferric nitrate is employed as a mordant. S02~n Ferric sulphate, S02-(/Fe'/'206)T1,901I2, occurs native in hexagonal . So^ crystals as the mineral coquimbite. It is best prepared by dissolving 10 parts of ferrous sulphate in water, together with 4 parts of concen- trated sulphuric acid, and adding to the hot solution small quantities of nitric acid until a portion on testing with potassic ferri-cyanide no longer gives a blue precipitate. The reaction is as follows: h. Ferric Salts. 6S02Feo" + 3S02Ho2 + 2N02Ho = 3S306('Fe'"206)- Ferrous Sulphuric Nitric Ferric sulphate, sulphate. acid. acid. + 2'N//0 + 40H2. Nitric Water, oxide. By evaporation the anhydrous salt is obtained as a white mass. It is soluble in water, yielding a brown solution, but insoluble in concen- trated sulphuric acid. Basic ferric sulphates of varying composition are obtained by boiling the dilute solution of ferric sulphate or by adding to its solution a quantity of alkali insufficient for complete pre- cipitation. Dipotassic diferric tetrasulphate {lron alum), S02Ko-| ('Fe" W,24OHr- S022KqJ This compound is obtained when the calculated quantity of potassic sulphate is dissolved in a solution of ferric sulphate, and the concen- COMPOUNDS OF IRON. 661 trated solution is kept at a temperature of 0° C. The alum is depos- ited in violet octahedra, soluble in 5 parts of water at ordinary tem- peratures. Ferric phosphate, P202(/Fe///206)vi,40H2, is obtained as a white precipitate when hydric disodic phosphate is added to a solution of ferric chloride. It is insoluble in water and in acetic acid, but soluble in mineral acids. Ferric silicates.—A dihydric diferric disilicate, Si Ho (/^67/7206)v' oCClirs native as the mineral anthrosiderite. Ferric silicates also occur in combination with other silicates in a large number of minerals. THE FERRATES. Neither ferric acid, Fe02Ho2, nor its anhydride, Fe03, is known. When ferric acid is liberated from its salts, it is instantaneously decom- posed into ferric hydrate and free oxygen. JPotassie ferrate, Fe(),,Ko.r—This compound is prepared by suspend- ing freshly precipitated ferric hydrate in caustic potash and passing a rapid current of chlorine through the liquid, care being taken, however, that the temperature does not rise above 40° C. (104° F.). It is also formed when a positive electrode of cast iron is employed in the elec- trolysis of caustic potash, and when finely divided iron is fused with nitre. It forms small dark-red crystals, which appear almost black by reflected light. It dissolves in water, yielding a red solution which on standing deposits ferric hydrate and becomes colorless, oxygen being evolved. The same change takes place instantaneously on heating. So die ferrate, Feo2Nao2, is prepared like the potash salt, which it closely resembles. Baric ferrate, Feo2Bao//, is obtained as a red insoluble precipitate when baric chloride is added to the solution of the potash salt. It is moderately stable and may be heated to 100° C. without decomposition. COMPOUNDS OF IRON WITH SULPHUR. Ferrous sulphide, FeS", is formed by the direct union of its elements. Red-hot wrought iron or steel, but not cast iron, undergoes apparent fusion when brought in contact with a roll of sulphur, owing to the formation of the more fusible monosulphide. The same com- pound is formed with evolution of heat when a mixture of iron filings and sulphur is moistened with water and allowed to stand at ordinary temperatures. It is best prepared by throwing a mixture of 3 parts of iron filings and 2 parts of sulphur in small portions at a time into a red hot Hessian crucible. It is thus obtained as a black porous mass, which at a higher temperature fuses, solidifying to a grayish-yellow, crystalline, metallic mass, of sp. gr. 4.79. The alkaline sulphides precipitate from solutions of ferrous or ferric salts black amorphous ferrous sulphide. In this form it is readily oxidized if exposed to the air in a moist state. Dilute hydrochloric or sulphuric acid dissolves ferrous sulphide with evolution of sulphuretted hydrogen. r j»eS" Diferric trisulphide, < compound cannot be pre- pared by precipitating a ferric salt with amnionic sulphide, as under 662 INORGANIC CHEMISTRY. these circumstances a mixture of ferrous sulphide with sulphur is obtained. It is formed when iron is heated with its own weight of sulphur, avoiding too high a temperature. It is thus obtained as a yellowish metallic mass of sp. gr. 4.41. This compound may be regarded as the sulphanhydride of the sulpho-acid, | Fgg/zjjg* This acid is not known, but its salts have been prepared. Thus potassie sulphoferrite, | ‘s obtained in the form of red, lustrous, flexi- ble needles when a mixture of 1 part of finely-divided iron, 6 parts of dry potassie carbonate, and 6 parts of sulphur is fused and the cooled mass extracted with water. Copper pyrites, / ]?gg//(/Cu/2S//2)//, is the cuprous salt of this sulpho-acid TTeptnferric octosidphide (.Magnetic pyrites), Fe7S8, occurs native in brownish-yellow metallic, hexagonal crystals, more frequently, however, massive. This substance is attracted by the magnet, and is sometimes itself magnetic. Ferric Disulphide, FeS"2.—This compound occurs native in two distinct forms. As iron pyrites it is found in large quantities, either massive or in brass yellow crystals belonging to the regular system. It has a specific gravity of 5.185. The same compound is obtained artifi- cially by heating finely-divided iron with excess of sulphur to a tem- perature below redness. The native compound appears to have been formed by the reducing action of organic matter upon ferrous sulphate dissolved in water, and hence it is chiefly found along with the remains of organic matter such as coal, peat, etc. Sometimes it assumes the form of the piece of organic matter by which the reduction has been effected : thus wood, roots, ammonites, and other organized forms are found accurately reproduced in this material. Marcasite, or radiated pyrites, the second form of ferric disulphide, occurs in pale brass-yellow rhombic crystals with a sp. gr. of 4.68 to 4.85. Neither of the forms of iron pyrites is magnetic. It is not attacked by dilute acids or by cold concentrated sulphuric acid ; but hot concentrated sulphuric acid slowlv dissolves it with evolution of sulphurous anhydride. Hot nitric acid also oxidizes and dissolves it. When heated in a current of hydrogen it is reduced to the monosulphide. It burns with a flame when heated in air, yielding sulphurous anhydride and ferric oxide. In this way it is employed in enormous quantities in the manufacture of sulphuric acid. General Properties and Reactions of the Compounds of Iron : a. Ferrous salts.—The aquates of these salts are green, the anhydrous salts are white. Caustic alkalies precipitate white ferrous hydrate, which speedily oxidizes by exposure to the air and becomes green. Ammonia only partially precipitates solutions of ferrous salts as hydrate ; in presence of an excess of ammonic chloride no precipitate is produced by ammonia, but the araraoniacal solution absorbs oxygen from the air, and a film of ferric hydrate forms upon the surface. Sulphuretted hydro- gen does not precipitate ferrous salts in acid solution ; ammonic sulphide precipitates black hydrated ferrous sulphide, which is readily oxidized 663 COBALT. by exposure to air. Potass ic ferrocyanide gives a white precipitate of dipotassic ferrous ferrocyanide (Fe"Cy6Fe"K2), which rapidly oxidizes and becomes blue. Potassic ferricyanide occasions a deep-blue precipi- tate of ferrous ferricyanide (Turnbull’s blue) (Fe'/3/Fe'/'2Cyl2). Oxidiz- ing agents convert the ferrous into ferric salts. b. Ferric salts.—These have a yellow or reddish-brown color. Caustic alkalies and ammonia give a reddish-brown bulky precipitate of ferric hydrate, insoluble in excess. Sulphuretted hydrogen does not precipitate the iron but reduces it to the ferrous state, whilst finely divided white sulphur is deposited. Ammonic sulphide precipitates black ferrous sulphide with separation of sulphur. Potassicferrocyanide gives a deep-blue precipitate of ferric ferrocyanide (Prussian blue) (3Fe//Cy2,2'Fe///2Cy6). Potassic ferricyanide gives no precipitate with solutions of ferric salts ; but the color of the liquid changes from yellow to reddish-brown. Soluble thiocyanates give a blood-red coloration which is not destroyed by hydrochloric acid. Baric carbonate precipi- tates the whole of the iron in the cold as ferric hydrate with evolution of carbonic anhydride. Sadie acetate col.ors neutral solutions dark-red, and, on boiling, the whole of the iron is precipated as basic ferric acetate. The benzoates and succinates of the alkali-metals produce in neutral solutions bulky insoluble brown precipitates. All compounds of iron when heated with sodic carbonate on charcoal in the reducing flame yield a black magnetic powder. Borax and microcosmic salt give with iron compounds beads which in the reducing flame are bottle-green and in the oxidizing flame yellow, or, if the quantity of iron is very small, colorless. The compounds of iron do not color flame. The spark-spectrum of the metal contains many hundreds of bright lines coincident with lines of the solar spectrum. COBALT, Co. Atomic weight = 58.6. Molecular weight unknown. Sp. gr. 8.5 to 8.7. Atomicity", IV, andvi ? Also a pseudo-triad. Fvidence of atomicity : Cobaltous chloride, .... Oo"Cl2. Cobaltic disulphide .... CoivS"2. Cobaltic oxide, / Co/;/On * * * 1 Co'"Ou' History.—Cobalt was discovered by Brandt in 1735. Occurrence.—Metallic cobalt occurs in small quantity in meteoric- iron. Its chief ores, which are not very widely distributed, are the " ( As arsenides and arsenical sulphides, such as speiss-cobalt, ] AsCo//’ and (As glance-cobalt, o"S'')"2. In almost all the cobalt minerals a por- tion of cobalt is replaced by nickel, iron, and other isomorphous metals. Cobalt is present in the solar atmosphere. Extraction.—The ores of cobalt, which consist, as above stated, of mixed arsenides and sulphides of cobalt, nickel, and iron, and generally contain, in addition, copper, bismuth, and other metals, are first roasted 664 INORGANIC CHEMISTRY, in a current of air. In this way an impure cobaltous arsenate, known as zaffre, is obtained, whilst large quantities of arsenious anhydride are volatilized, this product being carefully condensed. The roasted mass is extracted with hydrochloric acid until the residue is free from cobalt. On evaporating the solution chlorine is evolved, the arsenic acid being reduced by the hydrochloric acid to arsenious acid, which crystallizes out. The remainder of the arsenic is got rid of by oxidizing the arseni- ous acid back to arsenic acid by the addition of bleaching powder, care- fully avoiding an excess, and then exactly neutralizing with milk of lime. In this way ferric hydrate is precipitated, carrying with it all the arsenic acid. The solution is then again acidified with hydrochloric acid, and treated with sulphuretted hydrogen in order to precipitate copper, bismuth, etc. The cobalt is then precipitated from the weak acid solution as cobaltic oxide, by the careful addition of bleaching powder. An excess of the precipitant is to be avoided, as this would bring down the nickel. The crude oxide, which still contains nickel and iron, is washed and ignited. It is thus converted into cobaltous dicobaltic tetroxide, { CoO000" in which form it is used in imparting a blue color to glass and porcelain. In order to obtain pure metallic cobalt, the commercial oxide is dis- solved in hydrochloric acid, and the solution evaporated to a small bulk. Arnmonic chloride and an excess of ammonia are then added. Any ferric hydrate which is precipitated is filtered off, and the solution is exposed to the air for some days until a portion of the liquid, when treated with excess of concentrated hydrochloric acid, does not become blue. Excess of concentrated hydrochloric acid is then added to the entire liquid, which is now heated to boiling and evaporated. Almost the whole of the cobalt separates as purpurea-cobalt chloride, Oo2Cl6,- in the form of a red crystalline powder. This, when heated in a current of hydrogen, is reduced to spongy metallic cobalt, which may be obtained in the form of a regulus by fusion in a crucible of lime or graphite. The oxides of cobalt are also reduced to the metallic state when heated in a current of hydrogen. Properties.—Metallic cobalt is almost white, with a faint reddish tinge, and is capable of taking a high polish. It is malleable and very tenacious. It is magnetic, and, unlike iron and nickel, is attracted by the magnet also when red hot. Its fusing-point lies somewhat lower than that of iron. The compact metal is oxidized neither in air nor in water at ordinary temperatures; but when heated in air it undergoes slow oxidation. It dissolves slowly in dilute sulphuric and hydrochlo- ric acids with evolution of hydrogen, and is readily soluble in dilute nitric acid. COMPOUNDS OF COBALT WITH THE HALOGEN. Cobaltous chloride, CoC12, is obtained by dissolving any of the oxides of cobalt in hydrochloric acid and evaporating. In the case of the oxides higher than cobaltous oxide the solution evolves chlorine. The concentrated liquid deposits dark-red monoclinic crystals of the formula CoC12,601I2. These, when heated to 120° C. (248° F.), are COMPOUNDS OF COBALT. 665 converted into a dark-bine crystalline powder possessing the formula CoC12,20H2, and at a temperature above 140° C. (284° F.), this salt be- comes anhydrous. The anhydrous salt sublimes in a current of chlorine, yielding dark-blue scales, which, when exposed to air, absorb moisture and become pink-colored. The anhydrous chloride dissolves slowly in water, yielding a pink-colored solution, and in absolute alcohol with a blue color, which becomes pink on the addition of water. Most cobal- tous salts exhibit this property of possessing a pink or rose-color in the highly hydrated condition, and a blue or violet color in the slightly hydrated or anhydrous condition. Owing to this property a solution of a cobaltous salt may be employed as a so-called sympathetic ink. Characters inscribed upon paper with a dilute solution of cobaltous chloride are invisible under ordinary conditions, but appear blue when the paper is warmed to expel the moisture, gradually disappearing again on cooling, owing to the absorption of moisture from the air. In like manner a not too dilute pink-colored solution of cobaltous chloride becomes blue on the addition of an excess of strong hydrochloric acid, owing to the abstraction of water from the salt in solution. Cobaltic chloride, /C02C16, is probably formed when cobaltic oxide is dissolved in cold hydrochloric acid, but the solution speedily evolves chlorine, and contains cobaltous chloride. Cobaltous bromide, Coßra, resembles the chloride in properties and mode of prepara- tion. The aquate, C08r2,60H2, is dark-red, the anhydrous salt green. Cobaltous iodide, Cola.—This compound is obtained by digesting finely divided cobalt with iodine and water. It forms either brownish-red prisms of the formula C01a,60 Ha, or small green very deliquescent crystalsof the formula C012.20H2. When heated to 130° C. (266° F.), the salt is converted into a black graphite-like mass of the anhydrous iodide. Cobaltous fluoride, CoF2,20H2, is obtained in rose-red crystals by dissolving the car- bonate in hydrofluoric acid and evaporating the solution. COMPOUNDS OF COBALT WITH OXYGEN. Cobaltous oxide, . . CoO. n r( O—Co—O Cobaltous dicobaltic te- f ,, r\ ri r/ r\ troxide, . . . .tCoOCo°- o=Co-Co=o. Cobaltic oxide, . . . / /\ I uou o=Co—Co=o. Cobaltous oxide, CoO, is formed when cobaltous hydrate or cobaltous carbonate is heated with careful exclusion of air. It is best prepared by strongly heating either of the higher oxides in a current of carbonic anhydride. It forms a greenish-brown powder, readily soluble in acids. When heated in hydrogen or carbonic oxide it is reduced to metal. Cobaltous hydrate, CoHo2.—On the addition of a caustic alkali to the solution of a cobaltous salt a blue basic salt is precipitated, which on boiling is converted into the rose-red hydrate. This, on exposure to air, speedily turns brown from oxidation. It is insoluble in caustic alkalies, but dissolves in ammonia with a reddish color. 666 INORGANIC CHEMISTRY. Cobaltous dicobaltic tetroxide, 'Co202Coo".—This compound is formed when either of the other oxides, or cobaltous nitrate, is strongly ignited in air. It forms a black non-magnetic powder. Cobaltic oxide, 'Co203, is prepared by gently igniting cobaltous nitrate. It is a dark-brown powder which dissolves in cold acids, yielding brown solutions of unstable cobaltic salts. On warming or evaporating the solutions decomposition ensues—in the case of the hydracids with evolution of halogen, in the case of the oxy-acids with evolution of oxygen—and a cobaltous salt remains in solution, Cobaltic hydrate, 'GoTIOg, is obtained as a black amorphous precipi- tate by adding an alkaline hypochlorite to the solution of a cobaltous salt. It behaves towards acids like cobaltic oxide. OXY-SALTS OF COBALT. Cobaltous nitrate, N2O4C00",60H2, forms red, very soluble deliques- cent monoclinic prisms. Cobaltous carbonate, COCoo//.—The anhydrous salt is obtained in bright-red micro- scopic octahedra by heating cobaltous chloride to 140° C. (284° F.) with a solution of hydric sodic carbonate which has been previously saturated with carbonic anhydride. An aquate of the formula COooo//2,60H2 is prepared by mixing asolution of cobaltous nitrate with the above solution of hydric sodic carbonate saturated with carbonic anhy- dride, and exposing the mixture for some time to a low temperature.—Normal alkaline carbonates precipitate from solutions of cobaltous salts blue or violet basic carbonates. Cobaltous sulphate, S02Coo".—This salt is prepared by dis- solving the oxide, hydrate, or carbonate in sulphuric acid. Its solutions deposit at ordinary temperatures dark-red monoclinic crystals of dihy- d.ric cobaltous sulphate, SOHo2Coo//,60H2, isomorphous with ferrous sulphate. The same salt occurs native as cobalt vitriol. Various other aquates are known.—Cobaltous sulphate forms with the sulphates of the alkalies double salts, which correspond exactly with the double sulphates of zinc, magnesia, etc., with the alkalies. Thus, dipotassic SO Ko cobaltous sulphate, SO2KoCoo'/j6OH2’ forms monoclinic crystals. Cobaltous phosphate.—The normal salt, P202C00//3, is obtained as a rose-red hydrated precipitate when hydric disodic phosphate isadded to thy solution of a cobaltous salt. —Hydric cobaltous phosphate, 2POHoCoo//,50H2, is prepared by dividing a quantity of the foregoing salt into two equal portions, dissolving the one portion in the smallest possible quantity of hydrochloric acid and then adding the other. It forms thin violet laminae. Cobaltous arsenate.—The normal salt, As202Coo//3,80H2, occurs native as cobalt-bloom or erythrine in peach-blossom-colored needles, or in earthy incrustations. This mineral has been formed by the spontaneous oxidation of speiss-cobalt and other native arsenites of cobalt. Zaffre is an impure basic arsenate of cobalt, prepared by roasting speiss- cobalt. It is employed in painting on glass and porcelain, for which purpose it must be free from iron. Silicates of Cobalt.—-These have not been prepared in a state of purity. When an alkaline silicate is added to the solution of a cobaltous salt a blue hydrated silicate of cobalt is precipitated. Smalt is a cobalt-potash glass'—a mixed silicate of cobalt and potassium. In a finely ground condition it is employed as a blue pigment. It is prepared on a large 667 COMPOUNDS OF COBALT. scale from speiss-cobalt or cobalt-glance. The ore is roasted at a low temperature, so as to oxidize the cobalt, leaving the nickel, iron, and other impurities, the presence of which would be detrimental to the purity of color of the smalt, as far as possible unaltered. The roasted ore is then fused with quartz-sand and potashes. The oxidized cobalt is taken up by the silica and unites with the potassic silicate to form smalt, whilst the nickel, iron, copper, bismuth, arsenic, etc., collect as a regulus at the bottom of the melting-pot. The glass is then finely ground under water. It contains from 6to 7 per cent, of cobalt and from 60 to 70 per cent, of silica. Smalt is less frequently employed as a pigment than formerly, owing to the introduction of artificial ultramarine; but it possesses the advantage over the latter pigment of not being altered by acids. Two other cobalt pigments are also manufactured: Thenard's blue or cobalt ultramarine, which is obtained by precipitating mixed solu- tions of alum and cobaltous sulphate with sodic carbonate and igniting the precipitate; and Pinmann’s green, which is prepared in a similar manner by igniting the precipitate produced by sodic carbonate in mixed solutions of cobaltous sulphate and zincic sulphate. Nothing is known concerning the constitution of these pigments. It has already been mentioned (p. 666) that the simple cobaltic salts are capable of existing only in solution. Double cobaltic salts are. however, known which possess a considerable degree of stability. Potassic cobaltic nitrite, N606(/Co"'206)vi,6NOKo.—This salt is formed as a yellow crystalline precipitate when potassic nitrite is added to the solution of a cobaltous salt acidified with acetic acid. Nitric oxide is evolved in the reaction ; 2CoC12 + IONOKo + INOHo = NG06('Co'"206)Ti,6NOK:o Cobaltous Potassic Nitrous Potassic cobaltic nitrite. Potassic Nitrous Potassic cobaltic nitrite. chloride. nitrite. acid. + 4KOI + 2'N"O + 20H2. Potassic Nitric Water. chloride. oxide. COMPOUNDS OF COBALT WITH SULPHUR. Cobaltous sulphide, CoS", is formed as a gray, metallic, crystal- line mass when cobalt is fused with sulphur. It may be obtained in long, thin, very lustrous needles of a yellowish-gray color by fusing a mixture of anhydrous cobaltous sulphate and baric sulphide with an excess of sodic chloride. Amnionic sulphide precipitates from solutions of cobaltous salts black amorphous hydrated cobaltous sulphide, scarcely soluble in cold dilute hydrochloric acid. Concentrated hydrochloric acid dissolves it with evolution of sulphuretted hydrogen. Other sulphides, and CoS"2, are obtained by heating cobalt- ous sulphide with sulphur in a current of hydrogen. Cobaltous dicobaltic tetrasulphide, Co2S"2Cos", occurs native in steel- gray or copper-red regular octahedra as the mineral cobalt pyrites. 668 INORGANIC CHEMISTRY. AMMONIUM CO3IPOUNDS OF COBALT {COBALTA- -311 NFS). The cobaltamines are of two classes—cobaltous and cobaltic. Their salts possess the empirical composition of additive compounds of one molecule of a cobaltous or a cobaltic salt with a certain number of molecules of ammonia. The salts of the first class are formed by the direct union of gaseous ammonia with anhydrous cobaltous salts. In the formation of the cobaltamines of the second class the oxygen of the air also plays a part. Thus the solution of a cobaltous salt in aqueous ammonia rapidly absorbs oxygen and is converted into a cobaltic ammo- nium base. Various bases belonging to this class are known. They all possess characteristic colors, and from these their names are derived. Cobaltosammonic chlorides.—Anhydrous cobaltous chloride absorbs dry ammonia gas, a. Cobaltous Ammonium Compounds. f NII3CI I nh3 | NH3 and is converted into the compound C0C12,6NH3 = -j Co// , which is thus ob- I net, I nh3 L nh3ci tained as a pale pink powder. The same compound is deposited in red octahedral crystals when the chloride is dissolved in concentrated aqueous ammonia and the solution allowed to stand in a well-stoppered bottle. When heated to 120° C. (248° F.) it parts with four molecules of ammonia and is converted into cobalioso-diammonic f N HjjCl dichloride, -I Co// ( NHgCI A nitrate of the empirical formula and a sulphate, S02Coo//,6NH3, have also been prepared. These may be divided into four principal series, of which the chlorides may serve as examples: b. Cobaltic Ammonium Compounds. Empirical formula. Dicohaltic-hexammonic (dichro-cobaltic) chloride, . /C02C166NH3. Dicobaltic-octammonic (praseo- and fusco-cobaltic) \ chloride, J 2 6’ 3‘ Dicobaltic-decammonic (roseo- and purpureo-cobaltic ] ™ ] qwjt chloride, / 2 6’ 3‘ Dicobaltic-dodecammonic (luteo-cobaltic) chloride, /Co2C16,12NH3. The coZor-names (see above) are given in brackets. It will be observed that some of these bases exist in isomeric modifications. The above compounds behave like chlorides of complex ammonium bases. Thus the chlorine may be replaced by hydroxyl, and the resulting compounds are hydrates possessing an alkaline reaction and a purely alkaline, as opposed to a metallic, taste. Again, the chlorides form double compounds with platinic and auric chlorides. r-NUjCI f Co-NHjCI I j QJ Dichro-cobaltic chloride {Dicobaltic hexammonic chloride) -{ nITCI This [ Co-nhJgi I—NHSCI compound is formed wdien a solution of cobaltous chloride in aqueous ammonia is exposed to the air until the separation of black cobaltic hydrate commences. On COMPOUNDS OF COBALT. 669 adding an excess of concentrated hydrochloric acid and allowing the liquid to stand for some time, the cobaltamine chloride is deposited in dark-colored laminse or feather-shaped crystals. The dichroism of this compound is best exhibited by break- ing a large crystal; the splinters display different colors. 2(NvH4)CI f Co-NH3CI Proseo-cobaltic chloride, j ,20112. This compound is generally [ Co-NHjCI >_NII2(N7II4)CI formed along with the preceding and other cobaltamines, remaining in the filtrate after these have been precipitated with concentrated hydrochloric acid. On saturat- ing the liquid with amnionic chloride the praseo-compound separates in bright-green glistening needles. I—NH2(NvH4)CI1—NH2(NvH4)C1 f Co-NH2(N7H4)CI i nh. Cl Roseo-cobaltic chloride, - Nfh Cl ,20H2.—If a solution of cobaltous chlo- L Co-niiJnut4)ci 2(NtH4)CI ride in aqueous ammonia be oxidized by exposure to the air until a sample on testing with excess of strong hydrochloric acid no longer assumes a blue color, the liquid contains the roseo-compound. This may be separated by supersaturating the solution with strong hydrochloric acid, carefully avoiding any rise of temperature when the roseo-salt is deposited as a brick-red powder. The two molecules of water, which in the above formulae are represented as water of crystallization, are in reality water of constitution, inasmuch as they cannot be expelled without converting the compound into purpureo-cobaltic chloride, a salt which, though differing totally in its properties from the roseo-salt, possesses the same chemical composition, excepting that it is anhydrous. The purpureo-salt cannot be converted into the roseo-salt merely by recrystallizing from water. The dry roseo-salt slowly changes at ordinary tempera- tures into the purpureo-salt. This change takes place more rapidly in solutions, and on boiling is practically instantaneous. A number of other roseo-salts have been prepared. Purpureo-cobaltic chloride.—This compound possesses, as above stated, the same com- position as the foregoing, less two molecules of water. It is obtained by the same pro- cess as the roseo-salt, except that after supersaturating with strong hydrochloric acid the liquid is heated to boiling. The red powder which separates is purified by re- crystallization from hot dilute hydrochloric acid. The compound is thus obtained in small purple crystals. It may be converted into the roseo-compound by dissolving in dilute aqueous ammonia and adding the solution drop by drop to carefully-cooled strong hydrochloric acid. The salts of the purpureo-base with the various other acids have also been pre- pared. i—NH2(NTII4)C1 f Co-NH2(NvH4)CI lAiteo-cobaltic chloride, -j XnH, (NvII4jci •—•method of preparing the luteo- or [ Co-NH2(N*H4)CI l_nH2(NtH4)CI dodecamine-compounds yielding perfectly certain results has yet been discovered. They are formed along with the other cobaltamines in the oxidation of ammoniacal solutions of cobaltous salts, especially in presence of ammonic chloride, and must be separated from these by systematic crystallization. Luteo-cobaltic chloride crystallizes in reddish-yellow monoclinic prisms. The above list includes only the principal cobaltamines. Various other complex bases of this class have been prepared. General Properties and Reactions of the Compounds of Cobalt.—The aquates of the cobaltous salts are usually red ; the an- hydrous salts are blue. With caustic alkalies their solutions yield in the cold a blue precipitate of a basic salt, which on boiling is converted into pink cobaltous hydrate. Ammonia produces a similar precipitate 670 INORGANIC CHEMISTRY. soluble in excess, yielding a reddish solution which absorbs oxygen from the air and becomes reddish-brown. In presence of salts of am- monia no precipitate is produced on addition of ammonia. Sulphu- retted hydrogen gives no precipitate in presence of strong acids; ammonia sulphide precipitates black hydrated cobaltous sulphide, insoluble in alkalies and alkaline sulphides, scarcely soluble in dilute hydrochloric acid, readily soluble in aqua-regia. Potassic ferrooyani.de gives a green precipitate of cobaltous ferrocyanide (Co"2Fe"Cy6), and potassia ferri- cyanide a reddish-brown precipitate of cobaltous ferricyanide (Co"3- 'Fe"'2Cy12). Potassia cyanide precipitates pale-brown cobaltous cyanide, which dissolves in an excess of the alkaline cyanide, yielding a double cyanide of potassium and cobalt. From this solution acids precipitate cobaltous cyanide. If, however, to the solution containing the double cyanide, together with an excess of potassic cyanide, a small quantity of hydrochloric acid insufficient to cause a precipitate be added, and the liquid be boiled, potassic cobalticyanide is formed, and in the solution of this salt neither acids nor amraonic sulphide occasion a precipitate. (Distinction between the compounds of cobalt and nickel.) All the compounds of cobalt when heated with sodic carbonate on char- coal in the reducing flame yield shining white metallic particles which are attracted by a magnet. Cobalt compounds color the borax and microcosmic salt beads deep-blue, both in the oxidizing and in the re- ducing flame. They do not yield a flame-spectrum. NICKEL, Ni. Atomic weight = 58.6. Molecular weight unknown. Sp.gr. 8.9. Atom- icity ", IV, and T1 ? Also a pseudo-triad. Evidence of atomicity: Nickelous chloride, .... .... Ni"CL. Nickelic disulphide, .... .... NiivS"2. Nickelic oxide, / Ni"'On • • • ’ \ History.—Metallic nickel was first obtained by Cronstedt in 1751. Occurrence.—Nickel occurs in the native state in meteoric iron, of which it is an invariable constituent. Its chief ores are its compounds with arsenic, antimony, and sulphur; and in these it is generally asso- ciated with cobalt. Kupfer nickel, so called from its copper-red color, f AsNi is a dinickelous diarsenide, < ; this is the most important ore of f As the metal. A nickelous diarsenide," AsNi’ also occurs as arsenical nickel. Other minerals are millerite or nickel blende, a nickelous sul- f As phide, NiS"; nickel glance, a snlph-arsenide, < (Ni"S")"2; and f SbNi breithauptite, a dinickelous diantimonide, < In New Caledonia a source of nickel has lately been discovered in the mineral garnierite, a nickelous silicate of the formula 2Si403Nio"5,30H2, which occurs NICKEL. 671 there in large quantities. This ore is remarkable as being free from cobalt. Nickel has been detected in the solar atmosphere. Extraction.—The process of extracting nickel from its ores is iden- tical with that employed in the extraction of cobalt (p. 663) up to the point at which the cobalt is precipited as cobaltic oxide by bleaching- powder. From the solution thus freed from cobalt the nickel is pre- cipitated as hydrate by the addition of milk of lime. The precipitate is ignited and afterwards, in order to remove the excess of lime, treated with dilute hydrochloric acid, in which the ignited oxide of nickel is in- soluble. The purified oxide is reduced by heating with carbon. Pure nickel may be prepared by heating pure nickelous oxalate with exclusion of air. The metallic powder thus obtained may be fused into a regulus in a lime crucible. | CQNio" = Ni + 2C02. Nickelous oxalate. Carbonic anhydride. Properties.—Nickel is almost silver-white, with a faint yellowish tinge. It is capable of taking a high polish. It is very hard, but at the same time malleable and ductile. Nickel fuses at a somewhat lower temperature than cobalt. It is attracted by the magnet, but loses this property at a high temperature. It is not oxidized either in air or water at ordinary temperatures, and is oxidized only with difficulty when heated in air. It decomposes steam slowly at a red-heat and is converted into nickelous oxide. It dissolves slowly in dilute hydro- chloric and sulphuric acids, but is readily soluble in dilute nitric acid. Concentrated nitric acid renders it “ passive ” like iron. The commercial metal contains carbon along with traces of cobalt, iron, copper, and other metals. The presence of carbon has the same effect upon nickel as upon iron : it diminishes the malleability and low- ers the fusing-point of the metal. Nickel plating.—Nickel may be electrolytically deposited in a coher- ent coating from a solution of pure diammonic nickelous sidphate, SQ2 ImoM A plate of pure nickel serves as the positive elec- trode. Iron and steel are frequently coated with nickel, both on ac- count of the beauty and permanence of the metallic surface thus obtained, and also as a protection against rust. Alloys of nickel.—Nickel yields with copper valuable alloys of a silver-white color. The material of the small coinage in the United States, in Germany, in Belgium, in Switzerland, and in Brazil, is an alloy of 1 part of nickel with 3of copper. As this alloy is more val- uable than copper, the coins are smaller and consequently more port- able than copper coins possessing an equal value, whilst, owing to the hardness of the alloy, this coinage is also very durable. Chinese pack- fong is an alloy of copper, nickel, and zinc. German silver or nickel silver is a similar alloy, consisting, as a rule, of 5 parts of copper, 2 parts of nickel, and 2 parts of zinc. When first prepared it is crystalline and brittle; but by rolling and hammering, heating and allowing to cool, it is rendered tenacious and malleable. 672 INORGANIC CHEMISTRY. COMPOUNDS OF NICKEL WITH THE HALOGENS. Nickelous chloride, NiCl2, is obtained as a yellow earthy mass by dissolving the oxide or the carbonate in hydrochloric acid and evap- orating the solution to dryness. It may be sublimed in a current of chlorine, and is thus obtained in lustrous golden-yellow laminae. It dissolves in water, yielding a green solution which deposits on evapo- ration green monoclinic prisms of the formula NiCl2,60H2. Nickelous bromide, Nißr2, is prepared by heating finely divided nickel in bromine vapor. Combination occurs with incandescence, and the nickelous bromide sublimes in golden-yellow scales. The compound deliquesces in moist air. The green aqueous solution deposits on evaporation deliquescent needles of the aquate Nißr2,30H2. Nickelous iodide, Nil2, is obtained in a similar manner by heating spongy nickel in iodine vapor. It forms black lustrous laminae which dissolve in water, yielding a green solution. The aquate Nißr2,60H2 forms bluish-green deliquescent prisms. Nickelous fluoride, NiF2, is prepared by evaporating the solution of the carbonate in hydrofluoric acid. Bluish-green crystals of the formula NiF2,30H2 are deposited, which on boiling with pure water are decomposed with separation of an insoluble oxy- fluoride. COMPOUNDS OF NICKEL WITH OXYGEN Nickelous oxide, . . . NiO /°\ ( NiO / \ Nickelic oxide, . . . j o=Ni—Ni=o. Nickelous oxide, NiO, occurs native as the rare mineral bunsenite in green, translucent, regular octahedra. It may be obtained artificially in crystals by heating a mixture of nickelous sulphate and potassic sul- phate to a high temperature. In the crystallized condition it is with difficulty attacked by acids. By igniting the hydrate or carbonate it is obtained as a gray amorphous powder, readily soluble in acids. Nickelous hydrate, NiHo2, is an apple-green precipitate, obtained by adding caustic alkalies to the boiling solution of a nickelous salt. The precipitate is washed with hot water and dried. Acids dissolve it readily. It is insoluble in potassic hydrate and sodic hydrate, but ammonia dissolves it, yielding a blue solution, from which it is repre- cipitated as a green crystalline powder on expelling the ammonia by boiling. Nickelic oxide, 'NijOg, is prepared by careful ignition of the nitrate. It is a black powder which dissolves in hydrochloric acid with evolution of chlorine, and in sulphuric acid with evolution of oxy- gen. Ammonia dissolves it with evolution of nitrogen. 3'NLjO,, + 2NH3 = GNiHo2 + N2 + 30H2. Nickelic oxide. Ammonia. Nickelous hydrate. Water. Nickelic hydrate, 'NiglTog.—This compound is obtained as an amor- phous black powder when chlorine is passed through water (or prefer- ably through a solution of an alkaline hydrate or carbonate) in which nickelous hydrate is suspended ; or by warming a solution of a nickelous salt with an alkaline hypochlorite. Towards acids and ammonia it behaves like nickelic oxide. COMPOUNDS OF NICKEL. 673 Nickdous nitrate, N2O4Nio//,60H2, is obtained by dissolving the metal, the oxide, or the carbonate in nitric acid. It crystallizes in deliquescent green raonoclinic prisms. OX Y-SALTS OF NICKEL Nickdous nitrite, N2O.2Nio//, is prepared by decomposing nickelons sulphate with baric nitrite and evaporating the filtrate over sulphuric acid. It forms reddish-yellow crystals, which, when dry, may be heated to 100° C. without decomposition, but in solution are decomposed at 80° C. (176° F.) with separation of a basic salt.—It forms with the nitrites of other metals double salts: thus potassie nickdous nitrite, N202Nio//,- 4NOKo, and baric nickelons nitrite, N2o2Nio//,2N202Bao//. On adding potassie nitrite to the mixed solution of a nickel salt with a salt of barium, strontium, or calcium, the triple salts, NANio" N.CVBao" 2NOKo ; N?o,Nio" N2O*Sro",2NOKo; and Na0INio//,NaO,Cao//,2N0K0, are precipitated as sparingly soluble yellow crystalline powders, consisting of minute octahedra. These salts closely resemble in appearance potassie cobaltic nitrite. Owing to the formation of these salts it is not possible, in presence of the metals of the alkaline earths, to separate cobalt from nickel by means of potassie nitrite. Nickdous carbonate.—The anhydrous salt, CONio//, forms pale-green microscopic octahedra; the aquate, CONio//,60H-2, crystallizes in minute rhombohedra or prisms Both are obtained like the corresponding cobalt compounds (p. 666). Nickelous sulphate.—This salt is obtained by dissolving metallic nickel, or its oxide, hydrate, or carbonate, in dilute sulphuric acid. At ordinary temperatures it crystallizes from neutral solutions in green rhombic prisms of the formula SOHo2Ni0",60H2, isomorphous with magnesic sulphate. At temperatures between 30° and 40° C. (86-104° F.), or from solutions containing an excess of acid, bluish-green quadratic pyramids of the formula S()Ho2Nio;/,50H2, are obtained. This aquate is also deposited at temperatures above 50° C. (122° F.), but in monoclinic forms. 100 parts of water at 16° C. (60° F.) dis- solve 37.4 parts of anhydrous salt.—Anhydrous nickelons sulphate absorbs dry ammonia and is converted into a violet-white powder of the composition S02Nio",6NH3.—Nickelons sulphate forms double salts with the sulphates of the alkali metals and ammonia. Diammonic nickdous sulphate, gQ2^™^Nio,/,60H2, a salt employed in nickel electro-plating (p. 671), is prepared by adding amnionic sulphate to a concentrated solution of pure nickelons sulphate. The double salt separates as a crystalline powder and is purified by recrystallization. Nickdous phosphate, P2o2Nic//3, is formed as a pale-green hydrated precipitate when hydric disodic phosphate is added to the solution of a nickel salt. On heating, it becomes anhydrous and turns yellow. Nickelons arsenate, occurs native as the mineral 'nickel bloom, in green capillary crystals or as an efflorescence. COMPOUNDS OF NICKEL WITH SULPHUR. Nickelous sulphide, NiS", occurs native as miUerite or capillary pyrites in brass-yellow hexagonal and generally capillary crystals. It is formed when nickel and sulphur are heated together. Ammonic sulphide precipitates from solutions of nickel salts black hydrated 674 INORGANIC CHEMISTRY. amorphous niokelous sulphide, and if yellow aramonic sulphide has been employed, an excess of* this precipitant dissolves a portion of the precipitate, yielding a brown solution. The precipitate is slowly oxi- dized by exposure to the air when moist. The precipitated compound dissolves with difficulty in hydrochloric acid; and this acid is totally without action upon the native sulphide or upon that prepared in the dry way. Nickelic disulphide, NiS//2, is obtained by fusing a mixture of nickelous carbonate and sodic carbonate with an excess of sulphur. On extracting the mass with water the disulphide remains as a dark iron-gray impalpable powder. General Properties and Reactions of the Compounds of Nickel.—The aquates of the salts of nickel are of an apple-green color; the anhydrous salts are yellow. Caustic alkalies precipitate pale-green nickelous hydrate, which is not altered either by boiling or by exposure to air. Ammonia gives a similar precipitate, soluble in excess, yield- ing a greenish-blue liquid ; in presence of salts of ammonia no precipi- tate is formed. Sulphuretted hydrogen produces no precipitate in solu- tions with strong acids ; ammonia sulphide precipitates black hydrated nickelous sulphide, slightly soluble in excess, yielding a brown solution. The sulphide is scarcely soluble in dilute hydrochloric acid, readily soluble in aqua-regia. Potassic ferr o cyanide precipitates greenish-white nickelous ferrocyanide (Ni//2Fe//Cy6); potassic ferricyanide precipitates yellowish-brown nickelous ferricyanide (Ni///3/Fe/,/2Cyl2). Potassic cyanide produces a yellowish-green precipitate of nickelous cyanide, soluble in excess of the precipitant with formation of a double salt. From this solution acids reprecipitate nickelous cyanide, and if the solution be warmed with sodic hypochlorite the nickel is precipitated as black hydrated nickelic oxide. (Cobalt is not precipitated under these circumstances by sodic hypochlorite). The compounds of nickel, when heated with sodic carbonate on charcoal in the reducing flame, yield white shining magnetic particles of metallic nickel. With borax and microcosmic salt the compounds of nickel yield characteristically colored fluxes. In the oxidizing flame the borax bead is violet while hot, reddish-brown when cold ; the microcosmic salt bead is red or red- dish-brown while hot, yellow, or reddish-yellow when cold. In the reducing flame the microcosmic salt bead undergoes no change, whilst the borax bead turns gray and clouded, owing to the separation of metallic nickel. The nickel compounds do not color flame. NORWEGIUM, Ng. Atomic weight = 214 ? Sp. gr. 9.441. Fuses at 254° C. (489° F,). This rare metal has been recently discovered by Dahll in a specimen of Norwegian nickel glance. Very little is yet known concerning it. In most of its properties it closely resembles bismuth, but differs from this element in the solubility of its oxide in an excess of potassic hydrate, or of alkaline carbonates, on boiling. Assuming the correctness of the above atomic weight, the oxide possesses the formula Ng203. Excess of water decomposes its salts with precipitation of basic compounds. INDEX *** In order that names of compounds may as far as possible appear under the headings of their respective elements, the numerical prefixes di, tri, etc., have been omitted in the Index, except in cases where they serve to distinguish compounds that might otherwise he confounded. Abratjmsalz, oil Absorption of gases by charcoal, 196 Accumulators, electrical, 106 Acid, arsenic, 372 Acid, metarsenic, 373 metastannic, 327 metatungstic, 626 molybdic, 621 muriatic, 156 arsenions, 371 auric, 554 bisnlphuretted hyposulplmric, 280 boracic, 191 boric, 191 bromic, 294 carbonic, 208 chlorhydric, 156 chloric, 181 chlorochromic, 638 chromic, 634 nitric, 214 nitrous, 223 Nordhausen sulphuric, 274 orthantimonic, 387 ortharsenic, 373 orthoboric, 191 orthophosphoric, 356 osmic, 601 parantimonic, 387 pentathionic, 281 perchloric, 183 perchromic, 633 periodic, 304 permanganic, 648 phosphomolybdic, 622 phosphoric, 356 phosphorosophosphoric, 358 phosphorous, 351 platinonitrous, 590 pyrantimonic, 387 pyrarsenic, 373 pyrophosphamic, 364 pyrophosphodiamic, 364 pyrophosphoric, 355 pyrophosphotriamic, 363 pyrosulphuric, 274 selenic, 287 selenious, 286 silicic, 318 silicon-oxalic, 313 silico-tungstic, 627 stannic, 327 sulphhydric, 249 sulphindic, 563 sulphocarbonic, 258 sulphodithionic, 279 sulphuretted hyposulphuric, 279 sulphuric, 267 sulphurous, 262 telluric, 289 tellurous, 288 chromous, 634 disulphodithionic, 280 dithionic, 278 dithionous, 278 ferric, 660 graphitic, 199 hydriodic, 298 hydrobromic, 292 hydrochloric, 156 hydrolluoboric, 190 hydrofluoric, 307 hydrofluosilicic, 315 hydroselenic, 285 hydrosulphuric, 249 hydrosulphurous, 278 h vpobromous, 294 hypochlorous, 179 hyponitrous, 221 hypophosphoric, 358 hypophosphorous, 350 hyposulphuric, 278 hyposulphurous, 276 iodic, 302 manganic, 647 metabismuthic, 395 metabismuthous, 394 metaboric, 192 metantimonic, 387 metantirnonic, of Fremy, 387 metantimonious, 885 metaphosphoric, 353 676 INDEX Acid, tetrathionic, 280 thiosulphuric, 276 titanic, 332 trisulphodithionic, 281 trisulphuretted hydrosulphuric, 281 trithionic, 279 tungstic, 625 tungstic, colloidal, 626 tungsto-silicic, 627 vanadic, 366 Acids, definition of, 40 JEthiops mineralis, 535 Affinity, chemical, 34, 102 After-damp, 203 Agalmatolite, 572 Air, 237 analyses of, 239 not a compound, 242 Alabaster, 477 Albite, 572 Aldebaran, elements detected in, 406 Alkali waste, 244 Allophane, 572 Allotropy, 110 Alloys, 410 Alum, 569 shale, 570 stone, 570 Alumina, 567 Aluminates, 568 Aluminic bromide, 566 chloride, 566 fluoride, 566 hydrate, 568 hydrate, collodial, 568 iodide, 566 manganous sulphate, 647 nitrate, 568 Ammoniacal platinum compounds, 591 Ammonic borate, 446 bromate, 443 bromide, 441 carbonate, 443 chlorate, 443 chloride, 441 chlorostannate, 326 chromate, 637 dichromate, 637 di-iridic chloride, 596 dithionate, 445 ferrous sulphate, 659 fluoride, 442 heptasnlphide, 446 hydrate, 442 hyposulphite, 445 indie sulphate, 563 iodate, 443 iodide, 442 iridic chloride, 597 magnesic chromate, 637 nickelous sulphate, 673 nitrate, 442 nitrite, 443 palladic chloride, 594 pentasulphide, 446 perchlorate, 443 permanganate, 649 phosphate, 445 phosphomolybdate, 622 platinic chloride, 441 platinonitrite, 590 potassic sulphate, 444 pyrophosphate, 445 pyrosulphite, 445 silicofluoride, 442 sodic phosphate, 445 sodic sulphate, 444 sulphate, 444 sulphhydrate, 446 sulphide, 446 sulphite, 444 thiosulphate, 445 tungstate, 627 uranate, 618 Ammonium, 86, 235 amalgam, 235 oxide, 567 oxydihydrate, 568 oxytetrahydrate, 568 phosphate, 571 silicate, 571 sodic fluoride, 566 sulphate, 569 sulphide, 576 Aluminio-sodic fluoride, 426 Aluminite, 569 Aluminium, 564 bronze, 565 general properties and reactions of the compounds of, 576 Alums, 569 general properties and reactions of the salts of, 446 salts of, 440 Ammonoxyl, 86 Analcime, 572 Alunite, 570 Amalgamation process for extraction of silver, 448 Amalgams, 529 Amidogen, 86 Ammonia, 230 alum, 571 chrome alum, 634 gallium alum, 577 Ammonia-soda process, 429 Ammoniacal cobalt compounds, 668 mercury compounds, 536 Anatase, 332 Andalusite, 572 Anglesite, 612 Anhydride, antimonic, 386 antimonious, 384 arsenic, 372 arsenious, 370 auric, 554 bismuthic, 394 boracic, 190 boric, 190 carbonic, 200 677 INDEX Anhydride, carbonic, decomposition of by plants, 164 chlorous, non-existence of, 177 chromic, 632 hypochlorous, 177 hyponitrous, 220 iodic, 301 molybdic, 621 nitric, 219 nitrous, 222 Apatite, 357, 479 Apjohnite, 646 “ Aq,” use of symbol, 431 Aquafortis, 214 Aquamarine, 521 Aqua-regia, 218 Aquates, 45 Aqueous vapor, 240 Argentic amide, 459 ars dilate, 458 arsenite, 458 bromate, 457 bromide, 453 carbonate, 457 osmic, 601 permanganic, 645 persulphuric, 276 phosphoric, 266 phosphorous, 351 selenious, 286 silicic, 316 silicoformic, 314 stannic, 326 sulphantimonic, 389 sulphantimonious, 388 sulpharsenic, 376 sulpharsenious, 374 sulphuric, 265 sulphurous, 260 telluric, 289 tellurous, 288 titanic, 332 tungstic, 625 uranic, 616 vanadic, 365 Anhydrides, definition of, 40, 42 Anorthite, 319 Anthracite, 198 Anthrosiderite, 661 Antimonic chloride, 382 fluoride, 383 chlorate, 456 chloride, 452 chromate, 638 dichromate, 638 dithionate, 457 fluoride, 454 hyposulphite, 457 iodate, 457 iodide, 453 metaphosphate, 458 nitrate, 456 nitrite, 456 orthophosphate, 458 oxide, 454 periodate, 457 permanganate, 649 peroxide, 455 phosphate, 458 phosphide, 459 pyrophosphate, 458 sulphantimonite, 459 sulpharsenite, 459 sulphate, 457 sulphide, 459 sulphite, 457 thiosulphate, 457 Argentite, 459 Argentous oxide, 454 chloride, 453 Argillaceous iron ore, 651 Arragonite, .477 Arsenates, 373-376 Arsenic, 366 fluoride, 370 general properties and reactions of the compounds of, 376 poisoning, antidote for, 371 sulphide, 376 Arsenical iron, 366 oxytrichloride, 383 sulphide, 389 sulphotrichloride, 883 tetrethochloride, 378 Antimonious amylide, 381 argentide, 381 bromide, 383 chloride, 381 elhide, 381 fluoride, 383 hydride, 381 iodide, 383 oxide, 384 oxybromide, 383 oxychloride, 382 oxyfluoride, 383 oxyiodide, 383 sulphide, 388 zincide, 381 Antimoniuretted hydrogen, 380 Antimony, 378 amorphous, 379 copper glance, 389 crystalline, 378 general properties and reactions of, 390 nickel, 670 pyrites, 366 Arsenious bromide, 370 chlorhydrate, 369 chloride, 369 fluoride, 370 hydride, 367 iodide, 370 sulphide, 375 sulphide, colloidal, 375 Arsenites, 372, 376 Arseniuretted hydrogen, 367 ochre, 378 Antimonylic autimonate, 385 678 INDEX, Artiads, 79 Atacamite, 544 Baric oxide, 462 perchlorate, 465 Atmosphere, 237 composition of, 239 weight of, 238 Atom, definition of, 59 “ Atomic analogues,” 94 Atomic heat, 68 theory, 48 weight, definition of, 61 t weight, determination of by Avoga- dro’s law, 61 weight, determination of by means of isomorphism, 64 weight, determination by Neumann’s law, 71 weight, determination of by means of specific heat, 67 weights, list of, 38 volume, 96 volumes, curve of, 95 Atomicity, 78 active, 81 absolute, 81 permanganate, 649 peroxide, 463 platinate, 590 pyrosulphate, 466 silicofluoride, 462 sulphhydrate, 467 sulphate, 465 sulphide, 467 sulphite, 466 tetrasulphide, 467 thiosulphate, 466 Barium, 460 amalgam, 460 general properties and reactions of the compounds of, 468 Baryta, 462 water, 464 Bases, definition of, 43 Batteries, secondary, 106 storage, 106 Bauxite, 568 latent, 81 law of variation of, 80 of elements, 88 Atoms, 48 nature of, 51 Auric arnrnonic chloride, 554 chloride, 553 hydrate, 555 oxide, 554 potassic chloride, 554 sodic chloride, 554 Aurous amnionic sulphite, 555 chloride, 553 Bell metal, 542 Berthelot, laws of thermochemistry, 111 Berthierite, 389 Beryl, 521 Beryllia, 522 Beryllic aluminate, 568 bromide, 522 carbonate, 523 chloride, 521 fluoride, 522 hydrate, 522 iodide, 522 nitrate, 523 oxide, 522 iodide, 553 oxide, 554 sodic thiosulphate, 555 sulphide, 556 Avogadro’s law, 53 apparent exceptions to, 63 Azote, 211 Azurite, 547 phosphate, 523 silicate, 523 sulphate, 523 sulphide, 523 Beryllium, 521 general properties and reactions of the compounds of, 523 Bessemer process of steel making, 653 Bismuth, 391 general properties and reactions of the compounds of, 396 glance, 396 ochre, 393 Baking porcelain, 574 Baric bromide, 461 carbonate, 465 chlorate, 465 chloride, 461 chromate, 637 dichromate, 637 dithionate, 466 ferrate, 661 fluoride, 462 hydrate, 463 iodide, 462 manganate, 648 nickelous nitrate, 673 nitrate, 464 nitrite, 465 orthophosphate, 466 telluric, 396 Bismuthous bromide, 392 chloride, 391 dichlorethide, 391 ethide, 391 fluoride, 392 iodide, 392 nitrate, 394 nitrate dihydrate, 393, 394 oxide, 393 oxide, salts of, 394 oxybromide, 392 oxychloride, 392 oxyhydrate, 394 osmate, 602 oxyiodide, 392 679 INDEX Bismuth ous sulphide, 396 telluride, 396 uranate, 618 Bitter-spar, 510 Black ash, 429 Black band, 651 Black-lead, 199 Blanc fixe, 466 Bleaching, 476 Bleaching powder, 181, 476 Blister copper, 540 Blue malachite, 547 Blue vitriol, 547 Boiling points, 119 influence of pressure upon, 120 method of determining, 121 relation of to molecular weight, 121 Cadmic hydrate, 525 iodide, 525 nitrate, 525 oxide, 525 sulphate, 526 sulphide, 526 Cadmium, 524 amalgam, 530 general properties and reactions of the compounds of, 526 Classic antimonious chloride, 440 carbonate, 440 chloride, 440 hydrate, 440 nitrate, 440 platinic chloride, 440 sulphate, 440 Caesium, 439 general properties and reactions of the compounds of, 440 Calaite, 571 Calamine, 518 siliceous, 519 Calcic bromide, 473 carbonate, 477 Bolognian phosphorus, 467 Bonds, 78 Boracite, 512 Borates, 192 Borax, 434 Boric bromide, 189 chloride, 188 ethide, 185 fluoride, 189 chlorate, 475 chloride, 472 chlorohypochlorite, 181, 476 chlorophosphate, 335 chromate, 637 dithionate, 478 fluoride, 473 hydrate, 474 hypochlorite, 475 hypophosphite, 480 iodide, 473 iodohypiodite, 297 nitrate, 476 nitrite, 476 hydride, 187 nitride, 187 sulphide, 193 Borofluorides, 190 Boron, 185 adamantine, 185 amorphous, 186 graphitoid, 185 Boulangerite, 389 Boyle, law of, 52 Bournonite, 389 Bracket, use of, 76 Brass, 541 Braun ite, 643 Breithauptite, 670 Britannia metal, 323 Brittleness, 408 Brochantite, 547 Bromargyrite, 453 Bromates, 295 Bromides, 293 Bromine, 290 orthophosphate, 478 oxide, 474 oxychlorhydrate, 473 peroxide, 474 phosphate, 479 phosphide, 344, 483 potassic sulphate, 478 silicates, 480 silicofluoride, 473 sodic sulphate, 478 sulphate, 477 sulphide, 483 sulphite, 478 thiosulphate, 478 tungstate, 627 Calcined magnesia, 509 Calcite, 477 Calcium, 471 general properties and reactions of the compounds of, 484 Calc-spar, 477 Calomel, 530 Calorie, 68 Capillary pyrites, 673 Carat, definition ofj 553 hydrate, 291 Bronze, 542 Brookite, 332 Brown haematite, 658 Brown iron ore, 658 Brucite, 509 Brunswick green, 544 Brushite, 479 Bucholzite, 572 Bunsenite, 672 Butter of antimony, 382 Cadmic bromide, 525 carbonate, 525 chloride, 525 Carbon, 193 680 INDEX, Carbon, bisulphide of, 256 circulation of in nature, 202 Carbonates, 207 Carbonic disulphide, 256 oxide, 208 oxide, compound of with potassium, 210 Chrome ochre, 631 orange, 637 red, 637 yellow, 637 Chromic bromide, 630 chloride, 630 dioxide, 634 fluoride, 630 hydrate, 632 hydrate, colloidal, 632 nitrate, 634 nitride, 639 oxide, 631 oxychlorhydrate, 638 oxydichloride, 638 perfluoride, 631 sulphate, 634 sulphide, 639 Chromites, 634 Chromium, 629 oxydichloride, 211 oxysulphide, 258 Carbonylic chloride, 211 Carnallite, 508 Cassel yellow, 607 Cast iron, 652 Caustic potash, 415 soda, 427 Celestine, 47 0 Cementation process of steel making, 653 Ceric fluoride, 580 hydrate, .580 nitrate, 580 oxide, 580 sulphate, 581 Cerite, 578 Cerium, 578 Cerous chloride, 580 fluoride, 580 hydrate, 580 nitrate, 580 oxide, 580 general properties and reactions of the compounds of, 639 Chromosphere, 405 Chromous bromide, 630 chloride, 630 chromic oxide, 633 hydrate, 631 oxide, 631 phosphate, 633 sulphate, 633 Ohromylic chi or hydrate, 638 chloride, 638 Chrysoberyl, 568 Chrysocolla, 549 Cimolite, 572 phosphate, 581 potassic sulphate, 581 Cervantite, 385 Chalcedony, 319 Chalk, 477 Charcoal, 194 absorption of gases by, 196 animal, 195 Charles, law of, 53 Chemical action, modes of, 102 affinity, 102 combination, heat of, 111 equations, 78 formulae, 75 homogeneity, 108 nomenclature, 39 notation, 75 Chiastolite, 572 Chili saltpetre, 427 China, 574 China clay, 572, 573 Chlorates, 182 Chloraurates, 554 Cinnabar, 535 Clay, 573 Clay iron-stone, 651 Coal, 197 Coal-gas, purification of, 245 Coarse metal, copper, 539 Cobalt, 663 ammonium compounds of, 668 bloom, 666 general properties and reactions of the compounds of, 669 pyrites, 667 ultramarine, 667 vitriol, 666 Cobaltamines, 668 Cobaltic chloride, 665 hydrate, 665 oxide, 666 Cobaltosammonic chloride, 668 Cobaltoso-diammonic dichloride, 668 Cobaltous arsenate, 666 bromide, 665 carbonate, 666 Chloride of lime, 476 Chlorine, 151 hydrate, 154 oxygen compounds of, 177 Chloric peroxide, 178 Ohlorochromates, 639 Chloronitrous gas, 228 Chloropal, 320 Chloropernitric gas, 229 Chlorophyll, iron in, 651 Chromates, 635 Chrome alum, 634 chloride, 664 dicobaltic tetrasulphide, 667 dicobaltic tetroxide, 665 fluoride, 665 hydrate, 666 iodide, 665 iron ore, 635 nitrate, 666 INDEX Cobaltous oxide, 665 phosphate, 666 silicate, 666 sulphate, 666 sulphide, 667 Cohesive power, 407 Coke, 197 Cupric oxide, 546 oxychloride, 544 phosphate, 648 phosphide, 342, 550 ■ silicate, 549 silicide, 312 sulphate, 547 sulphide, 549 Cuprosammonic chloride, 543 Cuprous acetylide, 543 arsenide, 550 bromide, 543 Colloidal sulphides, 549 Colloids, 130 Colly rite, 572 Combination, 112 atomic, 87 by volume, 54 laws of, 45 molecular, 87 Combustibles, 165 chloride, 542 fluoride, 544 hydrate, 546 hydride, 542 iodide, 543 nitride, 550 oxide, 544 phosphide, 550 quadrantoxide, 544 sulphide, 549 Cuttle-fish, copper in blood of, 538 Cyanite, 572 Combustion, 164 supporters of, 165 Compound radicals, 85 Compounds, binary, 39 Common salt, 426 Condy’s disinfecting fluid, 648 Constant proportions, law of, 45 Conversion of volumes into weights, 137 “ Converted nitre,” 416 Copper, 538 alloys of, 541 amalgam, 629 compounds of with oxygen and hy- droxyl, 544 general properties and reactions of the compounds of, 550 glance, 549 pyrites, 538, 662 smelting, 539 Coprolites, 479 Coquimbite, 660 Corrosive sublimate, 531 Corundum, 567 Cotunnite, 607 Cream of tartar, 385 Crith, definition of, 137 Critical point, 121 Crookesite, 557 Cryohydrates, 118 Cryolite, 42§, 566 Crystallization, suspended, 128 fractional, 110 Dalton, atomic theory, 48 Dark red silver ore, 459 Decipium, 585 Decomposition, 103, 113 Dialysis, 129 Diamond, 199 Diantimonic tetroxide, 386 Diarsenious disulphide, 374 Diaspore, 568 Dibismuthous dioxide, 392 disulphide, 395 tetrachloride, 392 Dichro-cobaltic chloride, 668 Didymic oxide, 581 Didyrnium, 581 Didymous chloride, 581 hydrate, 581 nitrate, 581 oxide, 581 sulphate, 581 Diferric trisulphide, 661 Diffusion, 128 of gases, 109, 130 of liquids, 129 Di-iridic hexabromide, 596 hexaehloride, 596 hexahydrate, 597 trioxide, 597 trisulphide, 598 trisulphite, 598 Dimanganic dioxydihydrate, 643 hexaehloride, 642 trioxide, 643 Dimanganous manganite, 643 Dimercurammonic chloride, 537 oxide, 537 Dimolybdic trioxy-hexachloride, 621 Dimolybdous hexabromide, 620 hexaehloride, 620 water of, 88 Crystallography, 131 Crystalloids, 130 “ Crystals of the leaden chamber,” 268 Crystals, systems of, 132 Cupellation process for extraction of sil- ver, 448 Cuprammonic chloride, 544 sulphate, 548 Cupric arsenate, 548 arsenite, 548 bromide, 544 carbonate, 547 chloride, 544 fluoride, 544 hydrate, 544 nitrate, 544 682 INDEX Dimolybdous hexahydrate, 620 trioxide, 620 Dimorphism, 67 Diosmic hexachloride, 601 trioxide, 601 Diopside, 319 Dioptase, 549 Diphosphoric telrasulphide, 362 Diphosphorous tetriodide, 347 Diplumbic trioxide, 609 Dipotassic disulphide, 421 Dirhodic hexahydrate, 599 trioxide, 599 Diruthenic hexachloride, 603 hexahydrate, 604 hexiodide, 603 trioxide, 604 Diseases, zymodc, propagation of, 485 Disilicic hexabromide, 314 hexachloride, 313 hexafluoride, 316 hexiodide, 315 hydrotrioxide, 314 Diosdic dioxide, 427 Dissociation, 103 Distannic trioxide, 327 Distillation, fractional, 109 Disulphur dibromide, 256 dichloride, 255 diniodide, 256 Dithallic tetrachloride, 558 Dithionates, 279 Dititanic hexachloride, 330 trioxide, 332 Diuranic decachloride, 615 Diuranous hexachloride, 615 Dolomite, 510 Double decomposition, 114 Dry copper, 540 Ductility, 409 Dulong and Petit, law of, 68 Dulong and Petit’s law, exceptions to, 69 limit of validity of, 69 Dutch metal, 541 Equivalence of lieat and chemical change, law of, 112 Equivalent proportions, law of, 46 Equivalents, electrochemical, 107 Erbia, 584 Erbium, 584 Erbous hydrate, 585 nitrate, 585 oxide, 584 sulphate, 585 Ervthrine, 666 Estramadurite, 335 Ethylic orthosilicate, 312 silico-orthoformate, 312 Euxenite, 334 Expansion by heat, 398 Fahl ore, 389 Fayalite, 660 Farberite, 627 Feather ore, 389 Felspar, 320 Ferrates, 661 Ferric bromide, 657 chloride, 656 disulphide, 662 fluoride, 657 hydrate, 658 hydrate, colloidal, 658 iodide, 657 nitrate, 660 oxide, 658 phosphate, 661 silicate, 661 sulphate, 660 Ferrous- bromide, 656 carbonate, 659 chloride, 655 chromite, 635 diferric tetroxide, 657 fluoride, 656 hydrate, 657 iodide, 656 nitrate, 659 oxide, 657 Dyad elements, 160, 460, 524 phosphate, 660 silicate, 660 sulphate, 659 sulphide, 661 tungstate, 627 Fibrolite, 572 Fine metal, copper, 539 Fire-damp, 203 Flint, 319 Fluocerite, 578 Fluorides, 308 Fluorine, 306 Fluor-spar, 473 Force, 33 Forces, attractive, 35 Formulae, calculation of, 84 chemical, 75 constitutional, 77 Earthenware, 576 Ebullition, 119 percussive, 121 Electrolysis, 103 laws of, 104 Electro-silvering, 452 Electrum, 551 Elements and compounds, 37 classification of, 88 list of, 38 molecular weights of, 56 Emerald, 320, 521 Emery, 567 Enstatite, 513 Epsomite, 511 Epsom salt, 511 Equations, chemical, 76 Equivalence, 78 empirical, 77 683 INDEX Formulae, graphic, 82 molecular, 77 rational, 77 so-called equivalent, 108 Fowler’s solution, 372 Franco! ite, 357 Franklinite, 514 Fraunhofer lines, 405 Freezing-mixtures, 118 Frit (porcelain), 574 Fulminating gold, 555 silver, 459 Fusco-cobaltic chloride, 668 Fusible metal, Wood’s, 399 Fusing-point, influence of pressure upon, 117 Gold, general properties and reactions of the compounds of, 556 mining, hydraulic, 652 standard, 553 Graphite, 198 Gray antimony ore, 388 Greenockite, 526 Green salt of Magnus,” 591 Green vitriol, 659 “ Gros’ chloride,” 591 Grossularia, 320 Guanite, 612 Guignet’s green, 632 Gun metal, 542 Gunpowder, 416 Gnrolite, 580 Gypsum, 477 burnt, 477 Fusion, 117 change of volume accompanying, 117 latent heat of, 117 Gadolinite, 583 Gahnite, 514 Galena, 613 Gallic choride, 577 Haematite, brown, 658 red, 658 Haemocyanin, 539 Haemoglobin, iron in, 651 Haidingerite, 373 Haloid salts, definition of, 43 Hardness, 408 Hausmannite, 643 Heat, atomic, 68 molecular, 70 specific, 67 specific, table of, 73 unit of, 68 Heavy glass, Faraday’s, 613 Heavy-spar, 465 Hemihedral forms, 133 Ilepar sulphuris, 422 Heptaferric octosulphide, 662 Hexad elements, 243, 614, 629 Hexagonal system, 135 Homogeneity, chemical, 108 Horn-quicksilver, 630 Horn-silver, 452 Hubnerite, 627 Hydracids, definition of, 42 Hydrargillite, 568 Hydrate, definition of, 43 Hydric ammonic sodic phosphate, 445 oxide, 169 peroxide, 175 persulphide, 254 potassic sodic phosphate, 433 potassic tartrate, 385 Hydrogen, 140 displaceable, 41 liquefaction of, 148 occlusion of, by metals, 148 Hydrogen ium, 148 Hydroraagnesite, 510 Hydrosulphyl, 86, 254 Hydroxydimercurammonic iodide, 537 Hydroxyl, 86, 175 Hydroxylamine, 235 Hypiodous chloride, 300 oxide, 577 sulphate, 577 Gallium, 576 general properties and reactions of the compounds of, 577 Garnierite, 670 Gas carbon, 197 Gases, diffusion of, 130 expansion by heat, 52 liquefaction of, 123 relation of, to pressure, 52 solubility of, 124 Gay-Lussac, law of, 54 “ Gerhardt’s base,” chloride of, 591 German silver, 671 Gilding, 552 Glance-cobalt, 663 Glass, 480 annealing, 482 Bohemian, 480 bottle, 480 colored, 483 composition of, 483 crown, 480 devitrification of, 483 flint, 480 making, 480 plate, 480 potash, 480 soda,4Bo toughened, 482 unannealed, 481 window, 480 Glauberite, 478 Glauber’s salt, 430 Glucinum, 521 Gold, 551 fineness of, 553 fulminating, 555 684 INDEX, Hypochlorites, 181 Hypomolybdons bromide, 620 chloride, 619 oxide, 620 Hypopalladous oxide, 594 sulphide, 594 Hypophosphites, 350 Hyposulphurous chloride, 255 hydrosnlphate, 254 Hypotnngstons bromide, 624 chloride, 624 iodide, 624 Hypovanadic chloride, 365 oxide, 365 Hypovanadous chloride, 364 oxide, 365 Johannite, 616 Kaolin, 572, 573 of Ellenbogen, 572 Kelp, 296 Keramohalite, 569 Kerargyrite, 452 Kieserite, 511 Kobeflite, 396 Kupfer nickel, 670 Labradorite, 319 Lamp-black, 196 Lana philosophica, 517 Lanthanous chloride, 582 hydrate, 582 oxide, 582 sulphate, 582 Lanthanum, 582 Lapis lazuli, 573 Latent heat of fusion, 117 vapors, 122 Laughing gas, 220 Laurite, 605 Ice, 173 artificial production of, 232 Indie chloride, 562 hydrate, 563 nitrate, 563 oxide, 562 sulphate, 563 sulphide, 563 sulphite, 563 Indigo copper, 549 Indium, 561 ammonia alum, 563 general properties and reactions of the compounds of, 563 Introduction, 33 Induction tube, 166 Lead, 605 basic “ hyponitrate ” of, 610 compounds of, with oxygen, 608 desilverization of, 448 general properties and reactions of the compounds of, 613 Leblanc’s process for the manufacture of sodic carbonate, 428 lodargyrite, 453 lodates, 303 lodides, 300 lodine, 295 as a heptad, 305 lodous chloride, 300 lons, 103 Iridic bromide, 597 chloride, 597 hydrate, 597 iodide, 597 oxide, 597 sulphide, 598 Iridium, 595 black, 596 Lepidolite, 435, 572 Libethenite, 548 Liebigite, 614 Light red silver ore, 459 Lignite, 198 Lime, chloride of, 476 kilns, 474 milk of, 474 superphosphate of, 480 Limestone, 477 Liquids, diffusion of. 129 solubility of, 124 Litharge, 609 Lithia, 436 Lithic carbonate, 437 chloride, 436 dithionate, 437 fluoride, 436 hydrate, 436 iodide, 436 nitrate, 437 oxide, 436 perchlorate, 437 phosphate, 437 sulphate, 437 general properties and reactions of the compounds of, 598 Iridous sulphide, 598 Iron, 650 alum, 660 amalgam, 529 general properties and reactions of the compounds of, 662 meteoric, 650 passive state of, 655 pyrites, 662 telluric, 650 Irresolvable nebulae, spectra of, 406 Isomerism, 110 Ismorphism, 64 Lithium, 435 general properties and reactions of the compounds of, 437 Liver of sulphur, 422 INDEX, 685 Loadstone, 658 Lucifer matches, 340 Luminous paints, 467 Luteo-cobaltic chloride, 669 Manganic perflnoride, 642 peroxide, 644 peroxide, regeneration of, 645 sulphate, 647 Manganite, 643 Manganous bromide, 642 carbonate, 646 chloride, 641 chromite, 635 dimanganic textroxide, 643 dithionate, 646 Magnesia, 509 Magnesia alba, 510 Magnesia usta, 509 Magnesic aluminate, 568 ammonic arsenate, 512 ammonic carbonate, 510 ammonic chloride, 508 ammonic orthophosphate, 512 ammonic sulphate, 511 arsenate, 512 borate, 512 boride, 513 bromide, 508 calcic carbonate, 510 fluoride, 642 hydrate, 643 iodide, 642 nitrate, 646 oxide, 642 silicate, 647 sulphate, 646 sulphide, 649 tungstate, 627 Marble, 477 Marcasite, 662 Marsh’s test, 377 Matches, safety, 340 Matter, 33 Maximum work, law of, 112 Measures of capacity, 137 length, 136 surface, 136 weight, 137 Meerschaum, 319, 513 calcic chloride, 508 carbonate, 510 chloride, 508 chromate, 637 fluoride, 509 hydrate, 509 iodide, 509 nitrate, 510 nitride, 513 orthophosphate, 512 oxide, 508 phosphate, 512 potassic carbonate, 510 potassic chloride, 508 potassic orthophosphate, 512 potassic sulphate, 511 silicate, 513 silicide, 311, 513 sodic fluoride, 509 sodic orthophosphate, 512 sulphate, 510 sulphydrate, 513 sulphide, 513 Magnesite, 509 Mendeleef, arrangement of elements, 91, 92 Mercuramraonic chloride, 537 Mercuridiammonic dichloride, 537 Mercuric bromide, 532 carbonate, 534 chloride, 531 chromate, 638 fluoride, 532 iodide, 532 nitrate, 534 nitride, 536 oxide, 533 oxychloride, 632 phosphate, 535 potassic sulphide, 536 sulphate, 534 sulphide, 535 sulphochloride, 536 Mercurius solubilis Hahnemanni, 536 Mercurosammonic chloride, 536 nitrate. 536 Mercurosodiammonic dichloride, 537 Mercurous bromate, 634 bromide, 530 carbonate, 534 Magnesium, 507 general properties and reactions of the compounds of, 513 light, 508 Magnetic iron ore, 657 oxide, 657 properties of elements, 94 pyrites, 662 Malachite, 547 Malleability, 409 Malthacite, 572 Manganates, 647 Manganese, 640 alum, 647 black oxide of, 644 blende, 649 chlorate, 533 chloride, 530 fluoride, 531 iodide, 531 oxide, 533 nitrate, 533 characteristic properties and reactions of the compounds of, 650 Manganic dioxide, 644 disulphide, 650 perchloride, 642 perchlorate, 534 sulphate, 534 686 INDEX Mercurous sulphide, 535 Mercury, 527 general properties and reactions of the compounds of, 537 Metallic elements, distinguishing charac- teristics of the, 397 Metal slag, copper, 539 Metals, 397 colors of ignited liquid, 400 expansion of by heat, 398 fusibility of, 398 of the rare earths, 678 Molybdous oxide, 620 • sulphide, 623 Monad elements, 140, 290, 411-447 Monazite, 334 Monoclinic system, 134 Mortar, 475 hydraulic, 475 Mosaic gold, 329 Muntz metal, 541 Multiple proportions, law of, 46 Mysorin, 547 order of ductility of, 409 order of malleability of, 409 relations of, to gravity, 406 relations of, to heat, 398 relations of, to light, 399 relative tenacity of, 408 specific gravity of, 406 volatility of. 399 Metamerism, 110 Metaphosphates, 354 Metastannates, 327 Metatungstates, 628 Meteoric iron, 650 Nascent state, 55 Needle iron ore, 658 Needle ore, 396 Nessler’s solution, 537 Neutralization, change of volume in, 116 heat of, 116 Neumann, law of molecular heat, 70 Nickel, 670 alloys of, 671 bloom, 673 general properties and reactions of the compounds of, 674 glance, 670 plating, 671 silver, 671 of Lenarto, 148 Meyer, Lothar, curve of atomic volumes, 95 Miargyrite, 389 Microcosmic salt, 445 Milk of lime, 474 Millerite, 673 Miloschine, 572 Mimetesite, 613 Mineral chameleon, 648 Minium, 609 Nickeiic disulphide, 674 hydrate, 672 oxide, 672 Nickelous arsenate, 673 bromide, 672 carbonate, 673 chloride, 672 fluoride, 672 hydrate, 672 iodide, 672 Moiree mStallique, 322 Molecular heat, law of, 70 volume, 96 volume of gases, 96 volume of liquids, 98 volume of solids, 97 nitrate, 673 nitrite, 673 oxide, 672 phosphate, 673 silicate, 670 sulphate, 673 sulphide, 673 Niobium, 378 compounds of, 378 Nitrates, 218 Nitre, 416 plantations, 214 Nitric dioxychloride, 229 oxide, 224 peroxide, 226 Nitrification, 214 Nitrogen, 211 oxygen compounds of, 213 Nitrosyiic chloride, 228 Nitrous bromide, 237 chloride, 236 hydrodiniodide, 237 iodide, 237 oxide, 220 oxychloride, 228 Nitroxylic chloride, 229 Nomenclature, chemical, 39 volumes, calculation of, 99 volumes of liquids, table of, 101 weight, calculation of, 53 weight, determination of, 60 weights, 52 weights of elements, 55 work, law of, 111 Molecule, definition of, 59 Molecules, 48 size of, 52 Molybdates, 621 Molybdenite, 623 Molybdenum, 619 genera] properties and reactions of the compounds of, 623 Molybdic dioxvdibromide, 621 dioxydichloride, 621 oxy tetrachloride, 621 pentachloride, 620 persulphide, 623 sulphide, 623 Molybdous chloride, 620 hydrate, 621 iodide, 620 687 INDEX Non-metals, 140 Notation, chemical, 75 graphic, 82 symbolic, 75 Norwegium, 674 Permanganates, 648 Permanganic hexoxy-dichloride, 649 Perruthenates, 604 Peruranates, 618 Petalite, 435, 572 Pewter, 323 Pharmacolite, 373 Octad elements, 600 Okenite, 319, 480 Olivenite, 549 Opal, 317 Ophite, 319, 513 Ore-furnace, copper, 539 slag, copper, 539 Ornithite. 478 Phenacite, 319, 523 Phosgene gas, 210 Phospham, 363 Phosphamimide, 363 Phosphates, 356 Phosphine, 340 Phosphites, 352 Phosphochalcite, 548 Phosphonic bromide, 342 chloride, 342 Orpiment, 374 Orthite, 582 Osmates, 602 iodide, 342 Phosphor-bronze, 542 Phosphoretted hydrogen, gaseous, 340 liquid, 343 solid, 344 Phosphoric bromide, 347 chloride, 345 chloride, action of, upon organic com- pounds, 346 fluoride, 347 oxynitride, 363 oxytriamide, 363 oxytribromide, 360 oxytrichloride, 359 oxytrichloride, action upon organic compounds, 360 sulphide, 361 sulphotrichloride, 362 Phosphorite, 479 Phosphorosphosphates, 358 Phosphorous bromide, 347 chloride, 345 iodide, 347 Osmium, 600 general properties and reactions of the compounds of, 602 Osmic chloride, 601 hydrate, 601 oxide, 601 peroxide, 601 sulphide, 602 Osmiridium, 600 Osmous oxide, 601 sulphite, 602 Osteolite, 335, 478 Over-poling copper, 540 Oxygen, 160 allotropic, 166 diatomic molecule of, 176 Oxy-hydrogen flame, 165 Ozone, 166 Ozonizer, 166 Packfong, 671 Palladia chloride, 594 oxide, 594 sulphide, 362 Phosphorus, 335 amorphous, 338 compounds of with sulphur, 361 octahedral, 337 oxygen compounds of, 348 red, 338 rhombohedral, 339 / Phosphorylic chloride, 359 Phosphotungstates, 627 Photosphere, 405 “Pink-salt,” 326 Pitchblende, 616 Plaster of Paris, 477 sulphide, 595 Palladium, 592 spongy, 592 general properties and reactions of the compounds of, 595 hydride, 593 Fallacious bromide, 593 chloride, 593 hydride, 593 iodide, 593 nitrate, 694 oxide, 594 sulphate, 594 sulphide, 595 Passive iron, 655 Pentad elements, 211, 335, 581 Pentatitanic hexanitride, 333 Perchlorates, 184 Peridote, 319, 513 Periodates, 305 Periodic law, 90 Perissads, 79 Permanent white, 466 Platinamines, 591 Platinates, 590 Platinic bromide, 589 chloride, 589 hydrate, 590 iodide, 589 oxide, 590 sulphide, 590 Platiniridium, 595 Platinodiammonic chloride, 591 Platinonitrites, 590 688 INDEX Platinotetrammonic chloride, 591 Platinous bromide, 589 chloride, 587 hydrate, 590 iodide, 589 oxide, 590 sulphide, 590 sulphite, 590 Platinum, 586 black, 587 general properties and reactions of the compounds of, 591 spongy, 587 Platosodiaramonic chloride, 591 Potassic chlorate, 417 chloride, 414 chlorochromate, 639 chloroplatinate, 414 chromate, 636 chromic sulphate, 634 chroraous sulphate, 633 cobaltic nitrite, 667 cobaltous sulphate, 666 cupric sulphate, 548 dichromate, 636 di-iridic chloride, 596 diosmic chloride, 601 dioxide, 414 Platosotetrammonic chloride, 591 hydrate, 591 Plattnerite, 609 Plumbic amrnonic sulphate, 612 arsenate, 613 dithionate, 419 ferrate, 661 ferric chloride, 656 ferric sulphate, 660 ferrous chloride, 656 fluoride, 414 borate, 613 bromide, 608 carbonate, 610 chloride, 607 chromate 637 dithionate, 612 fluoride, 608 iodide, 608 hydrate, 415 hydride, 413 iodate, 417 iodide, 414 iridic chloride, 597 lithic sulphate, 437 magnesic chromate, 637 manganate, 647 manganic sulphate, 647 manganite, 644 manganous sulphate, 646 manganous, sulphide, 649 metantiraonate (of Fremy), 420 metaphosphate, 420 metastannate, 32/ molybdate. 622 nitrate, 610 nitrate nitrite, 610 nitrite, 610 oxide, 609 oxychloride, 608 oxyhydrate, 609 perchloride, 608 phosphate, 612 silicate, 613 sulphate, 612 sulphide, 613 sulphochloride 613 tungstate, 627 Plumbous oxide, 608 Poling copper, 540 molybdate, 622 nickelous nitrate, 673 nitrate, 416 nitride, 424 nitrite, 416 oxide, 415 orthophosphate, 419 osmate, 602 osraous sulphite, 602 • palladic chloride, 594 palladous chloride, 593 perchlorate, 417 periodate, 417 permanganate, 648 perruthenate, 604 persulphomolybdate, 623 phosphate, 419 phosphite, 420 phosphomolybdate, 622 platinic chloride, 414, 589 platinonitrite, 590 platinous sulphite, 590 platinous sulphodiplatinate, 590 polysulphides, 421 pyrantimonate, 420 pyrophosphate, 419 pyrosulphate, 418 pyrosulphite, 419 Polymerism, 110 Polytungstates, 626 Porcelain, 573 clay, 572 clay of Passau, 572 glazing, 574 hard, 574 tender, 575 Portland cement, 475 Potash, 415 alum, 570 Potassic aluminic bromide, 566 aluminic chloride, 566 amide, 423 antimonate, 420 antimonylic tartrate, 385 arsenate, 420 aurate, 554 borate, 420 bromate, 417 bromide, 414 carbonate, 417 ruthenate, 604 selenate, 419 INDEX, 689 Potassic selenite, 419 silicate, 420 silicofluoride, 414 siiico-tungstate, 627 sodic carbonate, 430 sodic pyrophosphate, 433 stannicofluoride, 326 sodic sulphate, 431 sulphantimonate, 423 sulpharsenate, 423 sulphate, 418 sulphhydrate, 421 sulphide, 420 sulphindate 563 sulphite, 418 sulphocarbonate, 257 sulphoferrite, 662 sulphomolybdate, 623 sulphostannate, 328 sulphothallate, 561 sulpho-tungstate, 628 tellurate, 419 tetrachromate, 636 tetroxide, 415 thiosulphate, 419 titanofluoride, 330 trichromate, 636 tungstate, 626 tungsto-tungstate, 628 uranate, 618 uranylic sulphate, 617 zincic chloride, 517 zirconofluoride, 333 Radiated pyrites, 662 Radicals, acid, chlorides of, 229 compound, 85 Rare earth metals, general properties and reactions of the compounds of, 585 Rare earths, metals of the, 578 Razoumofiskin, 572 Realgar, 373 Reaumur’s porcelain, 483 Red copper ore, 545 haematite, 658 lead, 609 phosphorus, 338 zinc ore, 517 Refinery slag, copper, 540 Regular system, 132 Reinsch’s test, 377 “ Reiset’s first base,” chloride of, 591 . ‘‘ Reiset’s second base,” chloride of, 591 Rhodic chloride, 599 nitrate, 599 oxide, 599 sulphate, 599 sulphite, 599 Rhodium, 598 general properties and reactions of the compounds of, 599 Rhodonite, 647 Rhodous oxide, 599 sulphite, 599 Rhombic system, 134 Rinmann’s green, 667 Rock crystal, 319 ! Roman alum, 570 cement, 475 Roseo-cobaltic chloride, 669 Rouge, 658 Rubidic borate, 439 bromide, 438 carbonate, 439 Potassium, 411 amalgam, 529 general properties and reactions of the compounds of, 424 Potassoxyl, 86 Pottery, 573 ware, 576 Praseo-cobaltic chloride, 669 Precipitation, fractional, 110 Prehnite, 573 “ Preparing salt,” 327 Proust! te, 459 Pseudo-alums, 646 Pucherite, 364 Puddling, 653 Purple of Cassius, 555 Purpureo-cobaltic chloride, 669 Pyrantimonates, 387 Pyrargyrite, 459 Pyrographitie Oxide, 199 Pyrolusite, 644 Pyromorphite, 613 Pyrophosphates, 355 Pyrophosphorylic chloride, 360 Pyrophyllite, 319 Pyrosulphurylic chloride, 283 chlorate, 439 chloride, 438 dithionate, 439 hydrate, 439 iodide, 438 nitrate, 439 perchlorate, 439 platinic chloride, 438 sulphate, 439 Rubidium, 438 general properties and reactions of the compounds of, 440 Ruby, 567 artificial, 567 Rupert’s drops, 481 Ruthenates, 604 Ruthenic chloride, 603 hydrate, 604 oxide, 604 peroxide, 604 sulphate, 604 sulphide, 605 Quadratic system, 134 Quartz, 316 Quicklime, 474 Ruthenium, 602 general properties and the compounds of, 605 INDEX, Euthenous chloride, 603 oxide, 604 Entile, 332 Silvering, 452 Sirius, elements detected in, 406 Slaked lime, 474 Smalt, 666 Soda alum, 571 Soda-ash process, 429 Sodic aluminate, 568 Sal alembroth, 531 Salt-cake process, 428 Saltpetre, 416 Salts, acid, definition of, 44 basic, definition of, 44 definition of, 43 haloid, definition of, 43 aluminic chloride, 566 amide, 435 antimonate, 434 antimonite, 434 argentic thiosulphate, 458 arsenate, 434 normal, definition of, 44 oxy-, definition of, 43 sulpho-, definition of, 44 Samarium, 585 Samarouschloride, 585 oxide, 585 Sand,3l9 Saponite, 573 Sapphire, 567 artificial, 567 borate, 434 bromate, 428 bromide, 426 carbonate, 428 chlorate, 428 chloride, 426 chromate, 636 dichromate, 636 di-iridic chloride, 596 dithionate, 432 ferrate, 661 fluoride, 426 hydrate, 427 hydride, 426, hyposulphite, 432 iodate, 428 iodide, 426 iridic chloride, 597 iridous sulphite, 598 manganate, 648 metaphosphate, 433 molybdate, 622 nitrate, 427 nitrite, 428 oxide, 427 Satin-spar, 477 Saturated vapors. 121 Saturation, fractional, 110 Scandium, 585 Scandous oxide, 585 Scheele’s green, 548 Scheelite, 627 Schlippe’s salt, 435. Schonite, 511 Schweinfurt green, 372 Sciences, classification of, 34 Secondary action in electrolysis, 105 Selenite, 477 Selenium, 283 Seleniuretted hydrogen, 285 Sellai'te, 509 Senarmontite, 384 Serpentine, 319 noble, 319, 513 Silica, 316 Silicates,. 319 Siliceous calamine. 519 perchlorate, 428 periodate, 428 permanganate, 649 peruranate, 618 phosphate, 432 polysulphides, 435 platinic chloride, 590 platinous sulphite, 590 pyrantimonate, 434 pyrarsenate, 434 pyrophosphate, 433 pyrosulphite, 431 selenate, 432 silicate, 434 silicate (Yorke’s), 319 silicofiuoride, 426 stannicoffuoride, 326 sulphantimonate, 435 sulpharsenate, 435 sulphate, 430 sulphhydrate, 435 sulphide, 435 sulphite, 431 sulphostannate, 329 tellurate, 432 Silicic bromide, 314 chloride, 313 fluoride, 315 hydride, 311 hydrotrichloride, 314 iodide, 315 sulphide, 320 trichlorsulphhydrate, 320 Silicium, 309 Siliciuretted hydrogen, 311 Silicofluorides, 316 Silicon, 309 bromoform, 314 chloroform, 314. iodoform, 314 Sillimanite, 572 Silver, 447 general properties and reactions of the compounds of, 459 glance, 459 standard, 451 thiosulphate, 432 tungstate, 626 INDEX 691 Sodic tungsto-tungstate, 628 uranate, 618 zincic chloride, 517 Sodium, 424 amalgam, 529 general properties and reactions of Strontic chlorate, 469 chloride, 468 chromate, 637 dithionate, 470 fluoi-ide, 469 hydrate, 469 iodide, 469 nitrate, 469 orthophosphate, 470 oxide, 469 peroxide, 469 silicofluoride, 469 sulphate, 470 sulphite, 470 thiosulphate, 470 Strontium, 468 the compounds of, 435 Solder, 323 Solidification, suspended, 119 Solids, solubility of, 125 Solubilities, diagram of, 126 Solubility of gases, 124 liquids, 124 solids, 125 Soluble soda glass, 435 Solution, 124 fractional, 110 Sombrerite, 478 Spathose iron ore, 659 general properties and reactions of the compounds of, 470 Struvite, 512 Substitution, 114 Sulphanhydride, molybdic, 623 Specific gravity, relation of to chemical composition, 96 heat, 68 heat equivalents, 74 heats, table of, 7 3 tungstic, 628 Sulphantimonates, 389 Sulphantimonites, 889 Sulpharsenates, 376 Sulpharsenites, 375 Sulphates, 273 Sulphhydrates, 252 Sulphides, 252 Sulphites, 263 Sulpho-acids, definition of, 42 Sulphobismnthites, 396 Sulphocarbonates, 258 Sulphophosphates, 362 Sulphostannates, 329 Sulphur, 243 allotropic modifications of, 246 halogen compounds of, 254 liver of, 422 oxygen compounds of, 259 plastic, 248 rhombic, 247 Sulphuretted hydrogen, 249 Sulphuric iodide, 256 oxychlorhydrate, 282 oxydichloride, 282 Sulphurous chloride, 255 oxydichloride, 282 Sulphurylic chlorhydrate, 282 Sulphury lie chloride, 282 Sun, elements detected in, 408 Superheated vapors, 121 Supersaturation, 128 Sylvine, 414 Syngenite, 478 Spectra of gases, 402 of solids and liquids, 402 solar and stellar, 405 Spectroscope, 400 Spectrum analysis, 400 delicacy of, 403 Specular iron ore, 658 pig iron, 652 Speculum metal, 542 Speiss cobalt, 663 Spiegeleisen, 652 Spinelle, 568 Spodnmene, 573 Stannates, 327 Stannic bromide, 326 chloride, 325 fluoride, 326 iodide, 326 oxide. 326 sulphide, 328 fluorides, 326 Stannous aurous stannate, 555 bromide, 325 chloride, 325 fluoride, 325 hydrate, 326 iodide, 325 stannate, 327 sulphide, 328 Stassfurtite, 512 Steatite, 319, 513 Steel, 653 natural, 651 oxide, 326 Tachvdrite, 508 Talc,‘3l9, 513 Tantalum, 378 tempering, 654 Stibnite, 388 Stolzite, 627 Stoneware, 575 Strontianite, 470 Strontic bromide, 468 compounds of, 378 Tartar emetic, 385 Tellurates, 289 Telluretted hydrogen, 288 Telluric bismuth, 396 carbonate, 470 692 IXDEX Tellurites, 289 Tellurium, 287 Tenacity, 415 Tenorite, 546 Tephroite, 647 Terbium, 585 Triamylstibine, 381 Triclinic system, 135 Tridymite, 317 Triethylstibine, 381 Triethylsulphinic iodide, 243 Triphyline, 435 Triplumbic tetroxide, 609 Trititanic tetranitride, 333 Tungstates, 626 Tetrad elements, 193, 309, 564, 578, 585, 605 Tetradymite, 228, 396 Tetraphosphorous trisulphide, 361 Tetrathallic hexachloride, 558 Thallic bromide, 558 Tungsten, 623 general properties and reactions of the compounds of, 628 chloride, 558 nitrate, 560 oxide, 558 sulphate, 559 sulphide, 560 Tungstic dioxydibromide, 626 dioxydichloride, 626 hexachloride, 625 oxytetrachloride, 626 pentachloride, 624 sulphide, 628 Tungsto-tungstates, 628 Tungstous chloride, 624 oxide, 625 sulphide, 628 Turpeth mineral, 535 Turquoise, 571 Type metal, 607 Thallium, 556 general properties and reactions of the compounds of, 561 Thallous bromide, 558 carbonate, 559 chloride, 557 fluoride, 558 hydrate, 559 iodide, 558 nitrate, 559 oxide, 558 oxyhydrate, 559 phosphate, 560 pyrophosphate, 560 sulphate, 560. sulphide, 560 zincic sulphate, 560 Thenard’s blue, 667 Thermochemistry, 111 Thick letters, use of, 77 Thio-acids, definition of, 42 Thionylic chloride, 281 Thorite, 330 Thorium, 330 compounds of, 330 Ultramarine, 573 Ultramarine, green, 573 Unit of heat, 68 thermal, 68 Uranates, 617 Uranic hydrate, 616 oxide, 616 pentachloride, 615 Uranium, 614 general properties and reactions of the compounds of, 618 mica, 614 vitriol, 616 yellow, 618 Uranous bromide, 615 chloride, 615 diuranate, 616 fluoride, 615 hydrate, 616 oxide, 616 Tin, 321 amalgam, 530 character and reactions of salts of, compounds of, 323 Tincal, 192. 434 Tinning, 323 Titanates, 332 Titanic chloride, 831 phosphate, 617 sulphate, 616 sulphide, 618 uranate, 616 cyanonitride, 333 nitride, 333 oxide, 332 Uranospherite, 618 Uranyl, radical, 616 Uranylic bromide, 616 chloride, 616 sulphide, 332 Titanium, 330 compounds of, 331 general character and reactions of, 333 nitrate, 617 pyrosulphate, 617 sulphate, 617 sulphide, 618 Titanous oxide, 332 Tombac, 541 Topaz, 573 Toughening copper, 540 Triad elements, 185, 551, 582 Valency, 78 Valentinite, 384 Vanadates, 366 Vanadinite, 36,6 Triads, 90 INDEX 693 Vanadium, 364 Yanadous chloride, 365 oxide, 365 Vapor density, determination of, 59 tension, 120 Vapors, latent heat of, 122 Verdigris, 547 Vermilion, 535 Vivianite, 660 Volborthite, 364 Volume-symbols, 56 Wrought iron, 653 Wulfenite, 622 Xenotime, 572 Xonaltite, 480 Yellow ultramarine, 637 Ytterbium, 585 Yttria, 584 Yttrium, 582 Yttrocerite, 584 Yttrous chloride, 584 fluoride, 584 hydrate, 584 nitrate, 584 oxide, 584 sulphate, 584 Wagnerite, 512 Water, 169 analysis, 486 maximum density of, 173 mineral, 484 of crystallization, 88 potable, 484 temporary hardness of, 477 Waters, ammonia present in, 491 average composition of, unpolluted potable, 504 chlorine in, 496 gases dissolved in, 486 hardness of, 497 mineral matters in suspension in, 500 natural, impurities occurring in, 484 nitrogen as nitrates and nitrites in, 492 Zaffre, 666 Zinc, 514 diamine, 520 general properties and reactions of the compounds of, 520 glass, 519 spinelle, 514, 568 Zincic aluminate, 568 ammonic sulphate, 519 antimonide, 520 arsenide, 520 organic carbon in, 488 organic matter in suspension in, 500 organic nitrogen, 489 potable, classification of, 501 potable, dangerous, 496 potable, safe, 495 potable, suspicious, 495 previous sewage or animal contami- nation in, 593 total combined nitrogen in, 492 total solid matters dissolved in, 488 Wavellite, 571 Weights and measures, 136 Weldon’s process for the regeneration of manganic peroxide, 645 Wernerite, 573 White arsenic, 870 lead, 611 lead, Dutch process of manufacturing, 611 blende, 519 bromide, 517 carbonate, 518 chloride, 516 chromate, 637 chromite, 635 fluoride, 517 hydrate, 517 iodide, 517 nitrate, 518 nitride, 520 oxide, 517 oxychloride, 517 pentasulphide, 520 potassic sulphate, 519 potassic sulphide, 519 phosphate, 519 phosphide, 520 silicate, 519 silicofluoride, 517 lead, Miller’s process of manufactur- ing, 612 lead, Thenard’s process of manufac- turing, 612 metal, copper, 539 precipitate, fusible, 537 vitriol, 518 sulphate, 518 sulphide, 519 Zincoxyl, 86 Zinkenite, 389 Zircon, 319, 333 Zirconia, 334 Zirconic bromide, 334 Willemite, 519 Witherite, 464 Wolfram, 627 chloride, 333 fluoride, 333 hydrate, 334 oxide, 334 Zirconium, 333 ochre, 625 Wollastonite, 480 Wood’s fusible metal, 399 Worthite, 572 Zoisite, 573 LEA BROTHERS CO.’S {Late HENRY C. LEA’S SON & CO.) CLASSIFIED CATALOGUE OP MEDICAL AND SURGICAL PUBLICATIONS. In asking the attention of the profession to the works advertised in the following pages, the publishers would state that no pains are spared to secure a continuance of the confi- dence earned for the publications of the house by their careful selection and accuracy and finish of execution. The large number of inquiries received from the profession for a finer class of bindings than is usually placed on medical books has induced us to put certain of our standard publications in half Russia; and, that the growing taste may be encouraged, the prices have been fixed at so small an advance over the cost of sheep as to place it within the means of all to possess a library that shall have attractions as well for the eye as for the mind of the reading practitioner. The printed prices are those at which books can generally be supplied by booksellers throughout the United States, who can readily procure for their customers any works not kept in stock. Where access to bookstores is not convenient, books will be sent by mail postpaid on receipt of the price, and as the limit of mailable weight has been removed, no difficulty will be experienced in obtaining through the post-office any work in this cata- logue. No risks, however, are assumed either on the money or on the hooks, and no pub- lications hut our own are supplied, so that gentlemen will in most cases find it more con- venient to deal with the nearest bookseller. Nos. 706 and 708 Sansom St., Philadelphia, September, 1885. LEA BROTHERS & CO. PROSPECTUS FOR 1885. A WEEKLY MEDICAL JOURNAL. The Medical News Five Dollars. The American Journal of the Medical Sciences . Five Dollars. SUBSCRIPTION RATES. The Medical News \ Nine Dollars per The American Journal of the Medical Sciences j annum, in advance. COMMUTATION RATES. THE MEDICAL NEWS. A National Weekly Medical Periodical, containing- 28 to 32 Quarto Pages of Reading Matter in Each Issue. The Medical News endeavors to render efficient assistance in the daily work of the practising physician, surgeon and specialist. Every department of medical science finds adequate representation in its columns, and its plan and arrangement are well calculated to suit the convenience and secure the comfort of its readers. In the 2 Lea Brothers & Co.’s Periodicals—Medical News, Am. Journal. THE MEDICAL NEWS-WEEKLY. Original Department its columns are replete with articles of the highest practical value; its Hospital Reports reflect the modes of treatment adopted in the most celebrated hospitals of the globe, and its Department of Progress contains judicious excerpts and translations from all the leading medical periodicals of the world. The Editorial Articles are from the pens of a large and able Editorial Staff, and are everywhere conceded to be the most instructive and scholarly productions of their class in the country. Maintaining a large corps of qualified correspondents in all the medical centres of both hemispheres, The News is in early receipt, by cable, telegraph and mail, of intelligence from all quarters. It thus unites the energy of a newspaper with the elaboration of a scientific magazine. Its reputation for enterprise in the past is the best guarantee for the future that nothing will be left undone to render it a faith- ful counsellor and indispensable assistant to every professional man in active practice. (Continued from first puga-) THE AMERICAN JOURNAL of the MEDICAL SCIENCES, Edited by I. MINIS HAYS, A. M., M. D,, Is published Quarterly, on the first days of January, April, July and October, each Number containing over Three Hundred Octavo Pages, fully Illustrated. In his contribution to “A Century of American Medicine,” published in 1876, Dr. John S. Billings, U. S. A., Librarian of the National Medical Library, Washington, thus graphically outlines the character and services of The American Journal—“The ninety-seven volumes of this Journal need no eulogy. They contain many original papers of the highest value; nearly all the real criticisms and reviews which we possess; and such carefully prepared summaries of the progress of medical science, and abstracts and notices of foreign works, that from this file alone, were all other productions of the press for the last fifty years destroyed, it would be possible to reproduce the great majority of the real contributions of the world to medical science during that period.” This opinion of a man pre-eminently qualified to judge is corroborated by the great circle of readers of the Journal, which includes the thinkers of the profession in all parts of the world. During the coming year the features of the Journal which have given unalloyed satisfaction to two generations of medical men, will be maintained in their vigorous maturity. The Original Department will consist of elaborate and richly illustrated articles from the pens of the most eminent members of the profession in all parts of the country and England. The Review Department will maintain its well-earned reputation for discernment and impartiality, and will contain elaborate reviews of new works and topics of the day, and numerous analytical and bibliographical notices by competent writers. Following these comes the Quarterly Summary of Improvements and Dis- coveries in the Medical Sciences, which, being a classified and arranged condensation of important articles appearing in the chief medical journals of the world, furnishes a compact digest of medical progress abroad and at home. The subscription price of The American Journal of the Medical Sciences has never been raised during its long career. It is still sent free of postage for Five Dollars per annum in advance. Taken together, the Journal and News combine the advantages of the elaborate prep- aration that can be devoted to a quarterly with the prompt conveyance of intelligence by the weekly; while, by special management, duplication of matter is rendered im- possible. It will thus be seen that for the very moderate sum of NINE DOLLARS in advance the subscriber will receive free of postage a weekly and a quarterly journal, both reflecting the latest advances of the medical sciences, and containing an equivalent of more than 4000 octavo pages, stored with the choicest material, original and selected, that can be furnished by the best medical minds of both hemispheres. It would be impossible to find elsewhere so large an amount of matter of the same value offered at so low a price. ♦ Lea Brothers & Co.’s Publications—Period., Compends, Manuals. 3 £@s“ The safest mode of remittance is by bank check or postal money order, drawn to the order of the undersigned; where these are not accessible, remittances for subscrip- tions may be made at the risk of the publishers by forwarding in registered letters. Address, Lea Brothers & Co., Nos. 706 and 708 Sansom St., Philadelphia. ** * Communications to both these periodicals are invited from gentlemen in all parts of the country. Original articles contributed exclusively to either periodical are liberally paid for upon publication. When necessary to elucidate the text, illustrations will be fur- nished without cost to the author. All letters pertaining to the Editorial Department of The Medical News and The American Journal of the Medical Sciences should be addressed to the Editorial Offices, 1004 Walnut Street, Philadelphia. All letters pertaining to the Business Department of these journals should be addressed exclusively to Lea Brothers & Co., 706 and 708 Sansom Street, Philadelphia. MARTSIIORFF, TIB FRY, A, M., M. D., LL. D., A Conspectus of the Medical Sciences; Containing Handbooks on Anatomy, Physiology, Chemistry, Materia Medica, Practice of Medicine, Surgery and Obstetrics. Second edition, thoroughly revised and greatly improved. In one large royal 12mo. volume of 1028 pages, with 477 illustrations. Cloth, $4.25; leather, $5.00. Lately Professor of Hygiene in the University of Pennsylvania. The object of this manual is to afford a conven- ient work of reference to students during the brief moments at their command while in attendance upon medical lectures. It is a favorable sign that it has been found necessary, in a short space of time, to issue a new and carefully revised edition. The illustrations are very numerous and unusu- ally clear, and each part seems to have received its due share of attention. We can conceive such a work to be useful, not only to students, but to practitioners as well. It reflects credit upon the industry and energy of its able editor.—Boston Medical and Surgical Journal, Sept. 3, 1874. We can say, with the strictest truth, that it is the best work of the kind with which we are acquaint- ed. It embodies in a condensed form all recent contributions to practical medicine, and is there- fore useful to every busy practitioner throughout our country, besides being admirably adapted to the use of students of medicine. The book is faithfully and ably executed.—Charleston Medical Journal. April, 1875. STUDFFTS’ SFRIFS OF MAFUALS. A Series of Fifteen Manuals, for the use of Students and Practitioners of Medicine and Surgery, written by eminent Teachers or Examiners, and issued in pocket-size 12mo. volumes of 300-540 pages, richly illustrated and at a low price. The following vol- umes are now ready: Bell’s Comparative Physiology and Anatomy, Gould’s Surgical Diagnosis, Robertson’s Physiological Physics, Bruce’s Materia Medica and Therapeutics, Power’s Human Physiology, Clarke and Lockwood’s Dissectors’ Manual, Ealfe’s Clinical Chemistry, Treves’ Surgical Applied Anatomy, Pepper’s Surgical Pathology, and Klein’s Elements of Histology. The following are in press: Bellamy’s Operative Surgery, Pepper’s Forensic Medicine, and Cubnow’s Medical Applied Anatomy. For separate notices see index on last page. SFRIFS OF CLIFICAL MAFUALS, In arranging for this Series it has been the design of the publishers to provide the profession with a collection of authoritative monographs on important clinical subjects in a cheap and portable form. The volumes will contain about 550 pages and will be freely illustrated by chromo-lithographs and woodcuts. The following volumes are just ready: Butlin on the Tongue; Treves on Intestinal Obstruction; and Savage on Insanity and Allied Neuroses; The following are in active preparation: Hutchinson on Syphilis; Bryant on the Breast; Morris on Surgical Diseases of the Kidney; Broadbent on the Pulse; Owen on Surgical Diseases of Children; Lucas on Diseases of the Urethra; Marsh on Diseases of the Joints, Pick on Fractures and Dislocations, and Ball on the Rectum and Anus. For separate notices see index on last page. FFILL, JO ILF, M. D., and SMITH, F, G„ M. JO., Late Surgeon to the Penna. Hospital. Prof, of the Institutes of Med. in the Univ. of Penna. An Analytical Compendium of the Various Branches of Medical Science, for the use and examination of Students. A new edition, revised and improved. In onelarge royal 12mo. volume of 974 pages, with 374 woodcuts. Cloth, $4; leather, $4.75. LUDLOW, J. L., M. D., A Manual of Examinations upon Anatomy, Physiology, Surgery, of Medicine, Obstetrics, Materia Medica, Chemistry, Pharmacy and Therapeutics. To which is added a Medical Formulary. 3d edition, thoroughly revised, and greatly enlarged. In one 12mo. volume of 816 pages, with 370 illustrations. Cloth, $3.25; leather, $3.75. The arrangement of this volume in the form of question and answer renders it espe- cially suitable for the office examination of students, and for those preparing for graduation. Consulting Physician to the Philadelphia Pfospital, etc. 4 Lea Brothers & Co.’s Publications—Dictionaries. BVJVGLTSOJY, ROBLFY, 31. D., MEDICAL LEXICON; A Dictionary of Medical Science; Containing a concise Explanation of the various Subjects and Terms of Anatomy, Physiology, Pathol- ogy, Hygiene, Therapeutics, Pharmacology, Pharmacy, Surgery, Obstetrics, Medical Juris- prudence and Dentistry, Notices of Climate and of Mineral Waters, Formulae for Officinal, Empirical and Dietetic Preparations, with the Accentuation and Etymology of the Terms, and the French and other Synonymes, so as to constitute a French as well as an English Medical Lexicon. Edited by Pickard J. Dunglison, M. D. In one very large and handsome royal octavo volume of 1139 pages. Cloth, $6.50; leather, raised bands, $7.50; very handsome half Russia, raised bands, $B. Late Professor of Institutes of Medicine in the Jefferson Medical College of Philadelphia. The object of the author, from the outset, has not been to make the work a mere lexi- con or dictionary of terms, but to afford under each word a condensed view of its various medical relations, and thus to render the work an epitome of the existing condition of medical science. Starting with this view, the immense demand which has existed for the work has enabled him, in repeated revisions, to augment its completeness and usefulness, until at length it has attained the position of a recognized and standard authority wherever the language is spoken. Special pains have been taken in the preparation of the present edition to maintain this enviable reputation. The additions to the vocabulary are more numerous than in any previous revision, and particular attention has been bestowed on the accentuation, which will be found marked on every word. The typographical arrangement has been greatly improved, rendering reference much more easy, and every care has been taken with the mechanical execution. The volume now contains the matter of at least four ordinary octavos. A book of which every American ought to be proud. When the learned author of the work passed away, probably all of us feared lest the book should not maintain its place in the advancing science whose terms it defines. Fortunately, Dr. Richard J. Dunglison, having assisted his father in the revision of several editions of the work, and having been, therefore, trained in the methods and imbued with the spirit of the book, has been able to edit it as a work of the kind should be edited—to carry it on steadily, without jar or inter- ruption, along the grooves of thought it has trav- elled during its lifetime. To show the magnitude of the task which Dr. Dunglison has assumed and carried through, it is only necessary to state that more than six thousand new subjects have been added in the present edition.—Philadelphia Medical Times, Jan. 3,1874. work has been well known for about forty years, and needs no words of praise on our part to recom- mend it to the members of the medical, and like- wise of the pharmaceutical, profession. The latter especially are in need of a work which gives ready and reliable information on thousands of subjects and terms which they are liable to encounter in pursuing their daily vocations, but with which they cannot be expected to be familiar. The work before us fully supplies this want.—American Jour- nal of Pharmacy, Feb. 1874. Particular care has been devoted to derivation and accentuation of terms. With regard to the latter, indeed, the present edition may be consid- ered a complete “Pronouncing Dictionary of Medical Science.” It is perhaps the most reliable work published for the busy practitioner, as it con- tains information upon every medical subject, in a form for ready access, and with a brevity as ad- mirable as it is practical.—Southern Medical Record, Feb. 1874. A valuable dictionary of the terms employed in medicine and the allied sciences, and of the rela- tions of the subjects treated under each head. It well deserves trie authority and popularity it has obtained.—British Med. Jour., Oct. 31, 1874. About the first book purchased by the medical student is the Medical Dictionary. The lexicon explanatory of technical terms is simply a sine qua non. In a science so extensive and with such col- laterals as medicine, it is as much a necessity also to the practising physician. To meet the wants of students and most physicians the dictionary must be condensed while comprehensive, and practical while perspicacious. It was because Dunglison’s met these indications that it became at once the dictionary of general use wherever medicine was studied in the English language. In no former revision have the alterations and additions been so great. The chief terms have been set in black letter, while the derivatives follow in small caps; an arrangement which greatly facilitates reference. —Cincinnati Lancet and Clinic, Jan. 10,1874. Few works of this class exhibit a grander monu- ment of patient research and of scientific lore.— London Lancet, May 13,1875. Dunglison’s Dictionary is incalculably valuable, and indispensable to every practitioner of medi- cine, pharmacist and dentist.— Western Lancet, March, 1874. It has the rare merit that it certainly has no rival in the English language for accuracy and extent of references.—London Medical Gazette. As a standard work of reference Dunglison’s HOBLYJY, BTC HARD I)., 31. I). A Dictionary of the Terms Used in Medicine and the Collateral Sciences. Revised, with numerous additions, by Isaac Hays, M. D., late editor of The American Journal of the Medical Sciences. In one large royal 12mo. volume of 520 double-columned pages. Cloth, $1.50; leather, $2.00. It is the best book of definitions we have, and ought always to be upon the student’s table—Southern Medical and Surgical Journal. ROBWFLL, G. F., F. R. A. 8., F. C. S., Lecturer on Natural Science at Clifton College, England. A Dictionary Of Science ; Comprising Astronomy, Chemistry, Dynamics, Elec- tricity, Heat, Hydrodynamics, Hydrostatics, Light, Magnetism, Mechanics, Meteorology, Pneumatics, Sound and Statics. Contributed by J. T. Bottomley, M. A., F. C. S., William Crookes, E.R.S., F.C.S., Frederick Guthrie, 8.A., Ph. D., E. A. Proctor, 8.A., F.E.A.S., G. F. Rodwell, Editor, Charles Tomlinson, F.E.S., F.C.S., and Richard Wornell, M.A., B.Sc. Preceded by an Essay on the History of the Physical Sciences. In one handsome octavo volume of 702 pages, with 143 illustrations. Cloth, $5.00. Lea Brothers & Co.’s Publications—Anatomy. 5 GRAY, HENRY, F. R. S., Lecturer on Anatomy at St. George’s Hospital, London. Anatomy, Descriptive and Surgical. The Drawings by H. Y. Cartes,, M. D., and Dr. Westmacott. The dissections jointly by the Authoe and Dr. Caster. With an Introduction on General Anatomy and Development by T. Holmes, M. A., Surgeon to St. George’s Hospital. Edited by T. Pickering Pick, F. R. G. S., Surgeon to and Lecturer on Anatomy at St. George’s Hospital, London, Examiner in Anatomy, Royal College of Surgeons of England. A new American from the tenth enlarged and improved London edition. To which is added the second American from the latest English edition ot Landmarks, Medical and Surgical, by Luther Holden, F. R. C. S., author of “ Human Osteology,” “ A Manual of Dissections,” etc. In one imperial octavo volume of 1023 pages, with 564 large and elaborate engravings on wood. Cloth, $6.00; leather, $7.00; very handsome half Russia, raised bands, $7.50. This work covers a more extended range of subjects than is customary in the ordinary text-books, giving not only the details necessary for the student, but also the application to those details to the practice of medicine and surgery. It thus forms both a guide for the learner and an admirable work of reference for the active practitioner. The engravings form a special feature in the work, many of them being the size of nature, nearly all original, and having the names of the various parts printed on the body of the cut, in place of figures of reference with descriptions at the foot. They thus form a complete and splendid series, which will greatly assist the student in forming a clear idea of Anatomy, and will also serve to refresh the memory of those who may find in the exigencies of practice the necessity of recalling the details of the dissecting-room. Combining, as it does, a complete Atlas of Anatomy with a thorough treatise on systematic, descriptive and applied Anatomy, the work will be found of great service to all physicians who receive students in their offices, relieving both preceptor and pupil of much labor in laying the groundwork of a thorough medical education. Landmarks, Medical and Surgical, by the distinguished Anatomist, Mr. Luther Holden, has been appended to the present edition as it was to the previous one. This work gives in a clear, condensed and systematic way all the information by which the practitioner can determine from the external surface of the body the position of interna] parts. Thus complete, the work, it is believed, will furnish all the assistance that can be rendered by type and illustration in anatomical study. This well-known work comes to us as the latest American from the tenth English edition. As its title indicates, it has passed through many hands and has received many additions and revisions. The work is not susceptible of more improvement. Taking it all in all, its size, manner of make-up, its character and illustrations, its general accur- acy of description, its practical aim, and its per- spicuity of style, it is the Anatomy best adapted to the wants of the student and practitioner.—Medical Record, Sept. 15, 1883. There is probably no work used so universally by physicians and medical students as this one. It is deserving of the confidence that they repose in it. If the present edition is compared with that issued two years ago, one will readily see how much it has been improved in that time. Many pages have been added to the text, especially in those parts that treat of histology, and many hew cuts have been introduced and old ones modified. —Journal of the American Medical Association, Sept. 1, 1883. HOLDEN, LUTHER, F. it. C, S., Also for sale separate— Landmarks, Medical and Surgical. Second American from the latest revised English edition, with additions by W. W. Keen, M. D., Professor of Artistic Anatomy in the Pennsylvania Academy of the Fine Arts, formerly Lecturer on Anatomy in the Phila- delphia School of Anatomy. In one handsome 12mo. volume of 148 pages. Cloth, $l.OO. Surgeon to St. Bartholomew’s and the Foundling Hospitals, London. This little book is all that can be desired within its scope, and its contents will be found simply in- valuable to the young surgeon or physician, since they bring before him such data as he requires at every examination of a patient. It is written in language so clear and concise that one ought almost to learn it by heart. It teaches diagnosis by external examination, ocular and palpable, of the body, with such anatomical and physiological facts as directly bear on the subject. It is eminently the student’s and young practitioner’s book.—Phy- sician and Surgeon, Nov. 1881. WILSON, ERASMUS, F. It. S. A System of Human Anatomy, General and Special. Edited by W. H. Go brecht, M. D., Professor of General and Surgical Anatomy in the Medical College of Ohio. In one large and handsome octavo volume of 616 pages, with 397 illustrations. Cloth, $4.00; leather, $5.00. Si/.I TIT, JT H., M. I)., and HORNER, WM, E.,3LD., Emeritus Prof, of Surgery in the Univ. of Penna., etc. Late Prof, of Anat. in the JJniv. of Penna. An Anatomical Atlas, Illustrative of the Structure of the Human Body. In one large imperial octavo volume of 200 pages, with 634 beautiful figures. Cloth, $4.50. CLELAND, JOHN, 31. />., F. R. S., Professor of Anatomy and Physiology in Queen’s College, Galway. A Directory for the Dissection of the Human Body. In one 12mo. volume of 178 pages. Cloth, $1.25. 6 Lea Brothers & Co.’s Publications—Anatomy. ALLEN, HARRISON, M. D., Professor of Physiology in the University of Pennsylvania. A System of Human Anatomy, Including Its Medical and Surgical Relations. For the use of Practitioners and Students of Medicine. With an Intro- ductory Section on Histology. By E. O. Shakespeare, M. D., Ophthalmologist to the Philadelphia Hospital. Comprising 813 double-columned quarto pages, with 380 illustrations on 109 full page lithographic plates, many of which are in colors, and 241 engravings in the text. In six Sections, each in a portfolio. Section I. Histology. Section 11. Bones and Joints. Section 111. Muscles and Fascial Section IY. Arteries, Veins and Lymphatics. Section V. Nervous System. Section VI. Organs of Sense, of Digestion and Genito-Urinary Organs, Embryology, Development, Teratology, Superficial Anatomy, Post-Mortem Examinations, and General and Clinical Indexes. Price per Section, each in a handsome portfolio, $3.50; also bound in one volume, cloth $23.00; very handsome half Russia, raised bands and open back, $25.00. For sale by subscription only. Apply to the Publishers. It is the design of this book to present the facts of human anatomy in the manner best suited to the requirements of the student and the practitioner of medicine. The author believes that such a book is needed, inasmuch as no treatise, as far as he knows, contains, in addition to the text descriptive of the subject, a systematic presentation of such anatomical facts as can be applied to practice. A book which will be at once accurate in statement and concise in terms; which will be an acceptable expression of the present state of the science of anatomy; which will exclude nothing that can be made applicable to the medical art, and which will thus embrace all of surgical importance, while omitting nothing of value to clinical medicine,—would appear to have an excuse for existence in a country where most surgeons are general practitioners, and where there are few general practitioners who have no interest in surgery. Extract from Introduction. It is to be considered a study of applied anatomy in its widest sense—a systematic presentation of such anatomical facts as can be applied to the practice of medicine as well as of surgery. Our author is concise, accurate and practical' in his statements, and succeeds admirably in infusing an interest into the study of what is generally con- sidered a dry subject. The department of Histol- ogy is treated in a masterly manner, and the ground is travelled over by one thoroughly famil- iar with it. The illustrations are made witn great care, and are simply superb. There is as much of practical application of anatomical points to the every-day wants of the medical clinician as to those of the operating surgeon. In fact, few general practitioners will read the work without a feeling of surprised gratification that so many points, concerning which they may never have thought before are so well presented for their con- sideration. It is a work which is destined to be the best of its kind in any language.—Medical Record, Nov. 25,1882. CLARKE,W.B.,F.R,C.S. & LOCKWOOD,C. 8,, E.R.C.S. Demonstrators of Anatomy at St. Bartholomew's Hospital Medical School, London. The Dissector’s Manual. In one pocket-size 12mo. volume of 396 pages, with 49 illustrations. Limp cloth, red edges, $1.50. Just ready. See Students’ Series of Manuals, page 3. This is a very excellent m anual for the use of the student who desires to learn anatomy. The meth- ods of demonstration seem to us very satisfactory. There are many woodcuts which, for the most part, are good and instructive. The book is neat and convenient. We are glad to recommend it.— Boston Medical and Surgical Journal, Jan. 17, 1884. TREVES, FREDERICK, E. R. C, So, Senior Demonstrator of Anatomy and Assistant Surgeon at the London Hospital. Surgical Applied Anatomy. In one pocket-size 12mo. volume of 540 pages, with 61 illustrations. Limp cloth, red edges, $2.00. Just ready. See Students’ Series oj Manuals, page 3. He has produced a work which will command a larger circle of readers than the class for which it was written. This union of a thorough, practical acquaintance with these fundamental branches, quickened by daily use as a teacher and practi- tioner, has enabled our author to prepare a work which it would be a most difficult task to excel.— The American Practitioner Feb. 1884. CURJSTOW, JOHN, M. D., F. R. C. I\, Professor of Anatomy at King's College, Physician at King's College Hospital. Medical Applied Anatomy. In one pocket-size 12mo. volume. Preparing. See Students’ Series of Manuals, page 3. BELLAMY, ED WARD, F, R. C. S., The Student’s Guide to Surgical Anatomy: Being a Description of the most Important Surgical Regions of the Human Body, and intended as an Introduction to operative Surgery. In one 12mo. volume of 300 pages, with 50 illustrations. Cloth, $2.25. Senior Assistant-Surgeon to the Charing-Cross Hospital, London. HARTSITORNB’S HANDBOOK OF ANATOMY AND PHYSIOLOGY. Second edition, revised, In one royal 12mo. volume of 310 pages, with 220 woodcuts. Cloth, $1.75. [ HORNER'S SPECIAL ANATOMY AND HISTOL- j OGY. Eighth edition, extensively revised and modified. In two octavo volumes of 1007 pages, | with 320 woodcuts. Cloth, $6.00. Lea Brothers & Co.’s Publications—Physics, Physiol., Anat. 7 DRAPER, JOLTY €., M, D., LL, 1)., Professor of Chemistry in the University of the City of New York. Medical Physics. A Text-book for Students and Practitioners of Medicine. In one octavo volume of 734 pages, with 376 woodcuts, mostly original. Cloth, $4. Just ready . The fact that a knowledge of Physics is indispensable to a thorough understanding of Medicine has not been as fully realized in this country as in Europe, where the admirable works of Desplats and Gariel, of Kobertson and of numerous German writers constitute a branch of educational literature to which we can show no parallel. A full appreciation of this the author trusts will be sufficient justification for placing in book form the sub- stance of his lectures on this department of science, delivered during many years at the University of the City of New York. Broadly speaking, this work aims to impart a 'knowledge of the relations existing between Physics and Medicine in their latest state of development, and to embody in the pursuit of this object whatever experience the author has gained during a long period of teaching this special branch of applied science. From the Preface. Certainly we have no text-book as full as the ex- cellent one he has prepared. It begins with a statement of the properties of matter and energy. After these the special departments of physics are explained, acoustics, optics, heat, electricity and magnetism, closing with a section on electro- biology. The applications of all these to physiology and medicine are kept constantly in view. The text is amply illustrated and the many difficult points of the subject are brought forward with re- markable clearness and ability.—Medical and Surg- ical Exporter, July 18,1885. q. The volume from beginning to end teems with useful information. Take the book as a whole and it is one of the most valuable scientific treatises given to the medical profession for a number of years. It is profusely and handsomely illustrated. The work should have a place upon every physician’s library shelf.—Maryland Medical Journal, July 18,1885. q. This is the only work with which we are ac- quainted in which physios is treated with reference to medicine. Preceptors who are axious that their pupils should have a scientific knowledge of med- icine, should make this work a text-book, and re- quire a thorough study of it.—Cincinnati Medical News, July 18, 1885. q. ROBERTSOY, J. McGREGOR, M. A., M, 8., Muirhead Demonstrator of Physiology, University of Glasgow. Physiological Physics. In one 12mo. volume of 537 pages, with 219 illustra- tions. Limp cloth, $2.00. Just ready. See Students’ Series of Manuals, page 3. The title of this work sufficiently explains the nature of its contents. It is designed as a man- nal for the student of medicine, an auxiliary to his text-book in physiology, and it would be pai’ticu- larly useful as a guide to his laboratory experi- ments. It will be found of great value to the practitioner. It is a carefully prepared book of reference, concise and accurate, and as such we heartily recommend it.—Journal of the American Medical Association, Dee. 6, 1884. DA LTOY, JOKY C., M. D., Doctrines of the Circulation of the Blood. A History of Physiological Opinion and Discovery in regard to the Circulation of the Blood. In one handsome 12mo. volume of 293 pages. Cloth, |2. Just ready. Professor Emeritus of Physiology in the College of Physicians and Surgeons, New York. Dr. Dalton’s work is the fruit of the deep research ! of a cultured mind, and to the busy practitioner it | cannot fail to be a source of instruction. It will inspire him with a feeling of gratitute and admir- ation for those plodding workers of olden times, who laid the foundation of the magnificent temple j of medical science as it now stands.—New Orleans Medical and Surgical Journal, Aug. 1885. In the progress of physiological study no fact ! was of greater moment, none more completely i revolutionized the theories of teachers, than the discovery of the circulation of the blood. This explains the extraordinary interest it has to all medical historians. The volume before us is one of three or four which have been written within a few years by American physicians. It is in several respects the most complete. The volume, though small in size, is one of the most creditable con- tributions from an American pen to medical history that has appeared.—Med. & Surg. Rep., Dec. 6,1884. THE SAME AUTHOR. The Topographical Anatomy of the Brain. In three very handsome quarto volumes comprising 178 pages of descriptive text. Illustrated with 48 full page photo- graphic plates of Brain Sections, with a like number of explanatory plates, as well as many woodcuts through the text. BELL, F. JEFFREY, M. A., Professor of Comparative Anatomy at King’s College, London. Comparative Physiology and Anatomy. In one 12mo. volume of 561 pages, with 229 illustrations. Limp cloth, $2.00. Just ready. See Students’ Series of Manuals, page 3. ELLIS, GEORGE 'FLYER, Emeritus Professor of Anatomy in University College, London. Demonstrations of Anatomy. Being a Guide to the Knowledge of the Human Body by Dissection. From the eighth and revised London edition. In one very handsome octavo volume of 716 pages, with 249 illustrations. Cloth, $4.25; leather, $5.25. ROBERTS, JOKY 8., A. 3L, M. J)., Prof, of Applied Anat. and Oper. Surg. in Phila. Polyclinic and Coll, for Graduates in Medicine. The Compend of Anatomy. For use in the dissecting-room and in preparing for examinations. In one 16mo. volume of 196 pages. Limp cloth, 75 cents. 8 Lea Brothers & Co.’s Publications—Physiology, Chemistry. DALTON, JOHN C., M. I)., A Treatise on Human Physiology. Designed for the use of Students and Practitioners of Medicine. Seventh edition, thoroughly revised and rewritten. In one very handsome octavo volume of 722 pages, with 252 beautiful engravings on wood. Cloth, $5.00; leather, $6.00; very handsome half Russia, raised bands, $6.50. Professor' of Physiology in the College of Physicians and Surgeons, New York, etc. The merits of Professor Dalton’s text-book, his smooth and pleasing style, the remarkable clear- ness of his descriptions, which leave not a chapter obscure, his cautious judgment and the, general correctness of his facts, are perfectly known. They have made his texhbook the one most familiar to American students.—Med. Record, March 4, 1882. more compact form, yet its delightful charm is re- tained, and no subject is thrown into obscurity. Altogether this edition is far in advance of any previous one, and will tend to keep the profession posted as to the most recent additions to our physiological knowledge.—Michigan Medical News, April, 1882. Certainly no physiological work has ever issued from the press that presented its subject-matter in a clearer and more attractive light. Almost every page bears evidence of the exhaustive revision that has taken place. The material is placed in a One can scarcely open a college catalogue that does not have mention of Dalton’s Physiology as the recommended text or consultation-book.' For American students we would unreservedly recom- mend Dr. Dalton’s work.- Va. Med. Monthly, July,’B2. DO STEM, MICHAEL, M, D., F. E. S., Text-Book of Physiology. Third American from the fourth English edition, with notes and additions by E. T. Reichert, M. D. In one handsome royal 12mo. volume of 908 pages, with 271 illustrations. Cloth, $3.25 ; leather, $3.75. Just ready. Prelector in Physiology and Fellow of Trinity College, Cambridge, England. Dr. Foster’s work upon physiology is so well- known as a text-book in this country, that it needs but little to be said in regard to it. There is scarcely a medical college in the United States where it is not in the hands of the students. The author, more than any other writer with whom we are acquainted, seems to understand what portions of the science are essential for students to know and what may be passed overby them as not important. From the beginning to the end, physiology is taught in a systematic manner. To this third American edition numerous additions, corrections and alterations have been made, so that in its present form the usefulness of the book will be found to be much increased.— Cincinnati Medical News, July 1885. DO WEE, HE NET, 31. 8., F. E. C. S., Examiner in Physiology, Royal College of Surgeons of England. Human Physiology. In one handsome pocket-size 12mo. volume of 396 pages, with 47 illustrations. Cloth, $1.50. See Students’ Series of Manuals, page 3. The prominent character of this work is that of judicious condensation, in which an able and suc- cessful effort appears to have been made by its accomplished author to teach the greatest number of facts in the fewest possible words. The result is a specimen of concentrated intellectual pabu- lum seldom surpassed, which ought to be care- fully ingested and digested by every practitioner who desires to keep himself well informed upon this most progressive of the medical sciences. The volume is one which we cordially recommend to every one of our readers.— The American Jour- nal of the Medical Sciences, October, 1884. This little work is deserving of the highest praise, and we can hardly conceive how the main facts of this science could have been more clearly or concisely stated. The price of the work is such as to place it within the reach of all, while the ex- cellence of its text will certainly secure for it most favorable commendation —Cincinnati Lancet and Clinic, Feb. 16,1884. CAE BEN TEE, WM. 8., 31. D., F. E. S., F. G. S., F. L. S., Registrar to the University of London, etc. Principles of Human Physiology. Edited by Henry Power, M. 8,, Loud., F. R. C. S., Examiner in Natural Sciences, University of Oxford. .A new American from the eighth revised and enlarged edition, with notes and additions by Francis G. Smith, M, D., late Professor of the Institutes of Medicine in the University of Pennsylvania. In one very large and handsome octavo volume of 1083 pages, with two plates and 373 illus- trations. Cloth, $5.50; leather, $6.50; half Russia, $7. FOWNES, GEOEGE, Eh. D. A Manual of Elementary Chemistry; Theoretical and Practical. Em- bodying Watts’ Inorganic Chemistry. New American edition. In one large royal 12mo. volume of over 1000 pages, with 200 illustrations on wood and a colored plate. Cloth, $2.75; leather, $3.25. In a few days. A notice of the previous edition is appended. The book opens with a treatise on Chemical Physics, including Heat, Light, Magnetism and Electricity. These subjects are treated clearly and briefly, but enough is given to enable the stu- dent to comprehend the facts and laws of Chemis- try proper. It is the fashion of late years to omit these topics from works on chemistry but their omission is not to be commended. As was required by the great advance in the science of Chemistry of late years, the chapter on the General Principles of Chemical Philosophy has been entirely rewrit- ten. The latest views on Equivalents, Quantiva- lence, etc., are clearly and fully set forth. This last edition is a great improvement upon its prede- cessors, which is saying not a little of a book that has reached its twelfth edition.—Ohio Medical Re- corder, Oct., 1878. Wohler’s Outlines of Organic Chemistry. Edited by Fittig. Translated by Ira Remsen, M. D., Ph. D. In one 12mo. volume of 550 pages. Cloth, $3. GALLOWAY’S QUALITATIVE ANALYSIS. New edition. LEHM ANN’S MANUAL OF CHEMICAL PHYS- lOLOGY. In one octavo volume of 327 pages, with 41 illustrations. Cloth, $2.25. CARPENTER’S PRIZE ESSAY ON THE USE A ND Abuse of Alcoholic Liquors in Health and Dis- ease. With explanations of scientific words. Small 12mo. 178 pages. Cloth, 60 cents, Lea Brothers & Co.’s Publications—Chemistry. 9 FRANK! ANF>, JE., D. C. L., F. R. S., & JAFF, F. li., F. I, C., Professor of Chemistry in the Normal School Assist. Prof, of Chemistry in the Normal of Science, London. School of Science, London. Inorganic Chemistry. In one handsome octavo volume of 600 pages, with 51 woodcuts and 2 lithographic plates. Cloth, $3.75; leather, $4.75. In a few days. This work on elementary chemistry is based upon principles of classification, nomen- clature and notation which have been proved by nearly twenty years experience in teaching to impart most readily a sound and accurate knowledge of the science. ATTFIFLJ), JOHN, Fh. Professor of Practical Chemistry to the Pharmaceutical Society of Great Britain, etc. Chemistry, General, Medical and Pharmaceutical; Including the Chem- istry of the U. S. Pharmacopoeia. A Manual of the General Principles of the Science, and their Application to Medicine and Pharmacy. A new American, from the tenth English edition, specially revised by the Author. In one handsome royal 12mo. volume of 728 pages, with 87 illustrations. Cloth, $2.50; leather, $3.00. A text-book which passes through ten editions in sixteen years must have good qualities. This remark is certainly applicable to Attfteld’s Chem- istry, a book which is so well known that it is hardly necessary to do more than note the appear- ance of this new and improved edition. It seems, however, desirable to point out that feature of the book which, in all probability, has made it so popular. There can be little doubt that it is its thoroughly practical character, the expression being used in its best sense. The author under- stands what the student ought to learn, and is able to put himself in the student’s place and to appre- ciate his state of mind.—American Chemical Jour- nal, April, 1884. It is a book on which too much praise cannot be bestowed. As a text-book for medical schools it is unsurpassable in the present state of chemical science, and having been prepared with a special view towards medicine and pharmacy, it is alike indispensable to all persons engaged in those de- partments of science. It includes the whole chemistry of the last Pharmacopoeia.—Pacific Medi- cal and Sugrieal Journal, Jan. 1884. BLOXA3T, CHARLES !., Professor of Chemistry in King's College, London. Chemistry, Inorganic and Organic. New American from the fifth Lon- don edition, thoroughly revised and much improved. In one very handsome octavo volume of 727 pages, with 292 illustrations. Cloth, $3.75; leather, $4.75. Comment from us on this standard work is al- most superfluous. It differs widely in scope and aim from that of Attfield, and in its way is equally beyond criticism. It adopts the most direct meth- ods in stating the principles, hypotheses and-facts of the science. Its language is so terse and lucid, and its arrangement of matter so logical in se- quence that the student never has occasion to complain that chemistry is a hard study. Much attention is paid to experimental illustrations of chemical principles and phenomena, and the mode of conducting these experiments. The book maintains the position it has always held as one of the best manuals of general chemistry In the Eng- lish language.—Detroit Lancet, Feb. 1884. The genera! plan of this work remains the same as in previous editions, the evident object being to give clear and concise descriptions of all known elements and of their most important compounds, with explanations of the chemical laws and principles involved. We gladly repeat now the opinion we expressed about a former edition, that we regard Bloxam’s Chemistry as one ot the best treatises on general and applied chemistry.—American Jour, of Pharmacy, Dee. 1883. SIMON, W., Flu 1)., 3! />., Professor of Chemistry and Toxicology in the College of Physicians and Surgeons, Baltimore, and Professor of Chemistry in the Maryland College of Pharmacy. Manual of Chemistry, A Guide to Lectures and Laboratory work for Beginners in Chemistry. A Text-hook, specially adapted for Students of Pharmacy and Medicine. In one Bvo. vol. of 410 pp., with 16 woodcuts and 7 plates, mostly of actual deposits, with colors illustrating 56 of the most important chemical reactions. Cloth, $3.00; also without plates, cloth, $2.50. Just ready. This book supplies a want long felt by students of medicine and pharmacy, and is a concise but thorough treatise on the subject. The long expe- rience of the author as a teacher in schools of medicine and pharmacy is conspicuous in the perfect adaptation of the work to the special needs of the student of these branches. The colored plates, beautifully executed, illustrating precipi- tates of various reactions, form a novel and valu- able feature of the book, and cannot fail to be ap- preciated by both student and teacher as a help over the hard places of the science.—Maryland Medical Journal, Nov. 22,1884. REMSFN, IRA, 31. J)., Fh. D., Professor of Chemistry in the Johns Hopkins University, Baltimore. Principles of Theoretical Chemistry, with special reference to the Constitu- tion of Chemical Compounds. Second and revised edition. In one handsome royal 12mo. volume of 240 pages. Cloth, $1.75. Just ready. The book is a valuable contribution to the chemi- cal literature of instruction. That in so few years a second edition has been called for indicates that many chemical teachers have been found ready to endorse its plan and to adopt its methods. In this edition a considerable proportion of the book has been rewritten, much new matter has been added and the whole has been brought up to date. We earnestly commend this book to every student of chemistry. The high reputation of the author assures its accuracy in all matters of fact, and its judicious conservatism in matters of theory, com- bined with the fulness with which, in a small compass, the present attitude of chemical science towards the constitution of compounds is con- sidered, gives it a value much beyond that accorded to the average text-books of the day.—American Journal of Science, March, 1884. 10 Lea Brothers & Co.’s Publications—Chemistry. CHARLES, T. CRANSTOVN, M. I)., /. C. S., M. S., Formerly Asst. Prof, and Demonst. of Chemistry and Chemical Physics, Queen's College, Belfast. The Elements of Physiological and Pathological Chemistry. A Handbook for Medical Students and Practitioners. Containing a general account of Nutrition, Foods and Digestion, and the Chemistry of the Tissues, Organs, Secretions and Excretions of the Body in Health and in Disease. Together with the methods for pre- paring or separating their chief constituents, as also for their examination in detail, and an outline syllabus of a practical course of instruction for students. In one handsome octavo volume of 463 pages, with 38 woodcuts and 1 colored plate. Cloth, $3.50. The work is thoroughly trustworthy, and in- formed throughout by a genuine scientific spirit. The author deals with the chemistry of the diges- tive secretions in a systematic manner, which leaves nothing to be desired, and in reality sup- plies a want in English literature. The book ap- pears to us to be at once full and systematic, and to show a just appreciation of the relative import- ance of the various subjects dealt with.—British Medical Journal, November 29, 1884. Dr. Charles’ manual admirably fulfils its inten- tion of giving his readers on the one hand a sum- mary, comprehensive but remarkably compact, of the mass of facts in the sciences which have be- come indispensable to the physician ; and, on the other hand, of a system of practical directions so minute that analyses often considered formidable may be pursued by any intelligent person.— Archives of Medicine, Dec. 1884. HOFFMANN, F., A.M., Fh.D., & ROWER E. 8., Fh.D., Public Analyst to the State of New York. Prof, of Anal. Chem. in the Phil. Coll, of Pharmacy. A Manual of Chemical Analysis, as applied to the Examination of Medicinal Chemicals and their Preparations. Being a Guide for the Determination of their Identity and Quality, and for the Detection of Impurities and Adulterations. For the use of Pharmacists, Physicians, Druggists and Manufacturing Chemists, and Pharmaceutical and Medical Students. Third edition, entirely rewritten and much enlarged. In one very handsome octavo volume of 621 pages, with 179 illustrations. Cloth, $4.25. We congratulate the author on the appearance of the third edition of this work, published for the first time in this country also. It is admirable and the information it undertakes to supply is both extensive and trustworthy. The selection of pro- cesses for determining the purity of the substan- ces of which it treats is excellent and the deserip- tion of them singularly explicit. Moreover, it is exceptionally free from typographical errors. We have no hesitation in recommending it to those who are engaged either in the manufacture or the testing of medicinal chemicals.—London Pharma- ceutical Journal and Transactions, 1883. CLOWES, FRANK, I). Sc., London, An Elementary Treatise on Practical Chemistry and Qualitative Inorganic Analysis. Specially adapted for.use in the Laboratories of Schools and Colleges and by Beginners. Third American from the fourth and revised English edition. In one very handsome royal 12mo. volume of about 400 pages, with about 50 illustrations. Cloth, $2.50. In a few days. Senior Science-Master at the High School, Newcastle-under-Lyme, etc. The style is clear, the language terse and vigor- ous. Beginning with a list of apparatus necessary for chemical work, he gradually"unfolds the sub- ject from its simpler to its more complex divisions. It is the most readable book of the kind we have yet seen, and is without doubt a systematic, intelligible and fully equipped laboratory guide and text book.—Medical Record, July 18, 1885. We may simply repeat the favorable opinion which wo expressed after the examination of the previous edition of this work. It is practical in its aims, and accurate and concise in its statements. —American Journal of Pharmacy, August, 1885. RALFE, CHARLES H., M. D., F. R. C. F., Assistant Physician at the London Hospital. Clinical Chemistry. In one pocket-size 12mo. volume of 314 pages, with 16 illustrations. Limp cloth, red edges, $1.50. See Students’ Series of Manuals, page 3. This is one of the most instructive little works that we have met with in a long time. The author is a physician and physiologist, as well as a chem- ist, consequently the book is unqualifiedly prac- tical, telling the physician just what he ought to know, of the applications of chemistry in medi- cine. Dr. Kalfe is thoroughly acquainted with the latest contributions to his science, and it is quite refreshing to find the subject dealt with so clearly and simply, yet in such evident harmony with the modern scientific methods and spirit.—Medical Record, February 2,1884. CLASSEN, ALEXANDER, Professor in the Royal Polytechnic School, Aix-la-Chapelle. Elementary Quantitative Analysis. Translated, with notes and additions, by Edgar F. Smith, Ph. D., Assistant Professor of Chemistry in the Towne Scientific School, University of Penna. In one 12mo. volume of 324 pages, with 36 illust. Cloth, $2.00. It is probably the best manual of an elementary nature extant insomuch as its methods are the best.*?; It teaches by examples, commencing with single determinations, followed by separations, I and then advancing to the analysis of minerals and | such products as are met with in applied chemis- try. It is an indispensable book for students in | chemistry.—Boston Journal of Chemistry, Oct. 1878. GREENE, WILLIAM JL, 31. D., Demonstrator of Chemistry in the Medical Department of the University of Pennsylvania. A Manual of Medical Chemistry, For the use of Students. Based upon Bow- man’s Medical Chemistry, In one 12mo. volume of 310 pages, with 74 illus. Cloth, $1.75. It is a concise manual of three hundred pages, I giving an excellent summary of the best methods of analyzing the liquids and solids of the body, both forthe estimation of their normal constituents and | the recognition of compounds due to pathological conditions. The detection of poisons is treated with sufficient fulness for the purpose of the stu- dentor practitioner.—Boston JL of Chem., June, ’SO. 11 Lea Brothers & Co.’s Publications—Pharm., Mat. Med., Therap. BBVJSTTON, T. 'LAUDKB, 31. D., D.Sc., F. 8.5., F.8.C.P., Lecturer on Materia Medica and Therapeutics at St. Bartholomew's Hospital, London, etc. A Text-book of Pharmacology, Therapeutics and Materia Medica; Including the Pharmacy, the Physiological Action and the Therapeutical Uses of Drugs. In one handsome octavo volume of about 1000 pages, with 188 illustrations. Cloth, $5.50; leather, $6.50. In press. It is with peculiar pleasure that the early appearance of this long expected work is announced by the publishers. Written by the foremost authority on its subject in Eng- land, it forms a compendious treatise on materia medica,, pharmacology, pharmacy, and the practical use of medicines in the treatment of disease. Space has been devoted to the fundamental sciences of chemistry, physiology and pathology, wherever it seemed necessary to elucidate the proper subject-matter of the book. A general index, an index of diseases and remedies, and an index of bibliography close a volume which will undoubtedly be of the highest value to the student, practitioner and pharmacist. It is a scientific treatise worthy to be ranked with the highest productions in physiology, either in our own or any other language. Everything is practical, the dry, hard facts of physiology being pressed into service and applied to the treatment of the commonest complaints. The information is so systematically arranged that it is available for immediate use. The index is so carefully compiled that a reference to any special point is at once obtainable. Dr. Brunton is never satisfied with vague generalities, but gives clear and pre- cise directions for prescribing the various drugs and preparations. Wo congratulate students on being at last placed in possession of a scientific treatise of enormous practical importance.—The Lancet, June 27, 1885. BABBISM, KDWABD, Late Professoi• of the Theory and Practice of Pharmacy in the Philadelphia College of Pharmacy. A Treatise on Pharmacy: designed as a Text-book for the Student, and as a Guide for the Physician and Pharmaceutist. With many Formulae and Prescriptions. Fifth edition, thoroughly revised, by Thomas S. Wiegand, Ph. G. In one handsome octavo volume of 1093 pages, with 256 illustrations. Cloth, $5; leather, $6. No thoroughgoing pharmacist will fail to possess himself of so useful a guide to practice, and no physician who properly estimates the value of an accurate knowledge of the remedial agents em- ployed by him in daily practice, so far as their miscibility, compatibility and most effective meth- ods of combination are concerned, can afford to leave this work out of the list of their works of reference. The country practitioner, who must always be in a measure, his own pharmacist, will find ” it indispensable.—Louisville Medical News, March 29, 1884. I This well-known work presents itself now based I upon the recently revised new Pharmacopoeia. | Each page bears evidence of the care bestowed upon it, and conveys valuable information from the rich store of the editor’s experience. In fact, all that relates to practical pharmacy—apparatus, processes and dispensing—has been arranged and described with clearness in its various aspects, so as to afford aid and advice alike to the student and to the practical pharmacist. The work is judi- ciously illustrated with good woodcuts—American Journal of_ Pharmacy, January, 1884. There is nothing to equal Parrish’s Pharmacy in this or any other language.—London Pharma- ceutical Journal. HKBMAJSTJSr, Dr. L., Professor of Physiology in the University of Zurich. Experimental Pharmacology. A Handbook of Methods for Determining the Physiological Actions of Drugs. Translated, with the Author’s permission, and with extensive additions, by Robert Meade Smith, M. D., Demonstrator of Physiology in the University of Pennsylvania. In one handsome 12mo. volume of 199 pages, with 32 illustrations. Cloth, $1.50. MAISCH, JOHN 31., JPhar. D., Professor of Materia Medica and Botany in the Philadelphia College of Pharmacy„ A Manual of Organic Materia Medica; Being a Guide to Materia Medica of the Vegetable and Animal Kingdoms. For the use of Students, Druggists, Pharmacists and Physicians. New (second) edition. In one handsome royal 12mo. volume of 550 pages, with 242 illustrations. Cloth, $3.00. Just ready. This work contains the substance,—the practical “kernel of the nut” picked out, so that the stu- dent has no superfluous labor. He can confidently accept what this work places before him, without any fear that the gist of the matter is not in it. Another merit is that the drugs are placed before him in such a manner as to simplify very much the study of them, enabling the mind to grasp them more readily. The illustrations are most excellent, being very true to nature, and are alone worth the price of the book to the student. To the practical physician and pharmacist it is a valuable work for handy reference and for keeping fresh in the memory the knowledge of materia medica and botany already acquired. We can and do heartily recommend it.—Medical and Surgical Re- porter, Feb. 14,1885. BBJJCK, J. MITCHELL, M. />., F. B. C. F, Physician and Lecturer on Materia Medica and Therapeutics at Charing Cross Hospital, London. Materia Medica and Therapeutics. An Introduction to Rational Treat- ment. In one pocket-size 12mo. volume of 555 pages. Limp cloth, $1.50. Just ready. See Students’ Series of Manuals, page 3. GBIFFITH, BOBKBT KGLESFIKLD, M. D. A Universal Formulary, containing the Methods of Preparing and Adminis- tering Officinal and other Medicines. The whole adapted to Physicians and Pharmaceut- ists. Third edition, thoroughly revised, with numerous additions, by John M. Maisch, Phar.D., Professor of Materia Medica and Botany in the Philadelphia College of Pharmacy. In one octavo volume of 775 pages, with 38 illustrations. Cloth, $4.50 ; leather, $5.50. 12 Lea Brothers & Co.’s Publications—Mat. Med., Tlierap. STILLF, AM. I)., LL. I)., & MAISCH, J. M., Thar. I)., Professor Emeritus of the Theory and Prac- tice of Medicine and of Clinical Medicine in the University of Pennsylvania. Prof, of Mat. Med. and Botany in Pkila. College of Pharmacy,Sec'ytotheAmeri- can Pharmaceutical Association. The National Dispensatory: Containing the Natural History, Chemistry, Phar- macy, Actions and Uses of Medicines, including those recognized in the Pharmacopoeias of the United States, Great Britain and Germany, with numerous references to the French Codex. Third edition, thoroughly revised and greatly enlarged. In one magnificent imperial octavo volume of 1767 pages, with 311 fine engravings. Cloth, $7.25; leather, $8.00; half Russia, open back, $9.00. With Denison’s “Beady Deference Index” $l.OO in addition to prlce.in any of above styles of binding. Just ready. In the present revision the authors have labored incessantly with the view of making the third edition of The National Dispensatory an even more complete represen- tative of the pharmaceutical and therapeutic science of 1884 than its first edition was of that of 1879. For this, ample material has been afforded not only by the new United States Pharmacopoeia, but by those of Germany and France, which have recently appeared and have been incorporated in the Dispensatory, together with a large number of new non- officinal remedies. It is thus rendered the representative of the most advanced state of American, English, French and German pharmacology and therapeutics. The vast amount of new and important material thus introduced may be gathered from the fact that the additions to this edition amount in themselves to the matter of an ordinary full-sized octavo volume, rendering the work larger by twenty-five per cent, than the last edition. The Therapeutic Index (a feature peculiar to this work), so suggestive and convenient to the practitioner, contains 1600 more references than the last edition—the General Index 3700 more, making the total number of references 22,390, while the 1 ist of illustrations has been increased by 80. Every effort has been made to prevent undue enlargement of the volume by having in it nothing that could be regarded as superfluous, yet care has been taken that nothing should be omitted which a pharmacist or physician could expect to find in it. The appearance of the work has been delayed by nearly a year in consequence of the determination of the authors that it should attain as near an approach to absolute ac- curacy as is humanly possible. With this view an elaborate and laborious series of examinations and tests have been made to verify or correct the statements of the Pharma- copoeia, and very numerous corrections have been found necessary. It has thus been ren- dered indispensable to all who consult the Pharmacopoeia. The work is therefore presented in the full expectation that it will maintain the position universally accorded to it as the standard authority in all matters pertaining to its subject, as registering the furthest advance of the science of the day, and as embody- ing in a shape for convenient reference the recorded results of human experience in the laboratory, in the dispensing room, and at the bed-side. Comprehensive in scope, vast in design and splendid in execution, The National Dispensatory may be Justly regarded as the most important work of its kind extant.-—Louisville Medical News, Dee. 6,1884. up to date. The work has been very well done, a large number of extra-pharmacopoeial remedies having been added to those mentioned in previous editions.—London Lancet, Nov. 22, 1884. Its completeness as to subjects, the comprehen- siveness of its descriptive language, the thorough- ness of the treatment of the topics, its brevity not sacrificing the desirable features of information for which such a work is needed, make this vol- ume a marvel of excellence.—Pharmaceutical Re- cord, Aug. 15,1884. We have much pleasure in recording the appear- ance of a third edition of this excellent work of reference. It is an admirable abstract of all that relates to chemistry, pharmacy, materia mediea, pharmacology and therapeutics. It may be re- garded as embodying the Pharmacopoeias of the civilized nations of the world, all being brought FARQVHARSOW, ROBERT, M. JO., A Guide to Therapeutics and Materia Mediea. Third American edition, specially revised by the Author. Enlarged and adapted to the U. S. Pharmacopoeia by Frank Woodbury, M. D. In one handsome 12mo. volume of 524 pages. Cloth, $2.25. Lecturer on Materia Mediea at St. Mary's Hospital Medical School. Dr. Farquharson’s Therapeutics is constructed upon a plan which brings before the reader all the essential points with reference to the properties of drugs. It impresses these upon him in such a way as to enable him to take a clear view of the actions of medicines and the disordered conditions in which they must prove useful. The double-col- nmned pages—one side containing the recognized physiological action of the medicine, and the other the disease in which observers (who are nearly al- ways mentioned) have obtained from it good re- sults—make a very good arrangement. The early chapter containing rules for prescribing is excel- lent.—Canada Med. and Stirg. Journal, Dec. 1882. STIELE, A LERE I), M. I)., EL. 1 Therapeutics and Materia Mediea. A Systematic Treatise on the Action and Uses of Medicinal Agents, including their Description and History. Fourth edition, revised and enlarged. In two large and handsome octavo volumes, containing 1936 pages. Cloth, $10.00; leather, $12.00; very handsome half Bussia, raised bands, $13.00. Professor of Theory and Practice of Med. and of Clinical Med. in the Univ. of Penna. We can hardly admit that it has a rival in the multitude of its' citations and the fulness of its research into clinical histories, and we must assign it a place in the physician’s library; not, indeed, a*fullyrepresentingthepresentstateofknowledge | in pharmacodynamics, but as by far the most corn- plete treatise upon the clinical and practical side of the question.—Boston Medical and Surgical Jour- nal, Nov. 5,1874. Lea Brothers & Co.’s Publications—Pathol., Histol. 13 COATS, JOSEPH, M. £>., F. F. F. S., Pathologist to the Glasgow Western Infirmary. A Treatise on Pathology. In one very handsome octavo volume of 829 pages, with 339 beautiful illustrations. Cloth, $5.50; leather, $6.50. The work before us treats the subject of Path- ology more extensively than it is usually treated in similar Mrorks. Medical students as well as physicians, who desire a work for study or refer- ence, that treats the subjects in the various de- partments in a very thorough manner, butwithout prolixity, will certainly give this one the prefer- ence to any with which we are acquainted. It sets forth the most recent discoveries, exhibits, in an interesting manner, the changes from a normal condition effected in structures by disease, and points out the characteristics of various morbid agencies, so that they can be easily recognized. But, not limited to morbid anatomy,it explains fully how the functions of organs are disturbed by abnormal conditions. There is nothing belonging to its de- partment of medicine that is notas fully elucidated as our present knowledge Mull admit.— Cincinnati Medical News, Oct. 1883. GREEN, T. HENRY, 31. />., Pathology and Morbid Anatomy, Fifth American from the sixth revised and enlarged English edition. In one very handsome octavo volume of 482 pages, with 150 line engravings. Cloth, §2.50. Lecturer on Pathology and Morbid Anatomy at Charing-Cross Hospital Medical School, London. The fact that this well-known treatise has so rapidly reached its sixth edition is a strong evi- dence of its popularity. The author is to be con- gratulated upon the tnoroughness with which he has prepared this work. It is thoroughly abreast with all the most recent advances in pathology. No work in the English language is so admirably adapted to the wants of the student and practi- tioner as this, and we would recommend it most earnestly to every one.—Nashville Journal of Medi- cine and Surgery, Nov. 1884. WOOD HEAD, G. SIMS, 31. D., F. R. C. P. E., Demonstrator of Pathology in the University of Edinburgh. • Practical Pathology. A Manual for Students and Practitioners. In one beau- tiful octavo volume of 497 pages, with 136 exquisitely colored illustrations. Cloth, $6.00. It forms a real guide for the student and practi- tioner who is thoroughly in earnest in his en- deavor to see for himself and do for himself. To the laboratory student it will be a helpful com- panion, and all those who may wish to familiarize themselves with modern methods of examining morbid tissues are strongly urged to provide themselves with this manual. The numerous drawings are not fancied pictures, or merely schematic diagrams, but they represent faithfully the actual images seen under the microscope. The author merits all praise for having produced a valuable work.—Medical Record, May 31, 1884. It is manifestly the product of one who has him- self travelled over the whole field and who is skilled not merely in the art of histology, but in the obser- vation and interpretation of morbid changes. The work is sure to command a wide circulation. It should do much to encourage the pursuit of path- ology, since such advantages in histological study have never before been offered.—The Lancet, Jan. 5, 1884. SCHAFER, EDWARD A., F. R. S., Assistant Professor of Physiology in University College, London. The Essentials of Histology. In one octavo volume of 246 pages, with 281 illustrations. Cloth, $2.25. Shortly. CORNIL, V., and RANVIER, L., Prof, in the Faculty of Med. of Paris. Prof, in the College of Prance, A Manual of Pathological Histology. Translated, with notes and additions, by E. O. Shakespeare, M. D., Pathologist and Ophthalmic Surgeon to Philadelphia Hospital, and by J. Henry C. Simes, M. I)., Demonstrator of Pathological Histology in the University of Pennsylvania. In one very handsome octavo volume of 800 pages, with 360 illustrations. Cloth, $5.50 ; leather, $6.50; half Russia, raised bands, $7. KLBIX, E., M. E., F. 11. S., Joint Lecturer on General Anat. and Phys. in the Med. School of St. Bartholomew's Hasp., London. Elements of Histology. In one pocket-size 12mo. volume of 360 pages, with 181 illus. Limp cloth, red edges, §1.50. See Students’ Series of Manuals, page 3. Although an elementary work, it is by no means superficial or incomplete, for the author presents in concise language nearly all the fundamental facts regarding the microscopic structure of tissues. The illustrations are numerous and excellent. We commend Dr, Klein’s Elements most heartily to the student.—Medical Record, Dec. 1,1883. FEPPER, A. J., 31. 8., M. S., F. R. C. S., Surgical Pathology. In one pocket-size 12mo. volume of 511 pages, with 81 illustrations. Limp cloth, red edges, $2.00. See Students’ Series of Manuals, page 3. Surgeon and Lecturer at St. Mary's Hospital, London. It is not pretentious, but it will serve exceed- ingly well as a book of reference. It embodies a freat deal of matter, extending over the whole _ eld of surgical pathology. Its form is practical, its language is clear, and the information set forth is well-arranged, well-indexed and well- illustrated. The student wall find in it nothing that is unnecessary. The list of subjects covers the whole range of surgery. The book supplies a very manifest want and should meet with suc- cess.—New York Medical Journal, May 31,1884. SCHAFER’S PRACTICAL HISTOLOGY. In one 1 handsome royal 12mo. volume of 308 pages, with I 40 illustrations. GLUGE’B ATLAS OF PATHOLOGICAL HISTOL- | OGY. Translated by Joseph Leidy, M. D. In one volume, very large imperial quarto, with S2B copper-plate figures, plain and colored and des- criptive letter-press. Cloth, $4.00 14 Lea Brothers & Co.’s Publications—Practice of Med. FLINT, AUSTIN, M. F., Prof, of the Principles and Practice of Med. and of Clin. Med. in Bellevue Hospital Medical College, N. Y. A Treatise on the Principles and Practice of Medicine. Designed for the use of Students and Practitioners of Medicine. With an Appendix on the Researches of Koch, and their bearing on the Etiology, Pathology, Diagnosis and Treatment of Phthisis. Fifth edition, revised and largely rewritten In one large and closely-printed octavo volume of 1160 pages. Cloth, $5.50; leather, $6.50; half Russia, $7. Koch’s discovery of the bacillus of tubercle gives promise of being the greatest boon ever conferred by science on humanity, surpassing even vaccination in its benefits to mankind. In the appendix to his work, Professor Flint deals with the subject from a practical standpoint, discussing its bearings on the etiology, pathology, diagnosis, prog- nosis and treatment of pulmonary phthisis. Thus enlarged and completed, this standard work will be more than ever a necessity to the physician who duly appreciates the re- sponsibility of his calling. A well-known writer and lecturer on medicine recently expressed an opinion, in the highest de- gree complimentary of the admirable treatise of Dr. Flint, and in eulogizing it, he described it ac- curately as “readable and reliable.” No text-book is more calculated to enchain the interest of the student, and none better classifies the multitudi- nous subjects included in it. It has already so far won its way in England, that no inconsiderable number of men use it alone in the study of pure medicine; and we can say of it that it is in every way adapted to serve, not only as a complete guide, but also as an ample instructor in the science and practice of medicine. The style of Dr. Flint is always polished and engaging. The work abounds in perspicuous explanation, and is a most valuable text-book of medicine.—London Medical News. This work is so widely known and accepted as the best American text-book of the practice of medicine that it would seem hardly worth while to give this, the fifth edition, anything more than a passing notice. But even the most cursory exami- nation shows that it is, practically, much more than a revised edition; it is, in fact, rather a new work throughout. This treatise will undoubtedly continue to hold the first place in the estimation of American physicians and students. No one of our medical writers approaches Professor Flint in clearness of diction, breadth of view, and, what we regard of transcendent importance, rational esti- mate of the value of remedial agents. It is thor- oughly practical, therefore pre-eminently the hook for American readers.—St. Louis Clin. Bee., Mar. ’Bl. lIAHTSHOHJSB, 11BJS11Y, M. 7>., LL. />., Lately Professor of Hygiene in the University of Pennsylvania. Essentials of the Principles and Practice of Medicine. A Handbook for Students and Practitioners. Fifth edition, thoroughly revised and rewritten. In one royal 12mo. volume of 669 pages, with 144 illustrations. Cloth, $2.75; half bound, $3.00. Within the compass of 600 pages it treats of the history of medicine, general pathology, general symptomatology, and physical diagnosis (including laryngoscope, ophthalmoscope, etc.), genera! ther- apeutics, nosology, and special pathology and prac- tice. There is a wonderful amount of information contained in this work, and it is one of the best of its kind that we have seen.—Glasgow Medical Journal, Nov. 1882. An indispensable book. No work ever exhibited a better average of actual practical treatment than this one; and probably not one writer in our day had a better opportunity than Dr. Hartshorne for condensing all the views of eminent practitioners into a 12mo. The numerous illustrations will be very useful to students especially. These essen- tials, as the name suggests, are not intended to supersede the text-books of Flint and Bartholow, but they are the most valuable in affording the means to see at a glance the whole literature of any disease, and the most valuable treatment.—Chicago Medical Journal and Examiner, April, 1882. Bilim OWE, JOHN SYFJR, M. 1)., F. 11. C. F., Physician and Joint Lecturer on Medicine at St. Thomas' Hospital. A Treatise on the Practice of Medicine. Second American edition, revised by the Author. Edited, with additions, by James H. Hutchinson, M.D., physician to the Pennsylvania Hospital. In one handsome octavo volume of 1085 pages, with illustrations. Cloth, $5.00; leather, $6.00; very handsome half Russia, raised bands, $6.50. The reader will find every conceivable subject connected with the practice of medicine ably pre- sented, in a style at once clear, interesting and concise. The additions made by Dr. Hutchinson are appropriate and practical, and greatly add to its usefulness to American readers.—Buffalo Med- ical and Surgical Journal, March, 1880. WATSON, Sill THOMAS, M. I)., Late Physician in Ordinary to the Queen. Lectures on the Principles and Practice of Physic. A new American from the fifth English edition. Edited, with additions, and 190 illustrations, by Henry Hartshorne, A. M., M. D., late Professor of Hygiene in the University of Pennsylvania. In two large octavo volumes of 1840 pages. Cloth, $9.00 ; leather, $ll.OO. LECTURES ON THE STUDY OF FEVER. By A. Hudson, M. D., M. R. I. A. In one octavo volume of 308 pages. Cloth, $2.50. STOKES’ LECTURES ON FEVER. Edited by John William Moore, M. D., P. K. Q. C. P. In one octavo volume of 280 pages. Cloth, $2.00. A TREATISE ON FEVER, By Robert D. Lyons, K. C. C. In one Bvo. yol. of 351 pp. Cloth, $2.25. LA ROCHE ON YELLOW FEVER, considered in its Historical, Pathological, Etiological and Therapeutical Relations. In two large and hand- some octavo volumes of 1468 pp. Cloth, $7.00. A CENTURY OF AMERICAN MEDICINE, 1776—1876. By Drs. E. H. Clarke, H. J. Bigelow, S. D. Gross, T. G. Thomas, and J. S. Billings. In one 12mo. volume of 370 pages. Cloth, $2.25. Lea Brothers & Co.’s Publications—Systems of Med. 15 For Sale by Subscription Only. A System of Practical Medicine. Edited by WILLIAM PEPPER, M. D., LL. D., AMERICAN AUTHORS. PROVOST AND PROFESSOR OF THE THEORY AND PRACTICE OP MEDICINE AND OP CLINICAL MEDICINE IN THE UNIVERSITY OP PENNSYLVANIA, Assisted by Louis Stare, M. D., Clinical Professor of tlie Diseases of Children in the Hospital of the University of Pennsylvania. In five imperial octavo volumes, containing about 1100 pages each, with illustrations. Price per volume, cloth, $5; leather, $6 ; half Russia, raised hands and open back, $7. Volume I. (General Pathology, Sanitary Science and General Diseases) contains 1094 pages, with 24 illustrations and is just ready. Volume 11. (General Diseases [con- tinued] and Diseases of the Digestive System) contains 1312 pages, with 27 illustrations, and is just ready. Volume III: (Diseases of the Eespiratory, Circulatory and Haematopoietic Systems) containing about 1050 pages, will he ready October Ist, and the subsequent volumes at intervals of four months thereafter. The publishers feel pardonable pride in announcing this magnificent work. For three years it has been in active preparation, and it is now in a sufficient state of forward- ness to justify them in calling the attention of the profession to it as the work in wdiich for the first time American medicine is thoroughly represented by its worthiest teachers, and presented in the full development of the practical utility which is its preeminent characteristic. The most able men—from the East and the West, from the North and the South, from all the prominent centres of education, and from all the hospitals which afford special opportunities of study and practice—have united in generous rivalry to bring together this vast aggregate of specialized experience. The distinguished. editor has so apportioned the work that each author has had assigned to him the subject which he is peculiarly fitted to discuss, and in which his views will be accepted as the latest expression of scientific and practical knowledge. The practitioner will therefore find these volumes a complete, authoritative and unfailing work of reference, to w7hich he may at all times turn with full certainty of finding what he needs in its most recent aspect, whether he seeks information on the general principles of medi- cine, or minute guidance in the treatment of special disease. So wide is the scope of the work that, with the exception of midwifery and matters strictly surgical, it embraces the whole domain of medicine, including the departments for which the physician is accustomed to rely on special treatises, such as diseases of women and children, of the genito-urinary organs, of the skin, of the nerves, hygiene and sanitary science, and medical ophthalmology and otology. Moreover, authors have inserted the formulas which they have found most efficient in the treatment of the various affections. It may thus be truly regarded as a Complete Library op Practical Medicine, and the general practitioner possessing it may feel secure that he will require little else in the daily round of professional duties. In spite of every effort to condense the vast amount of practical information fur- nished, it has been impossible to present it in less than 5 large octavo volumes, containing about 5500 beautifully printed pages, and embodying the matter of about 15 ordinary octavos. Illustrations are introduced wherever they serve to elucidate the text. As material for the work is substantially complete in the hands of the editor, the pro- fession may confidently await the appearance of the remaining volumes upon the dates above specified. A detailed prospectus of the work will be sent to any address on appli- cation to the publishers. It is a large undertaking, but quite justifiable in the case of a progressive nation like the United States. At any rate, if we may judge of future volumes from the first, it wdll be justified by the result. We have nothing but praise to bestow upon the work. The articles are the work of writers, m»ny of whom are already recognized in this country as authorities on the particular topics on which they deal, whilst the others show by the way they have handled their subjects that they are fully equal to the task they had undertaken, * * * A work which we cannot doubt will make a lasting reputation for itself.—London •Medical Times and Gazette, Blay 9, 1885. REYNOLDS, J. RUSSELL, M. D., Professar of the Principles and Practice of Medicine in University College, London. A System of Medicine. With notes and additions by Henry Hartshoene, A. M., M. D., late Professor cf Hygiene in the University of Pennsylvania. In three large and handsome octavo volumes, containing 3056 double-columned pages, with 317 illustra- tions. Price per volume, cloth, $5.00; sheep, $6.00; very handsome half Eussia, raised bands, $6.50. Per set, clotb, $l5; leather, $18; half Eussia, $19.50. Sold only by subscription. 16 Lea Brothers & Co.’s Publications—Clinical Med., etc. STILL E, ALI TI ED, M. I)., LL. JO., Professor Emeritus of the Theory and Practice of Med. and of Clinical Med. in the Univ. of Penna. Cholera: Its Origin, History, Causation, Symptoms, Lesions, Prevention and Treat- ment. In one handsome 12mo. volume of 163 pages, with a chart. Cloth, $1.25. Just ready. The threatened importation of cholera into the country renders peculiarly timely this work of an authority upon the subject so eminent as Professor Stille. The history of previous epidemics, their modes of propagation, the vast recent additions to our knowledge of the causation, prevention and treatment of the disease, all have been handled so skilfully as,to present with brevity the information which every practitioner should possess in advance of a visitation. This timely little work is full of the learning and good judgment which marks all that comes from the pen of its distinguished author. What he has to say on treatment is characterized by his.usual caution and his well-known preference for a rational system. Altogether, the monograph is one that will have an excellent influence on the professional mind.—Medical and Surgical Reporter, August 1, 1885. q. FLINT, AUSTIN, 31. I). Clinical Medicine. A Systematic Treatise on the Diagnosis and Treatment of Diseases. Designed for Students and Practitioners of Medicine. In one large and hand- some octavo volume of 799 pages. Cloth, $4.50; leather, $5.50; half Russia, $6.00. It is here that the skill and learning of the great clinician are displayed. He has given us a store- house of medical knowledge, excellentfor the stu- dent, convenient for the practitioner, the result of a long life of the most faithful clinical work,_ col- lected by an energy as vigilant and systematic as untiring, and weighed by a judgment no less clear than his observation is close.—Archives of Medicine, Dec. 1879. To give an adequate and useful conspectus of the extensive held of modern clinical medicine is a task ofnoordinary difficulty; but to accomplish this con- sistently with brevity and clearness, the different subjects and their several parts receiving the attention which, relatively to their importance, medical opinion claims for them, is still more diffi- cult. This task, we feel bound to say, has been executed with more than partial success by Dr. Flint, whose name is already familiar to students of advanced medicine in this country as that of the author of two works of great merit on special subjects, and of numerous papers exhibiting much originality and extensive research.— 7he Dublin Journal, Dee. 1879. the Same Author. Essays on Conservative Medicine and Kindred Topics. In one very hand- some royal 12mo. volume of 210 pages. Cloth, $1.38. BBOABBENT, W. IL, M. I)., E. B. C. P., Physician to and Lecturer on Medicine at St. Mary's Hospital. The Pulse. In one 12mo. volume. See Series of Clinical Manuals, page 3. SCHBELBEB, 88. JOSEPH. A Manual of Treatment by Massage and Methodical Muscle Ex- ercise. Translated by Walter Mendelson, M. D., of New York. In one handsome octavo volume of about 300 pages, with about 125 fine engravings. Preparing, lINLAYSON, JAMES, 31. I)., Editor, Physician and Lecturer on Clinical Medicine in the Glasgow Western Infirmary, etc. Clinical Diagnosis. A Handbook for Students and Practitioners of Medicine. With Chapters by Prof. Gairdner on the Physiognomy of Disease; Prof. Stephens on Diseases of the Female Organs; Dr. Robertson on Insanity; Dr. Gemmell on Physical Diagnosis; Dr. Coats on Laryngoscopy and Post-Mortem Examinations, and by the Editor on Case-taking, Family History and Symptoms of Disorder in the Various Systems. In one handsome 12mo. volume of 546 pages, with 85 illustrations. Cloth, $2.63. FENWICK, SAMUEL, 31.1)., The Student’s Guide to Medical Diagnosis. From the third revised and enlarged English edition. In one very handsome royal 12mo. volume of 328 pages, with 87 illustrations on wood. Cloth, $2.25. Assistant Physician to the London Hospital. TANNEB, THOMAS HAWKES, M. 1). A Manual of Clinical Medicine and Physical Diagnosis. Third American from the second London edition. Revised and enlarged by Tilbury F«x, M. D. In one small 12mo. volume of 862 pages, with illustrations. Cloth, $1.50. EOT HE BGILL, J. 31., M. I)., Edin., M. B. C. P., Loud., The Practitioner’s Handbook of Treatment; Or, The Principles of Thera- peutics. New edition. In one octavo volume. Preparing. Physician to the City of London Hospital for Diseases of the Chest. STURGEB’ INTRODUCTION TO THE STUDY OF CLINICAL MEDICINE. Being a Guide to the Investigation of Disease. In one handsome 12mo. volume of 127 pages. Cloth, $1.25. DAVIS’ CLINICAL LECTURES ON VARIOUS IMPORTANT DISEASES. By N. S. Davis, M. D. Edited by Frank H. Davis, M. D. Second edition. 12mo. 287 pages. Cloth, $1.75. Lea Brothers & Co.’s Publications—Hygiene, Electr., Pract 17 RICHARDSON, B. W., M.A., M.D., LL. I)., F.R.S., F.S.A. Fellow of the Royal College of Physic ians, London. Preventive Medicine. In one octavo volume of 729 pages. Cloth, $4; leather, $5; very handsome half Russia, raised bands, $5.50. Dr. Richardson has succeeded in producing a work which is elevated in conception, eomprehen- sive in scope, scientific in character, systematic in arrangement, and which is written in a clear, con- cise and pleasant manner. He evinces the happy faculty of extracting the pith of what is known on the subject, and of presenting it in a most simple, intelligent and practical form. There is perhaps no similar work written for the general public thatcontains such a complete, reliable and instruc- tive collection of data upon the diseases common to the race, their origins, causes, and the measures for their prevention. The descriptions of diseases are clear, chaste and scholarly; the discussion of the question of disease is comprehensive, masterly and fully abreast with the latest and best knowl- edge on the subject, and the preventive measures advised are accurate, explicit and reliable.—The American Journal of the Medical Sciences, April, 1884. This is a book that will surely find a place on the table of every progressive physician. To the medical profession, whose duty is quite as much to prevent as to cure disease, the book will be a boon. —Boston Medical and Surgical Journal, Mar. 6,1884. The treatise contains a vast amount of solid, valu- able hygienic information.—Medical and Surgical Reporter, Feb. 23,1884. BARTHOLOW, ROBFRTS, A. M., M. />., LL. />., Prof, of Materia Medica and General Therapeutics in the Jefferson Med. Coll, of Phila., etc. Medical Electricity. A Practical Treatise on the Applications of Electricity to Medicine and Surgery. Second edition. In one very handsome octavo volume of 292 pages, with 109 illustrations. Cloth, $2.50. The second edition of this work following so soon upon the first would in itself appear to be a sufficient announcement; nevertheless, the text has been so considerably revised and condensed, and so much enlarged by the addition of new mat- ter, that we cannot fail to recognize a vast improve- ment upon the former work. The author has pre- pared his work for students and practitioners—for those who have never acquainted themselves with the subject, or, having done so, find that after a time their knowledge needs refreshing. We think he has accomplished this object. The book is not too voluminous, but is thoroughly practical, sim- ple, complete and comprehensible. It is, more- over, replete with numerous illustrations of instru- ments, appliances, etc.—Medical Record, November 15, 1882. A most excellent work, addressed by a practi- tioner to his fellow-practitioners, and therefore thoroughly practical. The work now before us has the exceptional merit of clearly pointing out where the benefits to be derived from electricity must come. It contains all and everything that the practitioner needs in order to understand in- telligently the nature and laws of the agent he is making use of, and for its proper application in practice. In a condensed, practical form, it pre- sents to the physician all that he would wish to remember after perusing a whole library on medical electricity, including the results of the latest in- vestigations. It is the book for the practitioner, and the necessity for a second edition proves that it has been appreciated by the profession.—Physi- cian and Surgeon, Dec. 1882. THE YEAR-BOOK OF TREATMENT. A Comprehensive and Critical Review for Practitioners of Medi- cine. In one 12mo. volume of 320 pages, bound in limp cloth, with red edges, $1.25. This work presents to the practitioner not only a complete classified account of all the more important advances made in the treatment of Disease during the year ending Sept. 30, 1884, but also a critical estimate of the same by a competent authority. Each department of practice has been fully and concisely treated, and into the consideration of each subject enter such allusions to recent pathological and clinical work as bear directly upon treatment. As the medical literature of all countries has been placed under contri- bution, the references given throughout the work, together with the separate indexes of subjects and authors, will serve as a guide for those who desire to investigate any thera- peutical topic at greater length. In a few moments the busy practitioner can re- fresh his mind as to the principal advances in treatment for a year past. This kind of work is peculiarly useful at the present time, when current literature is teeming with innumerable so-called advances, of which the practitioner has not time to determine the value. Here he has, collected from many sources, a rtsumt of the theories and facts which are new, either entirely or in part, the decision as to their novelty being made by those who by wide reading and long experience are fully competent to render such a verdict.—Ameri- can Journd{ of the Medical Sciences, April, 1885. It is a complete account of the more important advances made in the treatment of disease. Ex- treme pains have been taken to explain clearly in the fewest possible words the views of each writer, and the details of each subject. One of the principle points about the book is its practical, yet concise language. Each editor has well per- formed his duty, and we can say with truth that it is a volume well worth buying for frequent use. Virginia Medical Monthly, March, 1885. HABERSHON, S. 0., M. D., Senior Physician to and late Jject. on Principles and Practice of Med. at Guy's Hospital, London. On the Diseases of the Abdomen; Comprising those of the Stomach, and other parts of the Alimentary Canal, (Esophagus, Csecum, Intestines and Peritoneum. Second American from third enlarged and revised English edition. In one handsome octavo volume of 554 pages, with illustrations. Cloth, $3.50. PAVY’S TREATISE ON THE FUNCTION OP DI- GESTION ; its Disorders and their Treatment. Prom the second London edition. In one octavo volume of 238 pages. Cloth, $2.00. 'CHAMBERS’ MANUAL OF DIET AND REGIMEN IN HEALTH AND SICKNESS. In one hand- some octavo volume of 302 pp. Cloth, $2.75. BARLOW’S MANUAL OF THE PRACTICE OF MEDICINE. With additions by D. F. Condie, M. D. 1 vol. Bvo., pp. 603. Cloth, $2.50. TODD’S CLINICAL' LECTURES ON CERTAIN ACUTE DISEASES. In one octavo volume of 320 pages. Cloth, $2.50. HOLLAND’S MEDICAL NOTES AND REFLEC- TIONS. 1 vol. Bvo., pp. 493. Cloth, $3.50. 18 Lea Brothers & Co.’s Publications—Throat, Lungs, Heart. COHEN, J. SOLIS, M. JO., Lecturer on Laryngoscopy and .Diseases of the Throat and Chest in the Jefferson Medical College. Diseases of the Throat and Nasal Passages. A Guide to the Diagnosis and Treatment of Affections of the Pharynx, CEsophagus, Trachea, Larynx and Naves. Third edition, thoroughly revised and rewritten, with a large number of new illustrations. In one very handsome octavo volume. Preparing. SEILER, CARL, M. I)., Lecturer on Laryngoscopy in the University of Pennsylvania. A Handbook of Diagnosis and Treatment of Diseases of the Throat, Nose and Naso-Pharynx. Second edition. In one handsome royal 12mo. volume of 294 pages, with 77 illustrations. Cloth, $1.75. It is one of the best of the practical text-books on this subject with which we are acquainted. The present edition has been increased in size, but its eminently practical character has been main- tained. Blany new illustrations have also been introduced, a case-record sheet has been added, and there are a valuable bibliography and a good index of the whole. For any one who wishes to make himself familiar with the practical manage- ment of cases of throat and nose disease, the book will be found of great value.—New York Medical Journal, June 9, 1883. The work before us is a concise handbook upon the essentials of diagnosis and treatment in dis- eases of the throat and nose. The art of laryngos- copy, the anatomy of the throat and nose and the pathology of the mucous membrane are discussed with conciseness and ability. The work is pro- fusely illustrated, excels in many essential feat- ures, and deserves a place in the office of the practitioner who would inform himself as to the nature, diagnosis and treatment of a class of dis- eases almost inseparable from general medical practice. With advanced students the book must be very popular on account of its condensed style. —Louisville Medical Neivs, June 26, 1883. BROWNE, LENNOX, E. R. C. S., Edin,, The Throat and its Diseases. Second American from the second English edi- tion, thoroughly revised. With 100 typical illustrations in colors and 50 wood engravings, designed and executed by the Author. In one very handsome imperial octavo volume of about 350 pages. Preparing. Senior Surgeon to the Central Bond,on Throat and Ear Hospital, etc. FLINT, AUSTIN, M. JO., Professor of the Principles and Practice of Medicine in Bellevue Hospital Medical College, N. Y. A Manual of Auscultation and Percussion; Of the Physical Diagnosis of Diseases of the Lungs and Heart, and of Thoracic Aneurism. Third edition. In one hand- some royal 12mo. volume of 240 pages. Cloth, $1.63. It is safe to say that there is not in the English language, or any other, the equal amount of clear, exact ana comprehensible information touching the physical exploration of the chest, in an equal number of words. Professor Flint’s language is precise and simple, conveying without dubiety the results of his careful study and ample ex- perience in such wise that the young will find it the best source of instruction, and the old the most pleasant means of reviving and complementing their knowledge. American Practitioner, June, 1883. Physical Exploration of the Lungs by Means of Auscultation and Percussion. Three lectures delivered before the Philadelphia County Medical Society, 1882-83. In one handsome small 12mo. volume of 83 pages. Cloth, $l.OO. THE SAME AUTHOR. A Practical Treatise on the Physical Exploration of the Chest and the Diagnosis of Diseases Affecting the Respiratory Organs. Second and revised edition. In one handsome octavo volume of 591 pages. Cloth, $4.50. Phthisis: Its Morbid Anatomy, Etiology, Symptomatic Events and Complications, Eatality and Prognosis, Treatment and Physical Diag- nosis ; In a series of Clinical Studies. In one handsome octavo volume of 442 pages. Cloth, $3.50. A Practical Treatise on the Diagnosis, Pathology and Treatment of Diseases of the Heart, Second revised and enlarged edition. In one octavo volume of 550 pages, with a plate. Cloth, $4. GROSS, S. JO., M.D., LL.JK, JD.C.L. Oxon., LL.JD. Cantab. A Practical Treatise on Foreign Bodies in the Air-passages. In one octavo volume of 452 pages, with 59 illustrations. Cloth, $2.75. FULLER ON DISEASES OF THE LUNGS AND AIR-PASSAGES. Their Pathology, Physical Di- agnosis, Symptoms and Treatment. From the second and revised English edition. In one octavo volume of 475 pages. Cloth, $3.50. SMITH ON CONSUMPTION; its Early and Reme- diable Stages. 1 vol. Bvo., pp. 253. Cloth, $2.25. LA ROCHE ON PNEUMONIA. 1 vol. Bvo. of 490 pages. Cloth, $3.00. SLADE ON DIPHTHERIA; its Nature and Treat- ment, with an account of the History of its Pre- valence in various Countries. Second and re vised edition. In one 12mo. vol., pp. 158. Cloth, $1.25. WILLIAMS ON PULMONARY CONSUMPTION; its Nature, Varieties and Treatment. With an analysis of one thousand eases to exemplify its duration. In one Bvo. vol. of 303 pp. Cloth, $2.50. WALSHE ON THE DISEASES OF THE HEART AND GREAT VESSELS. Third American edi- tion. In 1 vol. Bvo., 416 pp. Cloth, $3.00. JONES’ CLINICAL OBSERVATIONS ON FUNC- TIONAL NERVOUS DISORDERS. Second Am- erican edition. In one handsome octavo volume of 340 pages. Cloth, $3.25. Lea Brothers & Co.’s Publications—Nerv. and Ment. Dis., etc. 19 MITCHELL, S. WELL, M. I)., Physician to Orthopaedic Hospital and the Infirmary for Diseases of the Nervous System, Phila., etc. Lectures on Diseases of the Nervous System; Especially in Women. Second edition. In one 12mo. volume of 288 pages. Cloth, $1.75. Just ready. We feel sure that the new edition of Dr. Mitch- ell’s admirable lectures will be received on this side of the Atlantic with more than ordinary at- tention. His subject, the nervous disorders of women, is one that interests every practitioner, and his views on treatment are gradually receiving general acceptance.—London Medical, Times and Gazette, July 4, 1885. BOSS, JAMES, M.D., F.R. C.E., LL. D., Senior Assistant Physician to the Manchester Royal Infirmary. A Text-Book on Diseases of the Nervous System. In one handsome octavo volume of 600 pages, fully illustrated. Shortly. HAMILTON, ALLAN McLANE, M. D., Attending Physician at the Hospital for Epileptics and Paralytics, Blackwell's Island, N. V. Nervous Diseases; Their Description and Treatment. Second edition, thoroughly revised and rewritten. In one octavo volume of 598 pages, with 72 illustrations. Cloth, $4. When the first edition of this good book appeared we gave it our emphatic endorsement, and the Eresent edition enhances our appreciation of the ook and its author as a safe guide to students of clinical neurology. One of the best and most critical of English neurological journals, Brain, has characterized this book as the best of its kind in any language, which is a handsome endorsement from an exalted source. The improvements in the new edition, and the additions to it, will justify its purchase even by those who possess the old.— Alienist and Neurologist, April, 1882. TUKE, DANIEL HACK, Illustrations of the Influence of the Mind upon the Body in Health and Disease. Designed to elucidate the Action of the Imagination. New edition. Thoroughly revised and rewritten. In one handsome octavo volume of 467 pages, with Joint Author of The Manual of Psychological Medicine, etc. two colored plates. Cloth, $3.00. It is impossible to peruse these interesting chap- ters without being convinced of the author’s per- fect sincerity, impartiality, and thorough mental grasp. Dr. Tuke has exhibited the requisite amount of scientific address on all occasions, and the more intricate the phenomena the more firmly has he adhered to a physiological and rational method of interpretation. Guided by an enlight- ened deduction, the author has reclaimed for science a most interesting domain in psychology, previously abandoned to charlatans and empirics. This book, well conceived and well written, must commend itself to every thoughtful understand- ing.—New York Medical Journal, September 6,1884. CLOUS TON, THOMAS S., M. D., I\ R. C. i\, L. R. C. S., Olinical Lectures on Mental Diseases. With an Appendix, containing an Abstract of the Statutes of the United States and of the Several States and Territories re- lating to the Custody of the Insane. By Charles F. Folsom, M. D., Assistant Professor of Mental Diseases, Medical Department of Harvard University. In one handsome octavo volume of 541 pages, illustrated with eight lithographic plates, four of which are beautifully colored. Cloth, $4. Lecturer on Mental Diseases in the University of Edinburgh. The practitioner as well as the student will ac- cept the plain, practical teaching of the author as a forward step in the literature of insanity. It is refreshing to find a physician of Dr. Clouston’s experience and high reputation giving the bed- side notes upon which his experience has been founded and his mature judgment established. Such clinical observations cannot but be useful to the general practitioner in guiding him to a diag- nosis and indicating the treatment, especially in many obscure and doubtful oases of mental dis- ease. To the American reader Dr. Folsom’s Ap- pendix adds greatly to the value of the work, and will make it a desirable addition to every library. —American Psychological Journal, July, 1884. Jglgf’Dr. Folsom’s Abstract may also be obtained separately in one octavo volume of 108 pages. Cloth, $1.50. SAVAGE, GEORGE If., M. D., Insanity and Allied Neuroses, Practical and Clinical. In one 12mo. vol- ume of 551 pages, with 18 typical illustrations. Cloth, $2.00. Just ready. See Series of Clinical Manuals, page 3. Lecturer on Mental Diseases at Guy's Hospital, London. As a handbook, a guide to practitioners and stu- dents, the book fulfils an admirable purpose. The many forms of insanity are described with char- acteristic clearness, the illustrative cases are care- fully selected, and as regards treatment, sound common sense is everywhere apparent. We re- peat that Dr. Savage has written an excellent manual for the practitioner and student.—Am- erican Journal of Insanity, April, 1885. EL AYE AIR, W. S., M. D., E. R. C. I\, The Systematic Treatment of Nerve Prostration and Hysteria. In one handsome small 12mo. volume of 97 pages. Cloth, $l.OO. Blandford on Insanity and its Treatment: Lectures on the Treatment, Medical and Legal, of Insane Patients. In one very handsome octavo volume. 20 Lea Brothers & Co.’s Publications—Surgery. GROSS, S. D., 3L. D., LL. D., D. C. L. Oxon., LL. D, Cantab,, Emeritus Professor of Surgery in the Jefferson Medical College of Philadelphia. A System of Surgery: Pathological, Diagnostic, Therapeutic and Operative. Sixth edition, thoroughly revised and greatly improved. In two large and beautifully- printed imperial octavo volumes containing 2382 pages, illustrated by 1623 engravings. Strongly bound in leather, raised bands, $l5 ; half Russia, raised bands, $l6. Dr. Gross’ System of Surgery has long been the standard work on that subject for students and practitioners.—London Lancet, May 10, 1884. material has been introduced, and altogether the distinguished author has reason to be satisfied that he has placed the work fully abreast of the state of our knowledge.—Med. Record, Nov. 18,1882. His System of Surgery, which, since its first edi- tion in 1859, has been a standard work in this country as well as in America, in “the whole domain of surgery,” tells how earnest and labori- ous and wise a surgeon he was, how thoroughly he appreciated the work done by men in other countries, and how much he contributed to pro- mote the science and practice of surgery in his own. There has been no man to whom America is so much indebted in this respect as the Nestor of surgery.—British Medical Journal, May 10,1884. The work as a whole needs no commendation. Many years ago it earned for itself the enviable rep- utation of the leading American work on surgery, and it is still capable of maintaining that standard! The reason for this need only be mentioned to be appreciated. The author has always been calm and judicious in his statements, has based his con- clusions on much study and personal experience, has been able to grasp his subject in its entirety, and, above all, has conscientiously adhered to truth and fact, weighing the evidence, pro and con, accordingly. A considerable amount of new AS II HURST, JOHN, Jr,, M. I),, Professor of Clinical Surgery, Univ. of Penna., Surgeon to the Episcopal Hospital, Philadelphia. The Principles and Practice of Surgery. Fourth edition, enlarged and revised. In one large and handsome octavo volume of about 1100 pages, with about 575 illustrations. Shortly. GOULD, A, JPEARCE, 31. S., 31, 8., E. R. C. S„ Assistant Surgeon to Middlesex Hospital. Elements of Surgical Diagnosis. In one pocket-size 12mo. volume of 589 pages. Cloth, $2.00. Just ready. See Students’ Series of Manuals, page 3. The student and practitioner will find the principles of surgical diagnosis very satisfactorily set forth with all unnecessary verbiage elimi- nated. Every medical student attending lectures should have a copy to study during the intervals, and if practitioners would devote a portion of their leisure to the study of it, they would receive immense benefit in the way of refreshing their knowledge and bringing it up to the present state of progress.—Cincinnati Medical News, Jan., 1885. GIBNEY, V. I*., 31, D., Orthopaedic Surgery. For the use of Practitioners and Students. In one hand- some octavo volume, profusely illustrated. Preparing. Surgeon to the Orthopaedic Hospital, New York, etc. ROBERTS, JOHN 8., A. M., M. />., Lecturer on Anatomy and on Operative Surgery at the Philadelphia School of Anatomy. The Principles and Practice of Surgery. For the use of Students and Practitioners of Medicine and Surgery. In one very handsome octavo volume of about 500 pages, with many illustrations. Preparing. BELLA 31Y, EDWARD, F, R, C. S., Surgeon and Lecturer on Surgery at Charing Cross Hospital, London. Operative Surgery. Shortly. See Students’ Series of Manuals, page 3. STI3ISON, LEWIS A., B, A,, 31. D,, Prof, of Pathol. Anat. at the Univ. of the City of New York, Surgeon and Curator to Bellevue Hosp. A Manual of Operative Surgery. New (second) edition. In one very hand- some royal 12mo. volume of about 500 pages, with about 350 illustrations. Cloth, $2.50. Shortly. A notice of the previous edition is appended. This volume is devoted entirely to operative sur- gery, and is intended to familiarize the student with the details of operations and the different modes of performing them. The work is hand- somely illustrated, and the descriptions are clear and well-drawn. It is a clever and useful volume; every student should possess one. This work does away with the necessity of pondering over larger works on surgery for descriptions of opera- tions, as it presents in a nutshell w’hat is wanted by the surgeon without an elaborate search to find it.—Maryland Medical Journal, August, 1878. SARGENT ON BANDAGING and OTHER OPERA- TIONS OF MINOR SURGERY. New edition, wjth a Chapter on military surgery. One 12mo. volume of 383 pages, with 187 cuts. Cloth, $1.75. MILLER’S PRINCIPLES OF SURGERY. Fourth American from the third Edinburgh edition. In one Bvo. vol. of 638 pages, with 340 illustrations. Cloth, $3.75.. PIRRIE’S PRINCIPLES AND PRACTICE OP SURGERY. Edited by John Neill, M. D. In one Bvo. vol. of 784 pp. with 316 iilus. Cloth, $3.75. COOPER’S LECTURES ON THE PRINCIPLES AND PRACTICE OF SURGERY. In one Svo.vol. of 767 pages. Cloth, $2.00. SKEY’S OPERATIVE SURGERY. In one vol. Bvo. of 661 pages, with 81 woodcuts. Cloth, $3.25. MILLER’S PRACTICE OF SURGERY. Fourth and revised American from the last Edinburgh edition. In one large Bvo. vol. of 682 pages, with 364 illustrations. Cloth, $8.75. GIBSON’S INSTITUTES AND PRACTICE OF SURGERY. Eighth edition. In two octavo vole, of 965 pages, with 34 plates. Leather $6.50. Lea Brothers & Co.’s Publications—Surgery. 21 BRICHSBN, JOHN 8., B. R. S., B. R. C. S., Professor of Surgery in University College, London, etc. The Science and Art of Surgery; Being a Treatise on Surgical Injuries, Dis- eases and Operations. From the eighth and enlarged English edition. In tivo large and beautiful octavo volumes of 2316 pages, illustrated with 984 engravings on wood. Cloth, $9; leather, raised bands, $11; half Russia, raised bands, $l2. Just ready. After the profession has placed its approval upon a work to the extent of purchasing seven editions, it does not need to be introduced. Simultaneous with the appearance of this edition a translation is being made into Italian and Spanish. Thus this favorite text-book on surgery holds its own in suite of numerous rivals at the end of thirty years. It is a grand book, worthy of the art in the interest of which it is written.—Detroit Lancet, Jan. 10,1885. After being before the profession for thirty years and maintaining during that period a re- putation as a leading work on surgery, there is not much to be said in the way of comment or criti- cism. That it still holds its own goes without say- ing. The author infuses into it his large experi- ence and ripe judgment. Wedded to no school, committed to no theory, biassed by no hobby, he imparts an honest personality in his observations, and his teachings are the rulings of an impartial judge. Such men are always safe guides, and their works stand the tests of time and experience. Such an author is Eriehsen, and such a work is his Surgery.—Medical Record, Feb. 21,1885. Bit YANT, THOMAS, B. It. C. S., Surgeon and Lecturer on Surgery at Guy's Hospital, London. The Practice of Surgery. Fourth American from the fourth and revised Eng- lish edition. In one large and very handsome imperial octavo volume of 1040 pages, with 727 illustrations. Cloth, $6.50; leather, $7.50; half Kussia, $B.OO. Just ready. The treatise takes in the whole field of surgery, that of the eye, the ear, the female organs, ortho- paedies, venereal diseases, and military surgery, as well as more common and general topics. All of these are treated with clearness and with sufficient fulness to suit all practical purposes. The illustrations are numerous and well printed. We do not doubt that this new edition will con- tinue to maintain the popularity of this standard work.—Medical and Surgical Reporter, Feb. 14, ’B5. This most magnificent work upon surgery has reached a fourth edition in this country, showing the high appreciation in which it is held by the American profession. It comes fresh from the pen of the author. That it is the very best work on surgery for medical students we think there can be no doubt. The author seems to have understood just what a student needs, and has prepared the work accordingly.—Cincinnati Medical News, January, 1885. This most magnificent work upon surgery has reached a fourth edition in this country, showing the high appreciation in which it is held by the American profession. It comes fresh from the pen of the author. That it is the very best work on surgery for medical students we think there can be no doubt. The author seems to have understood just what a student needs, and has prepared the work accordingly.—Cincinnati Medical News, January, 1885. Diseases of the Breast. In one 12mo. volume. Preparing. See Series of Clinical Manuals, page 3. the same Author. BTJTLIN, HENRY T., B. R. C. S., Assistant Surgeon to St. Bartholomew's Hospital, London. Diseases of the Tongue. In one 12mo. volume of 456 pages, with 8 colored plates and 3 woodcuts. Cloth, $3.50. Just ready. See Series of Clinical Manuals, page 3. ESMARCH, Dr. BRIBHRICH, Professor of Surgery at the University of Kiel, etc. Early Aid in Injuries and Accidents. Five Ambulance Lectures. Trans- lated by 11. 1:1. H. Princess Christian. In one handsome small 12mo. volume of 109 pages, with 24 illustrations. Cloth, 75 cents. The course of instruction is divided into five sections or lectures. The first, or introductory lecture, gives a brief account of the structure and organization of the human body, illustrated by clear, suitable diagrams. The second teaches how to give judicious help in ordinary injuries—contu- sions, wounds, haemorrhage and poisoned wounds. The third treats of first aid in eases of fracture and of dislocations, in sprains and in burns. Next, the methods of affording first treatment in cases of frost-bite, of drowning, of suffocation, of loss of consciousness and of poisoning are described; and the fifth lecture teaches how injured persons may be most safely and easily transported to their homes, to a medical man, or to a hospital. The illustrations in the book are clear and good.—Medi- cal Times and Gazette, Nov. 4,1882. TRBVBS, BRED ERICK, B. R. C. S., Assistant Surgeon to and Lecturer on Surgery at the London Hospital. illustrations. Limp cloth, blue edges, $2.00. Just ready. See Series of Clinical Manuals, page 3. Intestinal Obstruction. In one pocket-size 12mo. volume of 522 pages, with 6G A standard work on a subject that has not been so comprehensively treated by any contemporary English writer. Its completeness renders a full review difficult, since every chapter deserves mi- nute attention, and it is impossible to do thorough justice to the author in a few paragraphs. Intes- tinal Obstruction is a work that will prove of equal value to the practitioner, the student, the pathologist, the physician and the operating sur- geon.—British Medical Journal, Jan. 31, 1885. BALL, CHARLES 8., M. Ch., Hub., B. R. C. S. 8., Diseases of the Rectum and Anus. In one 12mo. volume of 550 pages. Preparing. See Series of Clinical Manuals, page 3. Surgeon and Teacher at Sir P. Dun's Hospital, Dublin. HRJJITT, ROBERT, M. R. C. S., etc. The Principles and Practice of Modern Surgery. From the eighth London edition. In one Bvo. volume of 687 pages, with 432 illus. Cloth, |4; leather, $6. 22 Lea Brothers & Co.’s Publications—Surgery. HOLMES, TIMOTHY, 31. A., Surgeon and Lecturer on Surgery at St. George's Hospital, London. • A System of Surgery; Theoretical and Practical. IN TREATISES BY VARIOUS AUTHORS. American edition, thoroughly revised and re-edited by John H. Packard, M. I)., Surgeon to the Episcopal and St. Joseph’s Hospitals, Philadelphia, assisted by a corps of thirty-three of the most eminent American surgeons. In three large and very handsome imperial octavo volumes containing 3137 double- columned pages, with 979 illustrations on wood and 13 lithographic plates, beautifully colored. Price per volume, cloth, $6.00; leather, $7.00 ; half Russia, $7,50. Per set, cloth, $18.00; leather, $21.00; half Russia, $22.50. Sold only by subscription. Volume I. contains General Pathology, Morbid Processes, Injuries in Gen- eral, Complications op Injuries and Injuries of Regions. Volume 11. contains Diseases of Organs of Special Sense, Circulatory Sys- tem, Digestive Tract and Genito-Urinary Organs. Volume HI. contains Diseases of the Respiratory Organs, Bones, Joints and Muscles, Diseases op the Nervous System, Gunshot Wounds, Operative and Minor Surgery, and Miscellaneous Subjects (including an essay on Hospitals). This great work, issued some years since in England, has won such universal confi- dence wherever the language is spoken that its republication here, in a form more thoroughly adapted to the wants of the American practitioner, has seemed to be a duty owing to the profession. To accomplish this, each article has been placed in the hands of a gentleman specially competent to treat its subject, and no labor has been spared to bring each one up to the foremost level of the times, and to adapt it thoroughly to the practice of the country. In certain cases this has rendered necessary the substitution of an entirely new essay for the original, as in the case of the articles on Skin Diseases, on Diseases of the Absorbent System, and on Anaesthetics, in the use of which American practice differs from that of England. The same careful and conscientious revision has been pursued throughout, leading to an increase of nearly one-fourth in matter, while the series of illustrations has been nearly trebled, and the whole is presented as a complete exponent of British and American Surgery, adapted to the daily needs of the working practitioner. In order to bring it within the reach of every member of the profession, the five vol- umes of the original have been compressed into three by employing a double-columned royal octavo page, and in this improved form it is offered at less than one-half the price of the original. It is printed and bound to match in every detail with Reynolds’ System of Medi- cine. The work will be sold by subscription only, and in due time every member of the profession will be called upon and offered an opportunity to subscribe. The authors of the original English edition are men of the front rank in England, and Dr. Packard has been fortunate in securing as his American coadjutors such men as Bart biolow, Hyde, Hunt, Conner, Stimson, Morton, Hodgen, Jewell and their colleagues. As a whole, the work will be solid and substantial, and a valuable addition to the library of any medical man. It is more wieldly and more useful than the English edition, and with its companion work—“ Reynolds’ System of Medi- cine”—will well represent the present state of our science. One who is familiar with those two works will be fairly well furnished head-wise and hand- wise.— The Medical News, Jan. 7, 1882. STIMSOW, LEWIS A., B. A., M. JO., A Practical Treatise on Fractures. In one very handsome octavo volume of 598 pages, with 360 beautiful illustrations. Cloth, $4.75 ; leather, $5.75. Professor of Pathological Anatomy at the University of the City of New York, Surgeon and Curator to Bellevue Hospital, Surgeon to (he Presbyterian hospital, New York, etc. The author has given to the medical profession in this treatise on fractures what is likely to be- come a standard work on the subject. It is certainly not surpassed by any work written in the English, or, for that matter, any other language. The au- thor tells us in a short, concise and comprehensive manner, all that is known about his subject. There is nothing scanty or superficial about it, as in most other treatises; on the contrary, everything is thor- ough. The chapters on repair of fractures and their treatment show him not only to be a profound stu- dent, but likewise a practical surgeon and patholo- gist. His mode of treatment of the different fract- ures is eminently sound and practical. We consider this work one of the beat on fractures; and it will be welcomed not only as a text-book, but also by the ■ surgeon in full practice.—N. O. Medical and Surgical Journal, March, 1883. The author gives i n clear language all that the practical surgeon need know of the science ot fractures, their etiology, symptoms, processes of union, and treatment, according to the latest de- velopments. On the basis of mechanical analysis the author accurately and clearly explains the clinical features of fractures, and by the same method arrives at the proper diagnosis snd rational treatment. A thorough explanation of the patho- logical anatomy and a careful description of the various methods of procedure make the book full of value for every practitioner.—Oentralblatt fur Chirurgie, May 19,1883. MAIiSH, HOW ABB, E. B. C. S., Senior Assistant Surgeon to and Lecturer on Anatomy at St. Bartholomew's Hospital, London. Diseases of the Joints, In one 12mo. volume. Preparing. See Series of Clinical Manuals, page 3. TICK, T. PICKEBIXG, E. B. C. S., Surgeon to and Lecturer on Surgery at St. George's Hospital, London. Fractures and Dislocations, In one 12mo. volume. Preparing. See Series of Clinical Manuals, page 3. Lea Brothers & Co.’s Publications—Frac., Disloc., Oplithal. 23 HAMILTON, THANK H., M. I)., XX. JJ., Surgeon to Bellevue Hospital, New York. A Practical Treatise on Fractures and Dislocations. Seventh edition, thoroughly revised and much improved. In one very handsome octavo volume of 998 pages, with 379 illustrations. Cloth, $5.50; leather, $6.50; very handsome half Russia, open back, $7.00. Just ready. Hamilton’s great experience and wide acquaint- ance with the literature of the subject have enabled him to complete the labors of Malgaigne and to place the reader in possession of the advances made during thirty years. The editions have fol- lowed each other rapidly, and they introduce us to the methods of practice, often so wise, of his American colleagues. More practical than Mal- gaigne’s work, it will serve as a valuable guide to the practitioner in the numerous and embarrass- ing cases which come under his observation.— Archives Centrales de Midecine, Paris, Nov. 1884. This work, which, since its first appearance twenty-five years ago, has gone through many editions, and been much enlarged, may now be fairly regarded as the authoritative book of refer- ence on the subjects of fractures and dislocations. Each successive edition has been rendered of greater value through the addition of more re- cent work, and especially of the recorded re- searches and improvements made by the author himself and his countrymen.—British Medical Journal, May 9, 1885. With its first appearance in 1859, this work took rank among the classics in medical literature, and has ever since been quoted by surgeons the world over as an authority upon the topics of which it treats. The surgeon, if one can be found who does not already know the work, will find it scientific, forcible and scholarly in text, exhaustive in detail, and ever marked by a spirit of wise con- servatism.—Louisville Medical News, Jan. 10,1885. For a quarter of a century the author has been elaborating and perfecting his work, so that it now stands as the best of its kind in any lan- guage. As a text-book and as a book of reference and guidance for practitioners it is simply invalu- able.—New Orleans Med. andSurg. Jourril, Nov. 1884. JULER, HENRY E., T. R. C. S., Senior Ass't Surgeon, Royal Westminster Ophthalmic Hasp.; late Clinical Ass't, Moorfields, London. A Handbook of Ophthalmic Science and Practice. In one handsome octavo volume of 460 pages, with 125 woodcuts, 27 colored plates, and selections from the Test-types of Jaeger and Snellen. Cloth, $4.50 ; leather, $5.50. Just ready. This work is distinguished by the great num- ber of colored plates which appear in it for illus- trating various pathological conditions. They are very beautiful in appearance, and have been executed with great care as to accuracy. An ex- amination of the work shows it to be one of high standing, one that will be regarded as an authority among ophthalmologists. The treatment recom- mended is such as the author has learned from actual experience to be the best.—Cincinnati Medi- cal News, Dee. 1884. It presents to the student concise descriptions and typical illustrations of all important eye affections, placed in juxtaposition, so as to be frasped at a glance. Beyond a doubt it is the est illustrated handbook of ophthalmic science which has ever appeared. Then, what is still better, these illustrations are nearly ail original. We have examined this entire work with great care, and it represents the commonly accepted views of advanced ophthalmologists. We can most heartily commend this book to all medical stu- dents, practitioners and specialists. Detroit Lancet, Jan. 1885. WELLS, J. SOEEBERG, T. R. C. S., Professor of Ophthalmology in King's College Hospital, London, etc. A Treatise on Diseases of the Eye. Fourth American from the third London edition. Thoroughly revised, with copious additions, by Charles S. Bull, M. D., Surgeon and Pathologist to the New York Eye and Ear Infirmary. In one large octavo volume of 822 pages, with 257 illustrations on wood, six colored plates, and selections from the Test- types of Jaeger and Snellen. Cloth, $5.00; leather, $6.00; half Russia, $6.50. The present edition appears in less than three years since the publication of the last American edition, and yet, from the numerous recent inves- tigations that have been made in this branch of medicine, many changes and additions have been required to meet the present scope of knowledge upon this subject. A critical examination at once shows the fidelity and thoroughness with which the editor has accomplished his part of the work, The illustrations throughout are good. This edi- tion can be recommended to all as a complete treatise on diseases of the eye, than which proba- bly none better exists.—Medical Record, Aug. 18,’53, NETTLE SLUT, EDWARD, E. B. C. S., The Student’s Guide to Diseases of the Eye. Second edition. With a, chap- ter on the Detection of Color-Blindness, by William Thomson, M. D., Ophthalmologist to the Jefferson Medical College. In one royal 12mo. volume of 416 pages, with 138 illustrations. Cloth, $2.00. Ophthalmic Surg. and Lect. on Ophth. Surg. at St. Thomas' Hospital, London. This admirable guide bids fair to become the favorite text-book on ophthalmic surgery with stu- dents and general practitioners. It bears through- out the imprint of sound judgment combined with vast experience. The illustrations are numerous and well chosen. This book, within the short com- pass of about 400 pages, contains a lucid exposition of the modern aspect of ophthalmic science.— Medical Record, June 23, 1883. BJIOWNE, EDO All A., How to Use the Ophthalmoscope. Being Elementary Instructions in Oph- thalmoscopy, arranged for the use of Students. In one small royal 12mo. volume of 116 pages, with 35 illustrations. Cloth, $l.OO. Surgeon to the Liverpool Eye and Ear Infirmary and to the Dispensary for Skin Diseases. LAWSON ON INJURIES TO THE EYE, ORBIT AND EYELIDS: Their Immediate and Remote Effects. 8 vo., 404 pp., 92 illus. Cloth, $3.50. LAURENCE AND MOON’S HANDY BOOK OF OPHTHALMIC SURGERY, for the use of Prac- titioners. Second edition. In one octavo vol- ume of 227 pages, with 65 illust. Cloth, $2.75. CARTER’S PRACTICAL TREATISE ON DISEAS- ES OF THE EYE. Edited by John Green, M.D. In one handsome octavo volume. 24 Lea Brothers & Co.’s Publications—Otol., Urin. Dis.,Dent. B URNETT, CHARLES H., A. M,, M. IK, Professor of Otology in the Philadelphia Polyclinic ; President of the American Otological Society. The Ear, Its Anatomy, Physiology and Diseases. A Practical Treatise for the use of Medical Students and Practitioners. New (second) edition. In one handsome octavo volume of 580 pages, with 107 illustrations. Cloth, $4.00 ; leather, $5.00. Just ready. We note with pleasure the appearance of a second edition of this valuable work. When it first came out it was accepted by the profession as one of the standard works on modern aural surgery in the English language; and in his second edition Dr. Burnett has fully maintained his reputation, for the book is replete with valuable information and suggestions. The revision has been carefully carried out, and much new matter added. Dr. Burnett’s work must be regarded as a very valua- ble contribution to aural surgery, not only on account of its comprehensiveness, but because it contains the results of the careful personal observa- tion and experience of this eminent aural surgeon. —London Lancet, Feb. 21, 1885. EOLITZER, ARAM, A Text-Book of the Ear and its Diseases. Translated, at the Author’s re- quest, by James Patterson Cassells, M. D., M. E. C. S. In one handsome octavo vol- ume of 800 pages, with 257 original illustrations. Cloth, $5.50. Imperials Royal Prof, of Aural Therap. in the Univ. of Vienna. The work itself we do not hesitate to pronounce the best upon the subject of aural diseases which has ever appeared, systematic without being too diffuse on obsolete subjects, and eminently prac- tical in every sense. The anatomical descriptions of each separate division of the ear are admirable, and profusely illustrated by woodcuts. They are followed immediately by the physiology of the section, and this again by the pathological physi- ology, an arrangement which serves to keep up the interest of the student by showing the direct ap- plication of what has preceded to the study of dis- ease. The whole work can be recommended as a reliable guide to the student, and an efficient aid to the practitioner in his treatment.—Boston Med- ical and Surgical Journal, -June 7, 1883. ROBERTS, WILLIAM, M. I)., A Practical Treatise on Urinary and Renal Diseases, including Uri- nary Deposits. Fourth American from the fourth London edition. In one hand- some octavo volume of 609 pages, with 81 illustrations. Cloth, $3.50. Just ready. Lecturer on Medicine in the Manchester School of Medicine, etc. The peculiar value and finish of the book are derived from its resolute maintenance of a clinical and practical character. This volume is an un- rivalled exposition of everything which relates I directly or indirectly to the diagnosis, prognosis j and treatment of urinary diseases, and possesses j a completeness not found elsewhere in our lan- | guage.—2Vie Medical Chronicle, July, 1885. q. GROSS, S. IK, M. IK, LL. JO., I). C. L., etc. A Practical Treatise on the Diseases, Injuries and Malformations of the Urinary Bladder, the Prostate Gland and the Urethra. Third edition, thoroughly revised by Samuel W. Gross, M. D., Professor of the Principles of Surgery and of Clinical Surgery in the Jefferson Medical College, Philadelphia. In one octavo volume of 574 pages, with 170 illustrations. Cloth, $4.50. MORRIS, HENRY, M. 8., E. R. C. S., Surgeon to and Lecturer on Surgery at Middlesex Hospital, London. Surgical Diseases of the Kidney. In one 12mo. volume. Preparing. See Series of Clinical Manuals, page 3. LEG AS, CLEMENT, M. 8., B. S., E. R. C. S,, Diseases of the Urethra. In one 12mo. volume. Preparing. See Series of Clinical Manuals, page 3. Senior Assistant Surgeon to Guy's Hospital, London. THOMPSON, SIR HENRY, Lectures on Diseases of the Urinary Organs. Second American from the third English edition. In one Bvo. volume of 203 pp., with 25 illustrations. Cloth, $2.25. Surgeon and Professor of Clinical Surgery to University College Hospital, London. On the Pathology and Treatment of Stricture of the Urethra and Urinary Fistulse. From the third English edition. In one octavo volume of 359 pages, with 47 cuts and 3 plates. Cloth, $3.50. the Same Author. COLEMAN, A., L. R. C. I\, F. R. C, S,, Exam. L. I). S., Senior Pent. Surg. and Lect. on Dent. Surg. at St. Bartholomew's Hosp. and the Dent. Hosp., London. A Manual of Dental Surgery and Pathology. Thoroughly revised and adapted to the use of American Students, by Thomas C. Stellwagen, M. A., M. D., D. D. S., Prof, of Physiology at the Philadelphia Dental College. In one handsome octavo volume of 412 pages, with 331 illustrations. Cloth, $3.25. BASHAM ON RENAL DISEASES: A Clinical | Guide to their Diagnosis and Treatment. In | one 12mo. vol. of 304 pages, with 21 illustrations. Cloth, |2.00. Lea Brothers & Co.’s Publications—Venereal, Impotence. 25 JiUMSTUAD, F. J., M. I),, LL. />., and TAYLOB, B. W A. M., M. D., Late Professor of Venereal Diseases at the College of Physicians and Surgeons, New York, etc. Surgeon to Charity Hospital, New York, Prof, of Venereal and Skin Diseases in the University of Vermont, Pres, of the Am. Dermatological Ass’n. The Pathology and Treatment of Venereal Diseases. Including the results of recent investigations upon the subject. Fifth edition, revised and largely re- written, by Dr. Taylor. In one large and handsome octavo volume of 898 pages with 139 illustrations, and thirteen chromo-lithographic figures. Cloth, $4.75 ; leather, $5.75 ; very handsome half Russia, $6.25. It is a splendid record of honest labor, wide research, just comparison, careful scrutiny and original experience, which will always be held as a high credit to American medical literature. This is not only the best work in the English language upon the subjects of which it treats, but also one which has no equa. in other tongues for its clear, comprehensive and practical" handling of its themes.—American Journal of the Medical Sciences, Jan, 1884. It is certainly the best single treatise on vene- real in our own, and probably the best in any lan- guage.—Boston Medical and Surgical Journal, April 3, 1884. The character of this standard work is so well known that it would be superfluous here to pass in review its general or special points of excellence. The verdict of the profession has been passed; it has been accepted as the most thorough and com- plete exposition of the pathology and treatment of venereal diseases in the language. Admirable as a model of clear description, an exponent of sound pathological doctrine, and a guide for rational and successful treatment, it is an ornamentto the medi- cal literature of this country. The additions made to the present edition are eminently judicious, from the standpoint of practical utility.—Journal oj Cutaneous and Venereal Diseases, Jan. 1884. COBNIL, V,, Syphilis, its Morbid Anatomy, Diagnosis and Treatment. Specially revised by the Author, and translated with notes and additions by J. Henry C. Simes, M. IX, Demonstrator of Pathological Histology in the University of Pennsylvania, and J. William White, M. D., Lecturer on Venereal Diseases and Demonstrator of Surgery in the University of Pennsylvania. In one handsome octavo volume of 461 pages, with 84 very beautiful illustrations. Cloth, $3.75. Professor to the Faculty of Medicine of Paris, and Physician to the Lour cine Hospital. The anatomical and histological characters of the hard and soft sore are admirably described. The multiform cutaneous manifestations of the disease are dealt with histologically in a masterly way, as we should indeed expect them to be, and the accompanying illustrations are executed carefully and well. 'The various nervous lesions which are the recognized outcome of the syphilitic dyscrasia are treated with care and consideration. Syphilitic epilepsy, paralysis, cerebral syphilis and locomotor ataxia are subjects full of interest; and nowhere in the whole volume is the clinical experience of the author or the wide acquaintance of the translators with medical literature more evident. The anat- omy, the histology, the pathology and the clinical features of syphilis are represented in this work in their best, most practical and most instructive form, and no one will rise from its perusal without the feeling that his grasp of the wide and impor- tant subject on which it treats is a stronger and surer one.— The London Practitioner, Jan. 1882. HUTCHINSON, JONATHAN, F. B. S., F. B. C. S., Consulting Surgeon to the London Hospital. Syphilis. In one 12mo. volume. Preparing. See Series of Clinical Manuals, page 3. GBOSS, SAMUEL W., A. M., M. I)., Professor of the Principles of Surgery and of Clinical Surgery in the Jefferson Medical College. A Practical Treatise on Impotence, Sterility, and Allied Disorders of the Male Sexual Organs. Second edition, thoroughly revised. In one very hand- some octavo volume of 168 pages, with 16 illustrations. Cloth, $1.50. The author of this monograph is a man of posi- tive convictions and vigorous style. This is justi- fied by his experience and by his study, which has gone hand in hand with his experience. In regard to the various organic and functional disorders of the male generative apparatus, he has had ex- ceptional opportunities for observation, and his book shows that he has not neglected to compare his own views with those of other authors. The result is a work which can be safely recommended to both physicians and surgeons as a guide in the treatment of the disturbances it refers to. It is the best treatise on the subject with which we are acquainted.—The Medical News, Sept. 1, 1883. This work will derive value from the high stand- ing of its author, aside from the fact of its passing so rapidly into its second edition. This is, indeed, a book that every physician will be glad to place in his library, to be read with profit to himself, and with incalculable benefit to his patient. Be- sides the subjects embraced in the title, which are treated of in their various forms and degrees, spermatorrhoea and prostatorrhoea are also fully considered. The work is thoroughly practical in character, and will be especially useful to the general practitioner.—Medical Record, Aug. 18, 1883. CULLEBIFB, A., & BUMSTEAD, F. J., M.H., LL.F>., Surgeon to the Sdpiial du Midi. Late Professor of Venereal Diseases in the College of Physicians and Surgeons, New York. An Atlas of Venereal Diseases. Translated and edited by Freeman J. Bum- stead, M. D. In one imperial 4to. volume of 328 pages, double-columns, with 26 plates, containing about 150 figures, beautifully colored, many of them the size of life. Strongly bound in cloth, $17.00. A specimen of the plates and text sent by mail, on receipt of 25 cts. HILL ON SYPHILIS AND LOCAL CONTAGIOUS j DISORDERS. In one Svovol. 0f479 p. Cloth, $3.25. LEE’S LECTURES ON SYPHILIS AND SOME j FORMS OF LOCAL DISEASE AFFECTING PRINCIPALLY THE ORGANS OF GENERA- TION. In one Bvo. vol. of 246 pages. Cloth, $2.25. 26 Lea Brothers & Co.’s Publications—diseases of Skin. HYDE, J. XEVIXS, A. 31., M. J)., Professor of Dermatology and Venereal Diseases in Rush Medical College, Chicago. A Practical Treatise on Diseases of the Skin. For the use of Students and Practitioners. In one handsome octavo volume of 570 pages, with 66 beautiful and elab- orate illustrations. Cloth, $4.25; leather, $5.25. The author has given the student and practi- 1 tioner a work admirably adapted to the wants of each. We can heartily commend the book as a valuable addition to our literature and a reliable guide to students and practitioners in their studies and practice.—Am. Journ. of Med. Sci., July, 1883. Especially to be praised are the practical sug- gestions as to what may be called the common- sense treatment of eczema. It is quite impossible to exaggerate the judiciousness with which the formulae for the external treatment of eczema are selected, and what is of equal importance, the full I and clear instructions for their use.—London Medi- cal Times and Gazette, July 28, 1883. The work of Dr. Hyde will be awarded a high position. The student of medicine will find it ! peculiarly adapted to his wants. Notwithstanding the extent of the subject to which it is devoted, yet it is limited to a single and not very large vol- ume. without omitting a proper discussion of the topics. The conciseness of the volume, and the setting forth of only what can be held as facts will also make it acceptable to general practitioners. —Cincinnati Medical News, Feb. 1883. The aim of the author has been to present to his readers a work not only expounding the most modern conceptions of his subject, but presenting what is of standard value. He has more especially devoted its pages to the treatment of disease, and by his detailed descriptions of therapeutic meas- ures has adapted them to the needs of the physi- cian in active practice. In dealing with these questions the author leaves nothing to the pre- sumed knowledge of the reader, but enters thor- oughly into the most minute description, so that one is not only told what should be done under given conditions but how to do it as well. It is therefore in the best sense “ a practical treatise.” That it is comprehensive, a glance at the index will show.—Maryland Medical Journal, July 7,1883, Professor Hyde has long been known as one of the most intelligent and enthusiastic representa- tives of dermatology in the west. His numerous contributions to the literature of this specialty have gained for him a favorable recognition as a careful, conscientious and original observer. The remarkable advances made in our knowledge of diseases of the skin, especially from the stand- point of pathological histology and improved methods of treatment, necessitate a revision of the older texfrbooks at short intervals in order to bring them up to the standard demanded by the march of science. This last contribution of Dr. Hyde is an effort in this direction. He has at- tempted, as he informs us, the task of presenting in a condensed form the results of the latest ob- servation and experience. A careful examinatiod of the work convinces us that he has aecomplishen his task with painstaking fidelity and with a cred- itable result.—Journal of Cutaneous and Venereal Diseases, June, 1883. FOX, T., 3LD., F.R. €. F., and FOX, T. C,, 8.A., 3f.11. C.S., Physician to the Department for Skin Diseases, University College Hospital, London. Physician for Diseases of the Skin to the Westminster Hospital, London. An Epitome of Skin Diseases. With Formulae. For Students and Prac- titioners. Third edition, revised and enlarged. In one very handsome 12mo. volume of 238 pages. Cloth, $1.25. The third edition of this convenient handbook calls for notice owing to the revision and expansion which it has undergone. The arrangement of skin diseases in alphabetical order, which is the method of classification adopted in this work, becomes a positive advantage to the student. The book is one which we can strongly recommend, not only to students but also to practitioners who require a compendious summary of the present state of dermatology.—British Medical Journal, July 2,1883. We cordially recommend Fox’s Epitome to those whose time is limited and who wish a handy manual to lie upon the table for instant reference. Its alphabetical arrangement is suited to this use, for ail one has to know is the name of the disease, and here are its description and the appropriate treatment at hand and ready for instant applica- tion. The present edition has been very carefully revised and a number of new diseases are de- scribed, while most of the recent additions to dermal therapeutics find mention, and the formu- lary at the end of the book has been considerably augmented.—The Medical News, December, 1883. MOIUirS, MALCOLM, M. JL, Skin Diseases; Including their Definitions, Symptoms, Diagnosis, Prognosis, Mor- bid Anatom}/ and Treatment. A Manual for Students and Practitioners. In one 12mo. volume of 316 pages, with illustrations. Cloth, $1.75. Joint Lecturer on Dermatology at St. Mary's Hospital Medical School, London. To physicians who would like to know something about skin diseases, so that when a patient pre- sents himself for relief they can make a correct diagnosis and prescribe a rational treatment, we unhesitatingly recommend this little book of Dr. Morris. The affections of the skin are described in a terse, lucid manner, and their several charac- teristics so plainly set forth that diagnosis will be easy. The treatment in each case is such as the experience of the most eminent dermatologists ad- vises.—Cincinnati Medical News, April, 1880." for clearness of expression and methodical ar- rangement is better adapted to promote a rational conception of dermatology—a branch confessedly difficult and perplexing to the beginner.—St. Louis Courier of Medicine, April, 1880. The writer has certainly given in a small compass a large amount of well-compiled information, and his little book compares favorably with any other which has emanated from England, while in many points he has emancipated himself from the stub- bornly adhered to errors of others of his country- men. There is certainly excellent material in the book which will well repay perusal.—Boston Med. and Surg. Journ., March, 1880. This is emphatically a learner’s book; for we can safely say, that in the whole range of medical literature there is no book of a like scope which WILSON, ERASMUS, F. R. In one handsome small octavo volume of 535 pages. Cloth, $3.50. The Student’s Book of Cutaneous Medicine and Diseases of the Skin. SILLIER, THOMAS, 31. D., Handbook of Skin Diseases; for Students and Practitioners. Second Ameri- can edition. In one 12mo. volume of 353 pages, with plates. Cloth, $2.25. Physician to the Skin Department of University College, London. Lea Brothers & Co.’s Publications—Dis. of Women. 27 AN AMERICAN SYSTEM OF GYNAECOLOGY. A System of Gynaecology, in Treatises by Various Authors. Edited by Matthew D. Mann, M. I)., Professor of Obstetrics and Gynaecology in the Uni- versity of Buffalo, Y. Y. In two handsome octavo volumes, richly illustrated. In active preparation. LIST OF CONTRIBUTORS. WILLIAM H. BAKER, M. D., FORDYCE BARKER, M. D., ROBERT BATTEY, M. D., SAMUEL C. BUSEY, M. D., HENRY F. CAMPBELL, M. D., HENRY C. COE, M. D., EDWARD W. JENKS, M. D., WILLIAM T. LUSK, M. D., MATTHEW D. MANN, M. D., ROBERT B. MAURY, M. D., PAUL F. MUNDE, M. D., C. D. PALMER, M. D., WILLIAM M. POLK, M. D., THADDEUS A. REAMY, M. D., A. D. ROCKWELL, M. D., ALEX. J. C. SKENE, M. D., B. 0. DUDLEY, M. D., GEORGE J. BNGELMANN, M. D., HENRY F. GARRIGUES, M. D., WILLIABI GOODELL, M. D., EGBERT H. GRANDIN, M. D., SAMUEL W. GROSS, M. D., JAMBS B. HUNTER, M. D., A. REEVES JACKSON, M. D., R. STANSBURY SUTTON, A. M., M. D., T. GAILLARD THOMAS, M. D., ELI VAN DB WALKER, M. D., W. GILL WYLIE, M. D. THOMAS, T. GAILLARD, M. D., Professor of Diseases of Women in the College of Physicians and Surgeons, N. F. A Practical Treatise on the Diseases of Women. Fifth edition, thoroughly revised and rewritten. In one large and handsome octavo volume of 810 pages, with 266 illustrations. Cloth, $5.00; leather, $6.00; very handsome half Russia, raised bands, $6.50. The words which follow “fifth edition” are in this case no mere formal announcement. The alterations and additions which have been made are both numerous and important. The attraction and the permanent character of this book lie in the clearness and truth of the clinical descriptions of diseases; the fertility of the author in thera- vious one. As a book of reference for the busy practitioner it is unequalled.—Boston Medical any Surgical Journal, April 7,1880. It has been enlarged and carefully revised. It is a condensed encyclopedia of gynecological medi- cine. The style of arrangement, the masterly manner in which each subject is treated, and the peutic resources and the fulness with which the details of treatment are described; the definite character of the teaching; and last, but not least, the evident candor which pervades it. We would also particularize the fulness with which the his- tory of the subject is gone into, which makes the book additionally interesting and gives it value as a work of reference.—London Medical Times and Gazette, July 30, 1881. honest convictions derived from probably the largest clinical experience in that specialty of any in this country, all serve to commend it in the highest terms to the practitioner.—Nashville Jour, of Med. and Surg., Jan. 1881. That the previous editions of the treatise of Dr. Thomas were thought worthy of translation into German, French, Italian and Spanish, is enough to give it the stamp of genuine merit. At home it has made its way into the library of every obstet- rician and gynaecologist as a safe guide to practice. No small number of additions have been made to the present edition to make it correspond to re- cent improvements in treatment.—Pacific Medical and Surgical Journal, Jan. 1881. The determination of the author to keep his book foremost in the rank of works on gynaecology is most gratifying. Recognizing the fact that this can only be accomplished by frequent and thor- ough revision, he has spared no pains to make the present edition more desirable even than the pre- EDIS, ARTHUR W., M. D., Loud., F. R. C. F., M. R. C. S., The Diseases of Women. Including their Pathology, Causation, Symptoms, Diagnosis and Treatment. A Manual for Students and Practitioners. In one handsome Assist. Obstetric Physician to Middlesex Hospital, late Physician to British Lying-in Hospital. octavo volume of 576 pages, with 148 illustrations. Cloth, $3.00; leather, $4.00. It is a pleasure to.read a book so thoroughly good as this one. The special qualities which are conspicuous are thoroughness in covering the whole ground, clearness of description and con- ciseness of statement. Another marked feature of the book is the attention paid to the details of many minor surgical operations and procedures, as, for instance, the use of tents, application of leeches, and use of hot water injections. These are among the more common methods of treat- ment, and yet very little is said about them in many of the text-books. The book is one to be warmly recommended especially to students and general practitioners, who need a concise but com- plete rtsumt of the whole subject. Specialists, too, will find many useful hints in its pages.—Boston Med. and Surg. Journ., March 2, 1882. The greatest pains have been taken with the sections relating to treatment. A liberal selection of remedies is given for each morbid condition, the strength, mode of application and other details being fully explained. The descriptions of gynae- cological manipulations and operations are full, clear and practical. Much care has also been be- stowed on the parts of the book which deal with diagnosis—we note especially the pages dealing with the differentiation, one from another, of the different kinds of abdominal tumors. The prac- titioner will therefore find in this book the kind of knowledge he most needs in his daily work, and he will be pleased with the clearness and fulness of the Information there given.—The Practitioner, Feb. 1882. BABJSBS, BOBBBT, M. J)., P. B. C. P., A Clinical Exposition of the Medical and Surgical Diseases of Women. In one handsome octavo volume, with numerous illustrations. Yew edition. Preparing. Obstetric Physician to St. Thomas' Hospital, London, etc. WEST, CHARLES, M. D. Lectures on the Diseases Of Women. Third American from the third Lon- don edition. In one octavo volume of 543 pages. Cloth, $3.75; leather, $4.75. 28 Lea Brothers & Co.’s Publications—Dis. of Women, Midwty, EMMET, THOMAS ADDIS, M. D., LI. D., Surgeon to the Woman’s Hospital, New York, etc. The Principles and Practice of Gynaecology; For the use of Students and Practitioners of Medicine. New (third) edition, thoroughly revised. In one large and very handsome octavo volume of 880 pages, with 150 illustrations. Cloth, $5; leather, $6. (Just ready.) We are in doubt whether to congratulate the author more than the profession upon the appear- ance of the third edition of this well-known work. Embodying, as it does, the life-long experience of one who has conspicuously distinguished himself as a bold and successful operator, and who has devoted so much attention to the specialty, we feel sure the profession will not fail to appreciate the privilege thus offered them of perusing the views and practice of the author. His earnestness of purpose and conscientiousness are manifest. He gives not only his individual experience but endeavors to represent the actual state of gyme- cological science and art.—British Medical Jour- nal, May 16, 1885. No jot or tittle of the high praise bestowed upon the first edition is abated. It is still a book of marked personality, one based upon large clinical experience, containing large and valuable ad- ditions to our knowledge, evidently written not only with honesty of purpose, but with a conscien- tious sense of responsibility, and a book that is at once a credit to its author "and to American med- ical literature. We repeat that it is a book to be studied, and one that is indispensable to every practitioner giving any attention to gynaecology.— American Journal of the Medical Sciences, April, 1885. The time has passed when Emmet’s Gynaecology was to be regarded as a book for a single country or for a single generation. It has always been his aim to popularize gynseeology, to bring it within easy reach of the general practitioner. The orig- inality of the ideas, aside from the perfect con- fidence which we feel in the author’s statements, compels our admiration and respect. We may well take an honest pride in Dr. Emmet’s wort and feel that his book can hold its own against the criticism of two continents. It represents all that is most earnest and most thoughtful in American gynaecology. Emmet’s work will continue to reflect the individuality, the sterling integrity and the kindly heart of its honored author long after smaller books have been forgotten.—American Journal of Obstetrics, May, 1885. Any work on gynaecology by Emmet must always have especial interest and value. He has for many years been an exceedingly busy prac- titioner in this department. Few men have had his experience and opportunities. As a guide either for the general practitioner or specialist, it is second to none other. No one can read Emmet without pleasure, instruction and profit. —Cincinnati Lancet and Clinic, Jan 31, 1885. DUNCAN, J. MATTHEWS, M.D., LI. D., E. It. S. E., etc Clinical Lectures on the Diseases of Women; Delivered in Saint Bar- tholomew’s Hospital. In one handsome octant volume of 175 pages. Cloth, $1.50. They are in every way worthy of their author;' indeed, we look upon them as among the most valuable of his contributions. They are all upon matters of great interest to the general practitioner. Some of them deal with subjects that are not, as a rule, adequately handled in the text-books; others of them, while bearing upon topics that are usually treated of at length in such works, yet bear such a 'stamp of individuality that, if widely read, as they certainly deserve to be, they cannot foil to exert a wholesome restraint upon the undue eagerness with which many young physicians seem bent upon following the wild teachings which so infest the gynaecology of the present day.—N. Y. Medical Journal, March, 1880. MAY, CJSTABJLJES lIM. D. Late House Surgeon to Mount Sinai Hospital, New York. A Manual of the Diseases of Women. Containing a concise and systematic exposition of theory and practice. In one 12mo. volume of about 350 pages. In press. HODGE, HUGH L., M. D., Emeritus Professor of Obstetrics, etc., in the University of Pennsylvania. On Diseases Peculiar to Women; Including Displacements of the Uterus. Second edition, revised and enlarged. In one beautifully printed octavo volume of 519 pages, with original illustrations. Cloth, $4.50. the Same Author. The Principles and Practice of Obstetrics. Illustrated with large litho- graphic plates containing 159 figures from original photographs, and with numerous wood- cuts. In one large quarto volume of 542 double-columned pages. Strongly bound in cloth, $14.00. * * * Specimens of the plates and letter-press will be forwarded to any address, free by mail, on receipt of six cents in postage stamps. RAM SHOT HAM, FRANCIS 11., M. D. The Principles and Practice of Obstetric Medicine and Surgery; In reference to the Process of Parturition. A new and enlarged edition, thoroughly revised by the Author. With additions by W. Y. Keating, M. I)., Professor of Obstetrics, etc., in the Jefferson Medical College of Philadelphia. In one large and handsome imperial octavo volume of 640 pages, with 64 full-page plates and 43 woodcuts in the text, contain- ing in all nearly 200 beautiful figures. Strongly bound in leather, with raised bands, $7. ASHWELL’S PRACTICAL TREATISE ON THE DISEASES PECULIAR TO WOMEN. Third American from the third and revised London edition. In one Bvo. vol., pp. 520. Cloth, 53.50. I CHURCHILL ON THE PUERPERAL FEVER j AND OTHER DISEASES PECULIAR TO WO- MEN. In one Bvo. vol. 0f.464 pages. Cloth, $2.50. MEIGS ON THE NATURE, SIGNS AND TREAT- MBNT OF CHILDBED FEVER. In one Bvo. volume of 346 pages. Cloth, $2.00. Lea Brothers & Co.’s Publications—Midwifery. 29 PLAYFAIR, W. S., 31. T>., F. R. C. P., A Treatise on the Science and Practice of Midwifery. New (fourth) American, from the fifth English edition. Edited, with additions, by Robert P. Har- ris, M. D. In one handsome octavo volume of 654 pages, with 3 plates and 201 engrav- ings Cloth, $4; leather, |5 ; half Russia, $5.50. Just ready. Professor of Obstetric Medicine in King's College, London, etc. This excellent work needs no commendation. For many years it has maintained a deservedly high reputation among teachers as a text book, and in the profession as a guide to the practical experiences which attend the obstetrician. The present edition, under the supervision of Dr. Har- ris, has been carefully revised, and many portions rewritten, and the whole work has been adapted to the wants and circumstances of this continent.— Buffalo Medical and Surgical Journal, Aug. 1885. q. In the short time that this excellent and highly esteemed work has been before the profession it has reached a fourth edition in this country and a fifth one in England. This fact alone speaks in high praise of it, and it seems to us that scarcely more need be said of it in the way of endorsement of its value. As a text book for students and for the uses of the general practitioner there is no work on obstetrics superior to the work of Dr. Playfair. Its teachings are practical, written in plain language, and afford a correct understanding of the art of midwifery. No one can be disap- pointed in it.—Cincinnati Medical News, June, 1885. BABNJES, BOBEBT, M. J)., and FAJSTCOJIBT, M. J)., Phys. to the General Lying-in Hasp., Land. Obstetric Phys. to St. Thomas' Hasp., Land. A System of Obstetric Medicine and Surgery, Theoretical and Clin- ical. For the Student and the Practitioner. The Section on Embryology contributed by Prof. Milnes Marshall. In one handsome octavo volume of about 1000 pages, profusely illustrated. Cloth, $5; leather, $6. In a few days. BARKER, FORDYCE, A. M., 31. I)., LL. I). Ed hi., Clinical Professor of Midwifery and the Diseases of Women in the Bellevue Hospital Medical College, New York, Honorary Fellow of the Obstetrical Societies of London and Edinburgh, etc., etc. Obstetrical and Clinical Essays. In one handsome 12mo. volume of about 300 pages. Preparing. KING, A. F. A., 31. D., Professor of Obstetrics and Diseases of Women in the Medical Department of the Columbian Univer- sity, Washington, D. C., and in the University of Vermont, etc. A Manual of Obstetrics. Second edition. In one very handsome 12mo. volume of 331 pages, with 59 illustrations. Cloth, $2.00. It must be acknowledged that this is just what it pretends to be—a sound guide, a portable epit- ome, a work in which only indispensable matter has been presented, leaving out all padding and chaff, and one in which the student will find pure wheat or condensed nutriment.—New Orleans Med- ical and Surgical Journal, May, 1884. In a series of short paragraphs and by a con- densed style of composition, the writer has pre- sented a great deal of what it is well that every obstetrician should know and be ready to practice or prescribe. The fact that the demand for the volume has been such as to exhaust the first edition in a little over a year and a half speaks well for its popularity.—American Journal of the Medical Sciences, April”, 1884. LAS J) IS, HENRY (i., A. 31., 31. I)., Professor of Obstetrics and the Diseases of Women in Starling Medical College, Columbus, O. The Management of Labor. In one handsome 12mo. volume of about 300 pages, with. 30 illustrations. Shortly. BARNFS, FANCOURT, 31. 1)., Obstetric Physician to St. Thomas' Hospital, London. A Manual of Midwifery for Midwives and Medical Students. In one royal 12mo. volume of 197 pages, with 50 illustrations. Cloth, $1.25. PARVIN, TJIEOPHILUS, 31. J)., LL. I)., Professor of Obstetrics and the Diseases of Women and Children in the Jefferson Medical College. A Treatise on Midwifery. In one very handsome octavo volume of about 550 pages, with numerous illustrations. In press. PARRY, JOHN S., 31. !)., Obstetrician to the Philadelphia Hospital, Vice-President of the Obstet. Society of Philadelphia. Extra-Uterine Pregnancy; Its Clinical History, Diagnosis, Prognosis and Treatment. In one handsome octavo volume of 272 pages. Cloth, $2.50. TANNER, THO3IAS HA WKES, 31. J). On the Signs and Diseases of Pregnancy. First American from the second English edition. Octavo, 490 pages, with 4 colored plates and 16 woodcuts. Cloth, $4.25. WLNCKFL, F. A Complete Treatise on the Pathology and Treatment of Childbed, For Students and Practitioners. Translated, with the consent of the Author, from the second German edition, by J. R. Chadwick, M. D. Octavo 484 pages. Cloth, $4.00. 30 Lea Brothers & Co.’s Publications—Midwly., Dis. CMLdn. LEISHMAN, WILLIAM, M. L>., A System of Midwifery, Including the Diseases of Pregnancy and the Puerperal State. Third American edition, revised by the Author, with additions by John S. Parry, M. D., Obstetrician to the Philadelphia Hospital, etc. In one large and very handsome octavo volume of 740 pages, with 205 illustrations. Cloth, $4.50; leather, $5.50; very handsome half Russia, raised hands, $6.00. Regius Professor of Midwifery in the University of Glasgow, etc. The author is broad in his teachings, and dis- cusses briefly the comparative anatomy of the pel- vis and the mobility of the pelvic articulations. The second chapter is devoted especially to the study of the pelvis, while in the third the female organs of generation are introduced. The structure and development of the ovum are admirably described. Then follow chapters upon the various subjects embraced in the study of mid- wifery. The descriptions throughout the work are plain and pleasing. It is sufficient to state that in this, the last edition of this well-known work, every recenl/advancement in this field has been brought forward.—Physician and Surgeon, Jan. 1880. We gladly welcome the new edition of this ex- cellent text-book of midwifery. The former edi- tions have been most favorably received by the profession on both sides of the Atlantic. In the preparation of the present edition the author has made such alterations as the progress of obstetri- cal science seems to require, and we cannot but admire the ability with which the task has been performed. We consider it an admirable text- book for students during their attendance upon lectures, and have great pleasure in recommend- ing it. As an exponent of the midwifery of the present day it has no superior in the English lan- guage.—Canada Lancet, Jan. 1880. To the American student the work before us must prove admirably adapted. Complete in all its parts, essentially modern in its teachings, and with demonstrations noted for clearness and precision, it will gain in favor and be recognized as a work of standard merit. The work cannot fall to be popular and is cordially recommended.—N. O. Med. and Surg. Journ., March, 1880. SMITH, J. LEWIS, M. D., Clinical Professor of Diseases of Children in the Bellevue Hospital Medical College, N. Y. A Complete Practical Treatise on the Diseases of Children. Fifth edition, thoroughly revised and rewritten. In one handsome octavo volume of 836 pages, wi th illustrations. Cloth, $4.50; leather, $5.50; very handsome half Russia, raised bands, $6. This is one of the best books on the subject with which we have met and one that has given us satisfaction on every occasion on which we have consulted it, either as to diagnosis or treatment, It is now in its fifth edition and in its present form is a very adequate representation of the subject it treats of as at present understood. The important subject of infant hygiene is fully dealt with in the early portion of the book. The great bulk of the work is appropriately devoted to the diseases of infancy and childhood. We would recommend i any one in need of information on the subject to j procure the work and form his own opinion on it, 1 which we venture to say will be a favorable one.— Dublin Journal of Medical Science, March, 1883. There is no book published on the subjects of which this one treats that is its equal in value to the physician. While he has said just enough to impart the information desired by general praeti- tioners on such questions as etiology, pathology, prognosis, etc., he has devoted more attention to the diagnosis and treatment of the ailments which he so accurately describes ; and such information is exactly what is wanted by the vast majority of “ family physicians.”— Va. Med. Monthly, Feb. 1882. KEATING, JOHN M., M. I)., The Mother’s Guide in the Management and Feeding of Infants. In one handsome 12mo. volume of 118 pages. Cloth, $l.OO. Lecturer on the Diseases of Children at the University of Pennsylvania, etc. Works like this one will aid the physician im- mensely, for it saves the time he is constantly giv- ing his patients in instructing them on the sub- jects here dwelt upon so thoroughly and prac- tically. Dr. Keating has written a practical hook, has carefully avoided unnecessary repetition, and successfully instructed the mother in such details of the treatment of her child as devolve upon her. He has studiously omitted giving prescriptions, and instructs the mother when to call upon the doctor, as his duties are totally distinct from hers. —American Journal of Obstetrics, October, 1881. Dr. Keating has kept clear of the common fault of works of this sort, viz., mixing the duties of the mother with those proper to the doctor. There is the ring of common sense in the remarks about the employment of a wet-nurse, about the proper food for a nursing mother, about the tonic effects of a bath, about the perambulator versus the nurses, arms, and on many other subjects concerning which the critic might say, “surely this is obvi- ous,” but which experience teaches us are exactly the things needed to be insisted upon, with the rich as well as th e poor.—London Lancet, January, 281882. A book small in size, written in pleasant style, in language which can be readily understood by any mother, and eminently practical and safe; in fact a book for which we have been waiting a long time, and which we can most heartily recommend to mothers as the book on this subject.—New York Medical Journal and Obstetrical Review, Feb. 1882. OWEN, EDMUND, M. 8,, Surgeon to the Children's Hospital, Great Ormond St., London. E. R. C. S., Surgical Diseases of Children. In one 12mo. volume. Preparing. See Series of Clinical, Manuals, page 3. WEST, CHARLES, M. />., Lectures on the Diseases of Infancy and Childhood. Fifth American from 6th English edition. In one octavo volume of 686 pages. Cloth, $4.50; leather, $5.50. Physician to the Hospital for Sick Children, London, etc. the Same Author. On Some Disorders of the Nervous System in Childhood. In one small 12mo. volume of 127 pages. Cloth, $l.OO. CONDIE’S PRACTICAL TREATISE ON THE DISEASES OP CHILDREN. Sixth edition, re- I vised and augmented. In one octavo volume of ( 779 pages. Cloth, $5.25; leather, $6.26. Lea Brothers & Co.’s Publications—Med. Juris., Mlscel. 31 TIDY, CHARLES M EYMOTT, M. li., F. C. 8,, Professor of Chemistry and of Forensic Medicine and Public Health at the London Hospital, etc. Legal Medicine. Volume 11. Legitimacy and Paternity, Pregnancy, Abor- tion, Eape, Indecent Exposure, Sodomy, Bestiality, Live Birth, Infanticide, Asphyxia, Drowning, Hanging, Strangulation, Suffocation. Making a very handsome imperial oc- tavo volume of 529 pages. Cloth, $6.00; leather, $7.00. Volume I. Containing 664 imperial octavo pages, with twm beautiful colored plates. Cloth, $6.00; leather, $7.00. The satisfaction expressed with the first portion of this work is in no wise lessened by a perusal of the second volume. We find it characterized by the same fulness of detail and clearness of ex- pression which we had occasion so highly to com- mend in our former notice, and which render it so valuable to the medical jurist. The copious tables of cases appended to each division of the subject, must have cost the author a prodigious amount of labor and research, but they constitute one of the most valuable features of the book, especially for reference in medico-legal trials.— American Journal of the Medical Sciences, April, 1884. TAYLOR, ALFRED 8., M. D., Lecturer on Medical Jurisprudence and Chemistry in Guy's Hospital, London. A Manual of Medical Jurisprudence. Eighth American f A Manual of Medical Jurisprudence. Eighth American from the tenth Lon- don edition, thoroughly revised and rewritten. Edited by John J. Keese, M. D., Professor of Medical Jurisprudence and Toxicology in the University of Pennsylvania. In one large octavo volume of 937 pages, with 70 illustrations. Cloth, $5.00 ; leather, $6.00; half Russia, raised bands, $6.50. P-The American editions of this standard manual have for a long time laid claim to the attention of the profession in this country; and the eighth comes before us as embodying the latest thoughts and emendations of Dr. Taylor upon the subject to which he devoted his life with an assiduity and success which made him facile princeps among English writers on medical jurisprudence. Both the author and the book have made a mark too deep to be affected by criticism, whether it be censure or praise. In this case, however, we should only have to seek for laudatory terms.—American Journal of the Medical Sciences, Jan. 1881. This celebrated work has been the standard au- thority in its department for thirty-seven years, both in England and America, in both the profes- sions which it concerns, and it is improbable that it will be superseded in many years. The work is simply indispensable to every physician, and nearly so to every liberally-educated lawyer, and we heartily commend the present edition to both pro- fessions.—Albany Law Journal, March 26,1881. The Principles and Practice of Medical Jurisprudence. Third edition. In two handsome octavo volumes, containing 1416 pages, with 188 illustrations. Cloth, $10; leather, $l2. Just ready. the Same Author. For years Dr. Taylor was the highest authority in England upon the subject to which he gave especial attention. His experience was vast, his judgment excellent, and his skill beyond cavil. It is therefore well that the work of one who, as Dr. Stevenson says, had an “enormous grasp of all matters connected with the subject,” should be brought up to the present day and continued in its authoritative position. Tc accomplish this re- sult Dr. Stevenson has subjected it to most careful editing, bringing it well up to the times.—Ameri- can Journal of the Medical Sciences, Jan. 1884. the Same Author, Poisons in Relation to Medical Jurisprudence and Medicine. Third American, from the third and revised English edition. In one large octavo volume of 788 pages. Cloth, $5.50; leather, $6.50. PEPPER, AUGUSTUS JM. S., M. 8., F. R. C. S., Examiner in Forensic Medicine at the University of Lond.on. Forensic Medicine. In one pocket-size 12mo. volume. Preparing. See Students1 Series of Manuals, page 3. LFA, HENRY C. Superstition and Force: Essays on The Wager of Law, The Wager of Battle, The Ordeal and Torture. Third revised and enlarged edition. In one handsome royal 12mo. volume of 552 pages. Cloth, $2.50. This valuable work is in reality a history of civ- ilization as interpreted by the progress of jurispru- denee. . . In “Superstition and Force” we have a philosophic survey of the long period intervening between primitive barbarity and civilized enlight- enment. There is not a chapter in the work that should not be most carefully studied; and however well versed the reader may be in the science of jurisprudence, he will find much in Mr. Lea’s vol- ume of which he was previously ignorant. The book is a valuable addition to the literature of so- eial science.— Westminster Review, Jan. 1880. By the Same Author. Studies in Church History. The Rise of the Temporal Power—Ben- efit of Clergy—Excommunication. Hew edition. In one very handsome royal octavo volume of 605 pages. Cloth, $2.50. The author is pre-eminently a scholar. He takes up every topic allied with the leading theme, and traces it out to the minutest detail with a wealth of knowledge and impartiality of treatment that compel admiration. The amount of information compressed into the book is extraordinary. In no other single volume is the development of the Just ready. primitive church traced with so much clearness, and with so definite a perception of complex or conflicting sources. The fifty pages on the growth of the papacy, for instance, are admirable for con- ciseness and freedom from prejudice.—Boston Traveller, May 3,1883. Allen’s Anatomy . . . . '. 6 American Journal of the Medical Sciences . 3 American System of Gynaecology . . . 27 American System of Practical Medicine . 15 ♦Ashhurst’s Surgery ..... 20 Ashwell on Diseases of Women . . .28 Attfield’s Chemistry . . . . .9 Ball on the Rectum and Anus . . . 21 Barker’s Obstetrical and Clinical Essays, . 29 Barlow’s Practice of Medicine . . . 17 Barnes’ Midwifery . . . .29 ♦Barnes on Diseases of Women . . .27 Barnes’ System of Obstetric Medicine . . 29 Bartholow on Electricity .... 17 Basham on Renal Diseases .... 24 Bell’s Comparative Physiology and Anatomy . 3 7 Bellamy’s Operative Surgery . . . 3,20 Bellamy’s Surgical Anatomy ... 6 Blandford on Insanity . . .19 Bloxam’s Chemistry . ... 9 Bowman’s Practical Chemistry ... 9 ♦Bristowe's Practice of Medicine . . .14 Broadbent on the Pulse . . . . 3,16 Browne on the Ophthalmoscope . . .23 Browne on the Throat .... 18 Bruce’s Materia Medica and Therapeutics . 11 Brunton’s Materia Medica and Therapeutics . 11 Bryant on the Breast . . . . . 3,21 ♦Bryant’s Practice of Surgery . . .21 ♦Bumstead on Venereal Diseases . . .25 ♦Burnett on the Ear . . . . .24 Butlin on the Tongue . . . . .3,21 Carpenter on the Use and Abuse of Alcohol . 8 *Carpenter’s Human Physiology ... 8 Carter on the Eye ..... 23 Century of American Medicine . . .14 Chambers on Diet and Regimen . . .17 Charles’ Physiological and Pathological Chem. 10 Churchill on Puerperal Fever . . .28 Clarke and Lockwood’s Dissectors’ Manual . 6 Classen’s Quantitative Analysis . . .10 Cleland’s Dissector ..... 5 Clouston on Insanity .... 19 Clowes’ Practical Chemistry . • .10 Coats’ Pathology . . . . .13 Cohen on the Throat . . . . .18 Coleman’s Dental Surgery .... 24 Sondie on Diseases of Children . . .30 oopor’s Lectures on Surgery . . .20 Cornil on Syphilis .... .25 ♦Cornil and Ranvier’s Pathological Histology . 13 Cullerier’s Atlas of Venereal Diseases . . 25 Ournow’s Medical Anatomy . . . 3,6 Dalton on the Circulation .... 7 ♦Dalton’s HumanPhysiology ... 8 Dalton’s Topographical Anatomy of the Brain 7 Davis’ Clinical Lectures . . .16 Draper’s Medical Physics ... 7 Druitt’s Modern Surgery . . . .21 Duncan on Diseases of Women . . .28 ♦Dunglison’s Medical Dictionary ... 4 Bdis on Diseases of Women .... 27 Ellis’ Demonstrations of Anatomy . . 7 Emmet’s Gynaecology . . .28 ♦Erichsen’s System of Surgery . . .21 Esmarch’s Early Aid In Injuries and Accid’ts . 21 Farquharson’s Therapeutics and Mat. Med. . 12 Fenwick’s Medical Diagnosis . . .16 Finlayson’s Clinical Diagnosis . . .16 Flint on Auscultation and Percussion . . 18 Flint on Phthisis ..... 18 Flint on Physical Exploration of the Lungs . 18 Flint on Respiratory Organs . . .18 Flint on the Heart . ... 18 ♦Flint’s Clinical Medicine . . .16 Flint’s Essays . . . * . 16 ♦Flint’s Practice of Medicine . . .14 Folsom’s Laws of U. S. on Custody of Insane . 19 Foster’s Physiology ..... 8 ♦Fothergilrs Handbook of Treatment , . 16 Fownes’ Elementary Chemistry ... 8 Fox on Diseases of the Skin . ... 26 Frankland and Japp’s Inorganic Chemistry . 9 Fuller on the Lungs and Air Passages . . 18 Galloway’s Analysis , . ... 8 Qibney’s Orthopaedic Surgery . . .20 Gibson’s Surgery ..... 20 Gluge’s Pathological Histology, by Leidy . 13 Gould’s Surgical Diagnosis . . . . 3, 20 ♦Gray’s Anatomy ... ... 5 Greene’s Medical Chemistry . . . 10 Green’s Pathology and Morbid Anatomy . 13 Griffith’s Universal Formulary . . .11 Gross on Foreign Bodies in Air-Passages . 18 Gross on Impotence and Sterility . . . 25 Gross on Urinary Organs . . . 24 ♦Gross’ System of Surgery . . . 20 Habershon on the Abdomen . . 17 ♦Hamilton on Fractures and Dislocations . 23 Hamilton on Nervous Diseases . . .19 Hartshorne’s Anatomy and Physiology . . 6 Hartshorne’s Conspectus of the Med. Sciences . 3 Hartshorne’s Essentials of Medicine . . 14 Hermann’s Experimental Pharmacology . 11 Hill on Syphilis ...... 25 Hillier’s Handbook of Skin Diseases . . 26 Hoblyn’s Medical Dictionary ... 4 Hodge on Women . . . . .28 Hodge’s Obstetrics ..... 28 Hoffmann and Power’s Chemical Analysis 10 Holden’s Landmarks ..... 5 Holland’s Medical Notes and Reflections . 17 ♦Holmes’System of Surgery . . .22 Horner’s Anatomy and Histology . . 6 Hudson on Fever . . . ,14 Hutchinson on Syphilis . . . . 3,25 Hyde on the Diseases of the Skin . . .26 Jones (C. Handheld) on Nervous Disorders . 18 Jtiler’s Ophthalmic Science and Practice . 23 Keating on Infants , . . .30 King’s Manual of Obstetrics . . . . 29 Klein’s Histology . . . 3,13 Landis on Labor . . - . .29 La Roche on Pneumonia, Malaria, etc. . . 18 La Roche on Yellow Fever . . . .14 Laurence and Moon’s Ophthalmic Surgery . 23 Lawson on the Bye, Orbit and Eyelid . . 23 Lea’s Studies in Church History . . .31 Lea’s Superstition and Force . . .31 Lee on Syphilis . ... 25 Lehmann’s Chemical Physiology ... 8 ♦Leishman’s Midwifery . . . .30 Lucas on Diseases of the Urethra . . .3,24 Ludlow’s Manual of Examinations . . 3 Lyons on Fever ...... 14 Maisch’s Organic Materia Medica . . .11 Marsh on the Joints . . . .3,22 May on Diseases of Women .... 28 Medical News . . . . . 1 Meigs on Childbed Fever . . . .28 Miller’s Practice of Surgery . . . .20 Miller’s Principles of Surgery ... 20 Mitchell’s Nervous Diseases of Women . . 19 Morris on Diseases of the Kidneys . . 3,24 Morris on Skin Diseases . . . .26 Neill and Smith’s Compendium of Med. Sci. . 3 Nettleship on Diseases of the Eye . . . 23 Owen on Diseases of Children . . . 3, 30 * Parrish’s Practical Pharmacy ... 11 Parry on Extra-Uterine Pregnancy . .29 Parvin’s Midwifery . . ... .29 Pavy on Digestion and its Disorders . . 17 Pepper’s Forensic Medicine . . . . 8, 31 Pepper's Surgical Pathology . . . 3,13 Pick on Fractures and Dislocations . . 3,22 Pirrie’s System of Surgery .... 20 Playfair on Nerve Prostration and Hysteria . 19 ♦Playfair’s Midwifery . ... 29 Politzer on the Ear and its Diseases . . 24 Power’s Human Physiology . . . 3,8 Ralfe’s Clinical Chemistry . . . 3,10 Ramsbotham on Parturition . , 28 Remsen’s Theoretical Chemistry ... 9 ♦Reynolds’System of Medicine . . .15 Richardson’s Preventive Medicine . . 17 Roberts on Urinary Diseases- . . .24 Roberts’Principles and Practice of Surgery . 20 Robertson’s Physiological Physics . . 3,7 Rodwell’s Dictionary of Science ... 4 Ross on Nervous Diseases . . . .19 Sargent’s Minor and Military Surgery . , 20 Savage on Insanity, including Hysteria . . 3,19 Schafer’s Essentials of Histology, . . 13 Schafer’s Histology . 13 Schreiber on Massage . ... 16 Seiler on the Throat, Nose and Naso-Pharynx 18 Series of Clinical Manuals . . . .3 Simon’s Manual of Chemistry ... 9 Skey’s Operative Surgery .... 20 Slade on Diphtheria . .... 18 Smith (Edward) on Consumption . . .18 Smith (H. H.) and Horner’s Anatomical Atlas 5 ♦Smith (J. Lewis) on Children . . .30 Stille on Cholera . . . .16 ♦Stille & Maisch’s National Dispensatory . 12 ♦StillS’s Therapeutics and Materia Medica . 12 Stimson on Fractures . .... 22 Stimson’s Operative Surgery . . .20 Stokes on Fever ...... 14 Students’Series of Manuals .... 3 Sturges’ Clinical Medicine .... 16 Tanner on Signs and Diseases of Pregnancy . 29 Tanner’s Manual of Clinical Medicine . . 16 Taylor on Poisons . . . . .31 ♦Taylor’s Medical Jurisprudence . . .31 Taylor’s Prin. and Prac. of Med. Jurisprudence 31 ♦Thomas on Diseases of Women . . .27 Thompson on Stricture . . • . .24 Thompson on Urinary Organs ... 24 Tidy’s Legal Medicine. .... 31 Todd on Acute Diseases ... .17 Treves’ Applied Anatomy . . . . 3,6 Treves on Intestinal Obstruction . . . 3,21 Tuke on the Influence of Mind on the Body . 19 Walshe on the Heart ..... 18 Watson’s Practice of Physic . . . .14 * Wells on the Eye . . . .23 West on Diseases of Childhood . . .30 West on Diseases of Women . . .27 West on Nervous Disorders in Childhood . . 30 Williams on Consumption . . . .18 Wilson’s Handbook of Cutaneous Medicine . 26 Wilson's Human Anatomy .... 5 Wlnckel on Pathol, and Treatment of Childbed 29 Wohler’s Organic Chemistry ... 8 Woodhead’s Practical Pathology ... 13 Year-Rook of Treatment . . . .17 LEA BROTHERS & CO., Philadelphia. Books marked * are also bound in half Russia.