MANUAL O F CHEMISTRY. A GUIDE TO LECTURES AND LABORATORY WORK FOR BEGINNERS IN CHEMISTRY. A TEXT-BOOK SPECIALLY ADAPTED FOR STUDENTS OF MEDICINE AND PHARMACY. BY W. SIMON, Ph.D., M.D., PROFESSOR OF CHEMISTRY AND TOXICOLOGY IN THE COLLEGE OF PHYSICIANS AND SURGEONS; PROFESSOR OF CHEMISTRY AND ANALYTICAL CHEMISTRY IN THE MARYLAND COLLEGE OF PHARMACY, BALTIMORE, MD. FOURTH EDITION, THOROUGHLY REVISED. WITH FORTY-FOUR ILLUSTRATIONS AND SEVEN COLORED PLATES, REPRESENTING FIFTY-SIX CHEMICAL REACTIONS. PHILADELPHIxl: LEA BROTHERS & CO. 189 3. Entered according to Act of Congress, in the year 1893, by LEA BROTHERS & CO., In the Office of the Librarian of Congress at Washington, D. C. All rights reserved. DORN AN, PRINTER. PREFACE TO FOURTH EDITION. The demand for a new edition of this Manual within a twelvemonth affords sincere gratification to the author. To the reader it secures what- ever benefits may accrue from a prompt and thorough revision. In preparing the new edition the author has given careful attention to the kind suggestions of many teachers of chemistry who have regarded the work with favor, and he has furthermore made such changes and additions as were indicated by the advance of the science and by his increased experience of the needs of medical and pharmaceutical students. The fourth edition, therefore, differs from its predecessors not only in the greater quantity of matter, but also in some details of its presentation. While heartily in favor of the recommendations made by the American Association for the Advancement of Science regarding a uniform system of orthography and nomenclature in chemical terms, the author has felt impelled toward conservatism in the adoption of these changes. This seemed advisable because the new edition of the United States Pharma- copoeia will adopt only a portion of the proposed changes, and it would be unwise to have the student confronted by two different systems of orthography. As heretofore, the subject has been divided into seven parts, each one of which contains so much of the matter under consideration as is believed to be necessary for a fair understanding of the subject. At the same time care has been taken to place in the foreground all facts and data which are of direct interest to the physician and pharmacist, and to exclude, as far as is compatible with the presentation of a comprehensive view of chemistry, those portions which are of restricted interest only. In the first part the fundamental properties of matter are considered briefly, and to such an extent as is necessary for an understanding of chemical phenomena. The second part treats of those principles of chemistry which are the foundation of the science, and enters briefly into a discussion of theoretical views regarding the atomic constitution of matter. Though the author prefers to present these theories to his classes at the proper times during the course of lectures, he does not deem it desirable to have them scat- tered throughout the work, believing it better to assemble them compactly in print so that the student may be able to study them after having acquired some knowledge of chemical phenomena. The third and fourth parts are devoted to the consideration of the non- metallic and metallic elements and their compounds. While the periodic VI PREFACE TO FOURTH EDITION. law furnishes a most admirable basis for a scientific classification of ele- ments, yet their consideration according to a strict adherence to periodicity does not seem advisable in this book. For this reason the old classifica- tion of metals and non-metals, organic and inorganic compounds, ha3 been retained, since experience has shown it to be well adapted to the instruc- tion of beginners in chemistry. All our classifications of either natural objects or phenomena are imperfect, because Nature does not draw those distinct lines of demarcation which we adopt as necessary for our studies. The most simple and natural classification is therefore always to be pre- ferred, even if, as in the above case, the student might derive from it the impression that matter was thus separated into distinct groups. Of elements, those only are considered which have either intrinsically or in combination a practical interest, or which take an active part in the various chemical changes in nature. For the special benefit of pharmaceutical and medical students all chemicals mentioned in the United States Pharmacopoeia are included, and when of sufficient interest they are fully considered. The fifth part is devoted to analytical chemistry and will serve the student as a guide in his laboratory work. Qualitative methods are chiefly considered, but a chapter is added giving the principal methods for volumetric determinations. The sixth part treats of organic chemistry. Though it is impossible to include within the limits of this text-book an extended consideration of a branch of chemical science so highly developed, yet it is believed that an intelligent study of this part will familiarize the student with carbon compounds sufficiently to give him a clear understanding of their general character, and a knowledge of the bodies which are most important in medical science. The seventh and last part, giving some of the principal facts of physio- logical chemistry, was prepared for the benefit of the medical student in particular. A great deal of new matter has been added to these chapters, and special care has been taken to mention the most modern methods for chemical examination in clinical diagnosis. As an aid to laboratory work a number of experiments have been added which may readily be performed by students with a comparatively small outfit of chemical apparatus. The decimal system has been strictly adhered to in all weights and measures; degrees of temperature are expressed in the same system, the corresponding degrees of Fahrenheit being also mentioned. The plates showing the variously shaded colors of a number of chemi- cals and their reactions, that have proved so useful in previous editions, have been reproduced for this. Baltimore, March, 1893. w. s. CONTENTS. I. FUNDAMENTAL PROPERTIES OF MATTER. RESULTS OF THE ATTRACTION BETWEEN MASSES, SURFACES, AND MOLECULES. 1. Extension or figure. Matter—State of aggregation—Solids—Cohesion—Force—Crystal- lized, amorphous, polymorphous, isomorphous substances—Liquids— Gases—Law of Mariotte 17-21 2. Divisibility. Mechanical comminution—Action of heat on matter—Molecular theory—Law of Avogadro—Motion of molecules, heat—Melting, boil- ing, distillation, sublimation—Thermometers—Specific heat . . 21-28 3 Gravitation. Action of gravitation—Weight, specific weight—Hydrometers— Weight of gases—Barometer—Changes in the atmospheric pressure— Influence of pressure on state of aggregation ..... 28-31 4. Porosity. Nature of porosity—Surface, surface-action—Adhesion—Capillary attraction—Absorption—Diffusion—Dialysis—Indestructibility . . 32-36 PAGE II. PRINCIPLES OF CHEMISTRY. RESULTS OF THE ATTRACTION BETWEEN ATOMS. 5 Chemical divisibility. Decomposition by heat—Elements—Compound substances—Chem- ical affinity—Atoms—Chemistry—Atomic and molecular weight— Chemical symbols and formulas ...... . 37-42 6. Laws of chemical combination. Law of the constancy of composition—Law of multiple proportions —Law of chemical combination by volume—Law of equivalents— Quantivalence, valence 42-48 7. Determination of atomic and molecular weights. Determination of atomic weights by chemical decomposition, by means of specific weights of gases or vapors, by means of specific heat—Determination of molecular weights—Baoult’s law . . 48-53 viii CONTENTS. 8. Decomposition of compounds. Groups of compounds. Decomposition hy heat, light, and electricity—Mutual action of substances upon each other—The nascent state—Analysis and synthesis—Acid, basic, and neutral substances—Salts—Residues, radicals 53-60 9. General remarks regarding elements. Relative importance of different elements—Classification of ele- ments—Metals and non-metals—Natural groups of elements—Men- ■ delejeff’s periodic law—Physical properties of elements—Allotropic modifications—Relationship between elements and the compounds formed by their union—Nomenclature—Writing chemical equations —How to study chemistry 60-70 PAGE III. NON-METALS AND THEIR COMBINATIONS. Symbols, atomic weights, and derivation of names—Occurrence in nature—Time of discovery—Valence . . . . . . 71-72 10. Oxygen. History—Occurrence in nature—Preparation—Physical and chem- ical properties—Combustion—Ozone ...... 73-77 11. Hydrogen. History—Occurrence in nature—Preparation—Properties—Water —Mineral waters—Drinking-water—Distilled water—Hydrogen di- oxide 78-84 12. Nitrogen. Occurrence in nature—Preparation—Properties—Atmospheric air —Ammonia—Compounds of nitrogen and oxygen—Nitrogen mon- oxide—Nitric acid; tests for it . . . . . . . 84-91 13. Carbon. Occurrence in nature—Properties—Diamond—Graphite—Tests for carbon—Carbon dioxide—Carbonic acid—Tests for carbonic acid —Carbon monoxide—Compounds of carbon and hydrogen—Flame —Silicon—Silicic acid—Boron, boric acid; tests for it . . 92-99 14. Sulphur. Occurrence in nature—Properties—Crude, sublimed, washed, and precipitated sulphur—Sulphur dioxide—Sulphurous acid; tests for it—Sulphur trioxide—Sulphuric acid ; its manufacture and proper- ties—Tests for sulphates—Sulpho-acids — Pyrosulphuric acid — Thiosulphuric acid—Hydrogen sulphide ; tests for it—Disulphide of carbon—Selenium—Tellurium 100-108 15. Phosphorus. Occurrence in nature—Manufacture, properties, and modifications —Poisonous properties and detection in cases of poisoning—Oxides of phosphorus—Phosphorous acid ; tests for it—Metaphosphoric, pyrophosphoric, orthophosphoric acids; tests for them—Hypophos- phorous acid ; tests for it—Hydrogen phosphide .... 108-116 CONTENTS. IX 16. Chlorine. Haloids or halogens—Preparation and properties of chlorine— Chlorine water—Hydrochloric acid ; tests for it—Nitro-hydrochloric acid—Compounds of chlorine with oxygen—Hypochlorous acid— Chloric acid; tests for it ......... 116-122 17. Bromine. Iodine. Fluorine. Bromine—Hydrobromic acid —Tests for bromides—Hypobromic and bromic acid—Iodine—Hydriodic acid—Tests for iodine and iodides—Fluorine—Hydrofluoric acid ...... 122-127 PAGE IV. METALS AND THEIR COMBINATIONS. 18. General remarks regarding metals. Derivation of names, symbols, and atomic weights—Melting- points, specific gravities, time of discovery, valence, occurrence in nature, classification, and general properties of metals . . . 128-133 19. Potassium. General remarks regarding the alkali-metals—Occurrence in nature—Potassium hydroxide, carbonate, bicarbonate, nitrate, chlo- rate, sulphate, sulphite, hypophosphite, iodide, bromide—Analytical reactions . 133-140 20. Sodium. Occurrence in nature—Sodium hydroxide, chloride, carbonate, bicarbonate, sulphate, sulphite, thiosulphate, phosphate, nitrate, borate—Analytical reactions—Lithium 140-145 21. Ammonium. General remarks—Ammonium chloride, carbonate, sulphate, ni- trate, phosphate, iodide, bromide, and sulphide—Analytical reactions —Summary of analytical characters of the alkali-metals . . 145-148 22. Magnesium. General remarks—Occurrence in nature—Metallic magnesium— Magnesium carbonate, oxide, sulphate, sulphite—Analytical reactions 148-151 23. Calcium. General remarks regarding alkaline earths—Occurrence in nature —Calcium oxide, hydroxide, carbonate, sulphate, phosphate, acid phosphate, and hypophosphite—Bone-black and bone-ash—Chlorin- ated lime, calcium chloride and bromide—Sulphurated lime—Ana- lytical reactions for calcium—Barium and strontium; their salts and analytical reactions—Summary of analytical characters of the alka- line earth-metals . . . 151-158 X CONTENTS. 24. Aluminum. Occurrence in nature—Metallic aluminum—Alum—Aluminum hydroxide, oxide, sulphate, and chloride—Clay—Glass—Ultrama- rine—Analytical reactions—Cerium—Summary of analytical char- acters of the earth-metals and chromium ..... 158-163 25. Iron. General remarks regarding the metals of the iron group—Occur- rence in nature—Manufacture of iron—Properties—Reduced iron— Ferrous and ferric oxides, hydroxides, and chlorides—Dialyzed iron —Ferrous iodide, bromide, sulphide, and sulphate—Ferric sulphate and nitrate—Ferrous carbonate, phosphate, and hypophosphite— Analytical reactions 164-172 26. Manganese. Chromium. Cobalt. Nickel. Manganese; its oxides and sulphate—Potassium permanganate— Manganese' reactions—Chromium—Potassium dichromate—Chro- mium trioxide—Chromic oxide and hydroxide—Reactions for chro- mium compounds—Cobalt and nickel ...... 173-179 27. Zinc. Occurrence in nature—Metallic zinc—Zinc oxide, chloride, bro- mide, iodide, carbonate, sulphate, and phosphide—Analytical reac- tions—Antidotes—Cadmium—Summary of analytical characters of metals of the iron group 179-183 28. Lead. Copper. Bismuth. General remarks regarding the metals of the lead group—Lead— Lead oxides, nitrate, carbonate, iodide—Poisonous properties of lead —Antidotes—Lead reactions—Copper—Cupric and cuprous oxide— Cupric sulphate and carbonate—Ammonio-copper compounds—Poi- sonous properties and antidotes—Copper reactions—Bismuth—Bis- muthyl nitrate, carbonate, and iodide—Bismuth reactions . . 184-192 29. Silver. Mercury. Silver—Silver nitrate, oxide, iodide—Antidotes—Silver reactions —Mercury—Mercurous and mercuric oxides, chlorides, iodides, sul- phates, nitrates, sulphides—Ammoniated mercury—Antidotes—Mer- cury reactions—Summary of analytical characters of metals of the lead group . . 192-203 30. Arsenic. General remarks regarding the metals of the arsenic group— Arsenic—Arsenous and arsenic oxides and acids—Sodium arsenate —Hydrogen arsenide—Sulphides of arsenic—Arsenous iodide— Analytical reactions—Preparatory treatment of organic matter for arsenic analysis—Antidotes 204-213 31. Antimony. Tin. G-old. Platinum. Molybdenum. Antimony—Trisulphide, oxysulphide, and pentasulphide of anti- mony—Antimonious chloride and oxide—Antidotes—Antimony reactions—Tin—Stannous and stannic chloride—Tin reactions— Gold—Platinum—Molybdenum—Summary of analytical characters of metals of the arsenic group. 213-220 PAGE CONTENTS. XI Y. ANALYTICAL CHEMISTRY. 32. Introductory remarks and preliminary examination. General remarks—Apparatus needed for qualitative analysis— Reagents needed—General mode of proceeding in qualitative analysis — Use of reagents — Preliminary examination — Physical properties—Action on litmus—Heating on platinum foil—Heating on charcoal alone and mixed with sodium carbonate—Flame-tests— Colored borax-beads—Liquefaction of solid substances—Table I.: Preliminary examination 221-231 33. Separation of metals in different groups General remarks—Group reagents—Acidifying the solution— Addition of hydrogen sulphide—Separation of the metals of the arsenic group from those of the lead group—Addition of ammo- nium sulphide and ammonium carbonate—Table IT.: Separation of metals in different groups 232-237 34. Separation of the metals of each group. Table III.: Treatment of the precipitate formed by hydrochloric acid—Treatment of the precipitate formed by hydrogen sulphide— Table IV.: Treatment of that portion of the hydrogen sulphide precipitate which is insoluble in ammonium sulphide—Table V.: Treatment of that portion of hydrogen sulphide precipitate which is soluble in ammonium sulphide—Table VI.: Treatment of the precipitate formed by ammonium hydroxide and sulphide—Table VII.: Treatment of the precipitate formed by ammonium car- bonate—Table VIII.: Detection of the alkalies and of magnesium 237-240 35. Detection of acids. General remarks—Detection of acids by means of the action of strong sulphuric acid—Table IX.: Preliminary examination for acids—Detection of acids by means of reagents added to their neu- tral or acid solution—Table X.: Detection of the more important acids by means of reagents added to the solution—Table XI : Sys- tematically arranged table, showing the solubility and insolubility of inorganic salts and oxides—Table XII.: Table of solubility . 241-247 36. Detection of impurities in officinal inorganic chemical preparations. General remarks—Examination of sulphuric, sulphurous, nitric, phosphoric, and hydrochloric acids—Examination of the com- pounds of potassium, sodium, ammonium, calcium, magnesium, aluminum, iron, zinc, manganese, chromium, lead, copper, bismuth, silver, mercury, arsenic, and antimony ...... 248-255 37. Methods for quantitative determinations. General remarks—Gravimetric methods—Volumetric methods— Standard solutions—Different methods of volumetric determination —Indicators—Titration—Acidimetry and alkalimetry—Normal acid and alkali solution—Oxidimetry—Potassium permanganate and dichromate—Iodimetry—Solutions of iodine, sodium thiosulphate, silver nitrate, and sodium chloride—Gas analysis .... 256-274 PAGE xii CONTENTS. VI. CONSIDERATION OF CARBON COMPOUNDS, OR ORGANIC CHEMISTRY. 38. Introductory remarks. Elementary analysis. Definition of organic chemistry—Elements entering into organic compounds—General properties of organic compounds—Difference in the analysis of organic and inorganic substances—Qualitative analysis of organic substances—Ultimate or elementary analysis— Determination of carbon, hydrogen, oxygen, nitrogen, sulphur, and phosphorus—Determination of atomic composition from results obtained by elementary analysis—Empirical and molecular formulas —Eational, constitutional, structural, or graphic formulas . . 275-283 39. Constitution, decomposition, and classification of organic compounds. Radicals or residues—Chains—Homologous series—Types—Sub- stitution — Derivatives — Isomerism —Metamerism—Polymerism— Various modes of decomposition—Action of heat upon organic sub- stances—Dry or destructive distillation—Action of oxygen upon organic substances—Combustion—Decay—Fermentation and putre- faction — Antiseptics, disinfectants, and deodorizers — Action of chlorine, bromine, nitric acid, alkalies, dehydrating and reducing agents upon organic substances—Classification of organic compounds 284-295 40. Hydrocarbons. Occurrence in nature—Formation of hydrocarbons—Properties— Paraffin or methane series—Methane—Coal, coal-oil, petroleum— Illuminating gas—Coal-tar—Olefines—Benzene series or aromatic hydrocarbons—Volatile or essential oils 296-304 41. Alcohols. Constitution of alcohols—Occurrence in nature—Formation and properties of alcohols—Monatomic normal alcohols—Methyl alcohol —Ethyl alcohol—Alcoholic liquors—Wines, beer, spirits—Amyl alcohol—Glycerin—Nitro-glycerin—Phenols ..... 304-312 42. Aldehydes. Haloid derivatives. Aldehydes—Acetic aldehyde—Paraldehyde—Trichloraldehyde— Chloral hydrate — Chloroform — Bromoform — Iodoform — Ethyl bromide—Sulphonal . 312-319 43. Monobasic fatty acids. General constitution of organic acids—Occurrence in nature— Formation of acids—Properties—Fatty acids—Formic acid—Acetic acid—Vinegar—Reactions for acetates — Acetate of potassium, sodium, zinc, iron, lead, and copper—Acetone—Butyric acid—Vale- rianic acid and its salts Oleic acid ....... 319-328 PAGE xiii CONTENTS. 44. Dibasic and tribasic organic acids. Oxalic acid, oxalates, and analytical reactions—Tartaric acid ; ana- lytical reactions—Potassium tartrate—Potassium-sodium tartrate— Antimony-potassium tartrate—Action of certain organic -acids upon certain metallic oxides—Scale compounds—Citric acid ; analytical reactions—Citrates—Lactic acid—Ferrous lactate .... 329-336 45. Ethers. Constitution—Formation of ethers—Occurrence in nature—Gen- eral properties—Ethyl ether—Acetic ether—Ethyl nitrite—Amyl nitrite—Fats and fat oils—Soap—Lanolin ..... 337-344 46 Carbohydrates. Constitution—Properties—Occurrence in nature—Groups of car- bohydrates—Grape-sugar; tests for it—Fruit-sugar—Inosite—Cane- sugar — Milk-sugar — Starch—Dextrin—Gums—Cellulose — Nitro- cellulose—Glycogen—Glucosides—Digitalin—Myronic acid . . 345-353 47. Amines and amides. Cyanogen compounds Forms of nitrogen in organic compounds—Amines—Amides— Amido-acids—Formation of amines and amides—Occurrence in nature—Cyanogen compounds—Dicyanogen—Hydrocyanic acid— Potassium, silver, and mercuric cyanides—Eeactions for cyanides— Antidotes — Cyanic acid—Sulphocyanic acid — Metallocyanides— Potassium ferrocyanide—Eeactions for ferrocyanides—Potassium ferricyanide—Nitro-cyan-methane 353-362 48. Benzene series. Aromatic compounds General remarks—Benzene series of hydrocarbons—Benzene— Nitro-benzene—Benzene-derivatives—Phenols—Carbolic acid; tests for it—Creasote—Sulphocarbolic acid—Picric acid—Phenacetine— Eesorcin — Cymene — Terpenes — Eesins — India-rubber — Gutta- percha—Stearoptenes —Camphors — Menthol — Thymol — Benzoic acid—Oil of bitter almond—Salicylic acid—Salol—Phthalic acid— Gallic acid —Pyrogallol—Tannin—Naphthalene—Naphthol—San- tonin 362-378 49. Benzene-derivatives containing nitrogen. Aniline—Aniline dyes—Antifebrine—Antipyrine—Saccharine— Pyrrole—Pyridine—Quinoline—Kairine—Thalline . . . 378-383 50. Alkaloids. General remarks—General properties of alkaloids—Mode of ob- taining them—Antidotes—Detection in cases of poisoning—List of important alkaloids—Coniine—Nicotine—Opium — Morphine; its salts and reactions—Codeine—Narcotine and narceine—Meconic acid—Cinchona alkaloids—Quinine; its salts and reactions—Cincho- nine—Cinchonidine—Quinidine—Strychnine and its reactions— Brucine —Atropine — Hyoscyamine—Cocaine and its reactions— Aconitine —V eratrine — Ilydrastine — Berberine —Caffeine — Pto- maines—Leucomaines—Toxalbumins ...... 383-402 PAGE XIV CONTENTS. 51. Albuminous substances or proteids. Occurrence in nature—General properties—Analytical reactions— Classification—Native or true albumins—Globulins—Derived albu- mins or albuminates—Fibrins—Coagulated proteids—Albumoses— Peptones—Amyloid substance—Haemoglobin—Animal cryptolites— Pepsin—Gelatinoids 403-411 PAGE VII. PHYSIOLOGICAL CHEMISTRY. 52. Chemical changes in plants and animals. General remarks—Difference between vegetable and animal life— Formation of organic substances by the plant—Decomposition of vegetable matter in the animal system—Animal food—Nutrition of animals, digestion—Absorption—Respiration—Waste products of animal life—Chemical changes after death 412-421 53. Animal fluids and tissues Constituents of the animal body—Blood—Examination of blood- stains—Chyle—Lymph—Saliva—Gastrice juice ; clinical examina- tion of it—Bile—Biliary pigments—Biliary acids; tests for them— Cholesterin—Lecithin— Panci eatic juice — Feces—Bone—Teeth— Hair, nails, etc.—Mucus—Muscles—Brain 421-435 54. Milk. Properties and composition—Changes in milk—Butter—Cheese— Adulteration of milk—Testing milk—Lactometer, creamometer, lac- toscope—Analysis of milk 435-441 55. Urine and its normal constituents. Secretion of urine—General properties—Composition—Urea; its properties and determination—Uric acid ; tests for it—Hippuric acid ; tests for it 442-449 56. Examination of normal and abnormal urine. Points to be considered in the analysis of urine—Color—Odor— Reaction—Specific gravity—Determination of total solids, and of total organic and inorganic constituents—Detection and estimation of albumin—Blood—Detection and estimation of sugar—Detection of bile—Diazo-reaction—Urinary deposits—Urinary calculi—Micro- scopical examination of urinary sediments 449-473 APPENDIX. Table of weights and measures 475 Table of elements . 477 Index 479 LIST OF ILLUSTRATIONS. FIG- PAGE 1, 2. Structure of matter 22, 23 3. Thermometric scales 27 4. Dialyzer 35 ! 5. Apparatus for the decomposition of mercuric oxide ..... 37 I 6 Apparatus for generating oxygen ........ 75 7. Apparatus for generating hydrogen ........ 79 8. Apparatus for generating ammonia 86 9. Distillation of nitric acid . . . . . . . . . 90 10. Structure of flame 97 ill. Apparatus for making sulphurous acid 102 12. Apparatus for detection of phosphorus Ill 13-16. Detection of arsenic . 209-212 17-21. Apparatus for analytical operations 222, 223 22. Heating of solids in bent glass tube 227 23. Heating on charcoal by means of blowpipe 227 24. Washing and decanting in agate mortar 228 25. Platinum wire for blowpipe experiments ....... 229 26,27. Apparatus for generating hydrosulphuric acid ..... 234 28. Drying-oven 257 29. Desiccator 258 50. Watch-glasses for weighing filters 258 51. Litre flask 259 52. Pipettes 259 >3. Mohr’s burette and clamp 260 >4. Mohr’s burette and holder 60 >5. Gay Lussac’s burette . 261 !6. Titration ............. 265 !7. Flask for dissolving iron . 268 •8. Gas-furnace for organic analysis ......... 279 ■9. Flasks for fractional distillation . 297 0. Liebig’s condenser, with flask 308 1. Apparatus for estimation of urea 446 2. Urinometer 452 3. Nitric acid test for urine 457 4. Urinary sediments 471 COLORED PLATES. FACING PAGE Plate I. Compounds of iron and zinc 1”2 “ II. Compounds of manganese and chromium .... 178 “ III. Compounds of copper, lead, and bismuth • 188 “ IV. Compounds of silver and mercury . . ... 202 “ V. Compounds of arsenic, antimony, and tin ... 208 “ VI. Reactions of alkaloids 384 “ VII. Urine and tests for its constituents 450 ABBREVIATIONS. c. c. = Cubic centimeter. B. P. = Boiling-point. F. P. — Fnsing-point. Sp. gr. — Specific gravity. U. S. P. = United States Pharmacopoeia. PRACTICAL CHEMISTRY, PHARMACEUTICAL AND MEDICAL. I. INTRODUCTION. FUNDAMENTAL PROPERTIES OF MATTER. RESULTS OF THE ATTRACTIONS BETWEEN MASSES, SURFACES, AND MOLECULES. As the science of chemistry has for its object the study of the nature of all substances, or of all varieties of matter, it is necessary first to consider some of the properties which belong to every kind of matter, and are known as essential or fundamental properties. The fundamental properties of matter having a special interest for those studying chemistry are : extension, divisibility, gravitation, porosity, and indestructib ility. 1. EXTENSION OR FIGURE. Matter is anything occupying space, aud this property is known as extension. All bodies, without exception, fill a certain amount of space; they all have length, breadth, aud thickness. That portion of matter lying within the surrounding surface of a body is called its mass. We distinguish three different conditions of matter, namely : Solids, Liquids, and Gases. These conditions of matter are known as the three states of aggregation, and we will now consider the peculiar- ities of matter when existing in either of these states. Solid state. Solids are distinguished by a self-subsistent figure. A solid substance forms for itself, as it were, a casing in which its 18 INTRODUCTION. smallest particles1 are enclosed. The question arises, By what means are these particles connected? how are they kept together? No other answer can be given than that the particles themselves attract each other to such an extent that force is necessary to make them alter their relative position. We see, consequently, that some form of attraction or attractive power is acting between the particles of a solid mass, and we call this kind of attraction cohesion, to distinguish it from other forms of attraction. The external appearance or the figure of solid bodies is various. It may be an irregular or a natural regular figure. Of these two forms, only the latter is here of interest, as it includes all the different crys- tallized substances. Force may be defined as the action of one body upon another body, or as the action of particles of matter upon other particles either of the same or of another body. Strictly speaking, we may say that force is the cause tending to produce, change, or arrest motion ; or it is any action upon matter changing or tending to change its form or position. In many cases force manifests itself as an attractive power; for instance, in the case of cohesion mentioned above, but also in adhesion, gravitation, etc. Forces often give rise to motion (as in the case of heat and electricity), and also to a great variety of changes in matter. The three different states of aggregation are due to the relative intensity of two opposing forces, one—that of molecular attraction—which tends to draw the molecules together, and a second one—that of heat—which tends to separate the molecules from one another. Energy of a body is its capacity of doing work, and is measured by the pro- duct of the force acting and the distance through which it acts. Crystals are solid substances bounded by plane surfaces symmetri- cally arranged according to fixed laws. In explaining the formation of crystals we have to assume that the particles are endowed with the power of attracting one another in certain directions, thereby building themselves up into geometrical forms. The first condition essential to the formation of crystals is the possi- bility of free motion of the smallest particles of the matter to be crys- tallized ; in that case only will they be able to attract each other in such a way as to assume a regular shape, or form crystals. Particles of a solid mass can move freely only after they have been transferred 1 It will be shown later that all matter is supposed to consist of smallest particles, which we call molecules. EXTENSION OR FIGURE. 19 to the liquid or gaseous state. There are two different methods of liquefaction, viz., by means of heat (melting), or solution in some suitable agent (dissolving). In the liquid condition thus produced, the smallest particles can follow their own attraction, and unite to form crystals on removal of the cause of liquefaction (heat or solvent). If two or more (non-isomorphous) substances—for instance, common salt and Glauber’s salt—be dissolved together in water, and the solution be allowed to crystallize, the attraction of like particles for one another will be readily noticed by the formation of distinct crystals of common salt alongside of crystals of Glauber’s salt; neither do the particles of common salt help to build up a crystal of Glauber’s salt, nor the particles of the latter a crystal of common salt. Advantage is taken of this property in separating (by crystal- lization) solids from each other, when they are contained in the same solution. Not all matter can form crystals; some substances never have been obtained in a crystallized state, such as starch, gum, glue, etc. A solid substance showing no crystalline structure whatever, is called amorphous. Some substances capable of crystallization may be obtained also in an amorphous state (carbon, sulphur). Other substances are capable of assuming different crystalline shapes under different conditions. Thus sulphur, when liquefied by heat, assumes, on cooling, a shape different from the sulphur crystallized from a solution. One and the same substance under the same conditions always assumes the same shape. Substances capable of assuming in solidifying two or more different shapes or conditions, are said to be dimorphous and poly- morphous, respectively. When substances of different kinds crystallize in exactly the same form, we call them isomorphous (sulphate of magnesium aud sulphate of zinc). If two isomorphous substauces be contained in one solution, they will crystallize together, and the crys- tals are made up of particles of both substances. Characteristic properties of solids. Solid substances show a great variety of properties caused by the differences in the cohesion of the particles (molecules) composing the substances, and accordingly we distinguish between hard and soft, brittle, tenacious, malleable, and ductile substances. Hardness is that property in virtue of which some bodies resist attempts to force passage between their particles, or which enables solids to resist the dis- placement of their particles. Diamond and quartz are extremely hard, while wax and lead are comparatively soft. INTRODUCTION. Brittleness is that property of solids which causes them to be broken easily when external force is applied to them. Glass, sulphur, coal, etc., are brittle. Tenacity is that property in virtue of which solids resist attempts to pull their particles asunder. Iron is one of the most tenacious substances. Malleability, possessed by some solids, is the property in virtue of which they may be hammered or rolled into sheets. Gold is so malleable that it may be beaten into sheets so thin that it would require about 300,000 laid upon one another to measure one inch. Ductility is the property in virtue of which some solids may be drawn into wire or thin sheets—as, for instance, copper, iron, and platinum. Liquid state. The characteristic features of liquids are, that they have no self-subsistent figure; that they consequently require some vessel to hold them ; and that they present a horizontal surface. While in a solid substance the smallest particles are held together by cohesion to such an extent that they cannot change their relative position with- out force, iu a liquid this cohesion acts with much less energy and permits of a comparatively free motion of the particles; the repellant and attractive forces nearly balance each other in a liquid. That cohesion is not altogether suspended in a liquid is shown by the for- mation of drops or round globules, which, of course, consist of a large number of smallest particles. If there were uo cohesion at all between these particles of a liquid, drops could not be formed. The terms semi-solid and semi-liquid substances are used for bodies occupy- ing a position intermediate between true solids and fluids; butter, asphalt, amorphous sulphur, are instances of this kind. Gaseous state. Matter in the gaseous state has absolutely no self- subsistent figure. Gases fill any vessel or room entirely ; the smallest particles show the highest degree of mobility and move freely in every direction. Cohesion is entirely suspended in gases; indeed, the smallest particles exhibit toward each other an infinite repulsion, so that force is necessary to restrain them within any given bounds what- ever. It, therefore, follows that gases set up and maintain a pressure against the walls of vessels enclosing them. This characteristic pro- perty, possessed by all gases, is known as elasticity, or, better, as ten- sion, and is so unvarying that a law has been established in relation to it. This law is kuown as the Law of Mariotte (though really dis- covered by Boyle, of England, in 1661), and may be expressed thus: The volume of a gas is inversely as the pressure; the density and elastic force are directly as the pressure and inversely as the volume. For instance: If a vessel contains one cubic foot of a gas under a DIVISIBILITY. 21 pressure of ten pounds, the volume will be reduced to one-half, one- tenth, or one-hundredth of one cubic foot, if the pressure be increased to 20, 100, or 1000 pounds respectively. On the contrary, the gas will expand to 2, 10, or 100 cubic feet, if the pressure is reduced to 5, 1, or one-tenth pound respectively. Vapors, produced by evaporation of liquids or solids, have the same properties as gases. 2. DIVISIBILITY. Mechanical comminution. All matter admits of being subdivided into smaller particles, and this property is called divisibility. The processes by which we accomplish the comminution of a solid sub- stance may be of a mechanical nature, such as cutting, crushing, grinding, but beside these modes of subdivision we have other agents or causes by which matter may be divided into smaller particles, and one of these agents is heat. Action of heat on matter. Let us take a piece of ice and convert it, by means of mortar and pestle, into a very fine powder. When the smallest particle of this finely powdered ice is placed under the microscope and heat applied, we shall observe that it becomes liquid, thus proving that it was capable of further subdivision, that it con- sisted of smaller particles, which have now by the action of heat become movable. By further applying heat to the liquid particle of water we may convert it into a gas or vapor, which will escape into the air, or which we may collect in an empty flask. The flask will be filled completely by this water-gas (or steam) obtained by vaporizing that minute particle of ice-dust. This fact demonstrates that me- chanical comminution does not carry us beyond a certain degree of subdivision of matter. That is to say, the smallest fragment of the finest powder still consists of a very large number of much smaller particles. To the smallest particles which compose matter the name molecules has been given. Questions.—1. What is matter and what is force? 2. Mention the prin- cipal fundamental properties of matter. 3. Mention the three states of aggre- gation. 4. Describe the characteristic properties of matter in the solid, liquid, and gaseous states. 5. What is cohesion? 6. Give a definition of a crystal- lized substance. 7. Under what circumstances will matter crystallize? 8. state the difference between amorphous, polymorphous, and isomorphous substances. 9. What is meant by elasticity or tension of gases? 10. State ;he law of Mariotte. 22 INTRODUCTION. Molecular theory. The expression molecule is derived from the Latin word molecula—a little mass, and means the smallest particle of matter that can exist by itself, or into which matter is capable of being subdivided by physical actions. To explain more fully what is meant by the expression molecule, we will return to the conversion of water into steam. Fig. 1. When water boils at the ordinary atmospheric pressure it expands about 1800 times, or one cubic inch of water yields about 1800 cubic inches, equal to about one cubic foot of steam. In explaining this fact we have either to assume that the water, as well as the steam, is homogeneous matter (Fig. 1), or that the water consisted of small particles of a given size, which now exist in the steam again as such, with the only difference that they are more widely separated from each other (Fig. 2). Of the many proofs which we have of the fact that the latter assumption is correct, I will mention but one, viz., that the quantities of vapor formed by volatile liquids at any certain temperature above the boiling-point, in close vessels of the same size, are the same, no matter whether the vessel was entirely empty or contains the vapors of one, two, or more other substances. For instance: If we place one cubic inch of water in a flask holding one cubic foot, from which flask the air has been previously removed, and then heat the flask to the boiling-point, the cubic inch of water will evaporate, filling the vessel with steam. Upon now introducing into the flask a second and a third liquid—for instance, alcohol and ether—we find that of each of these liquids exactly the same quantity will evaporate which would DIVISIBILITY. 23 have evaporated if these liquids had been introduced into the empty flask.1 This fact is evidence that there must be small particles of steam which are not in close contact, that there are spaces between these particles which may be occupied by the particles of a second, third, or more substances. To these particles of matter we give the name molecules, and the spaces between them we call intermolecular spaces. Fig. 2. We have thus demonstrated the correctness, or, at least, the likeli- hood of the so-called molecular theory, but the proof given is but one of many. Of these molecules (though individually by far too small to make any impression whatever upon our senses), our conception is so perfect, that we have formed an idea of the actual size of these minute particles of matter. Very good reasons lead us to believe that the diameter of a molecule is equal to about g o"o o¥o'D~oO' °f one inch? and that one cubic inch of a gas under ordinary conditions contains about one hundred thousand million million millions of molecules. These figures at first glance appear to be beyond the limit of human conception, but in order to give some idea of the size of these mole- cules it may be mentioned that if a mass of water as large as a pea were to be magnified to the size of our earth, each molecule being magnified in the same proportion, these molecules would represent balls of about two inches in diameter. Whilst molecules consequently are exceedingly small particles, yet they are not entirely immeasurable ; they are, as Sir W. Thomson 1 As each gas, in consequence of its tension, exerts a certain pressure, the pressure in the flask rises with the introduction of every additional gas. 24 INTRODUCTION. says, pieces of matter of measurable dimensions, with shape, motion, and laws of action, intelligible subjects of scientific investigation. Before leaving the molecular theory, I will mention the Law (more correctly, the hypothesis) of Avogadro which may be stated as follows: All gases or vapors, without exception, contain, in the same volume, the same number of molecules, provided temperature and pressure are the same. Or, in other words: Equal volumes of different gases contain, under equal circumstances, the same number of molecules. The correct- ness of this law has good mathematical support deduced from the law of Mariotte, many other facts and considerations leading to the same assumption. We shall learn, hereafter, that the law of Avogadro is one of the greatest importance to the science of chemistry. Motion of molecules. Heat. If we place over a gas-flame a vessel containing a lump of ice of the temperature of 0° C., or 32° F., the ice gradually melts and becomes converted into water; but if we measure with a thermometer the temperature of the water at the moment when the last particle of ice is melted, we still find it at the freezing-point, 0° C. or 32° F. From the position of the vessel over the flame, as well as from the fact that the ice has been liquefied, we know that the vessel and its contents have absorbed heat. Yet vessel and water show the same temperature as before. If the heat of the flame is allowed to continue its action on the ice-cold water, the ther- mometer will soon indicate a rapid absorption of heat until it reaches 100° C., or 212° F. Then the water begins to boil and escapes in the form of steam, but the temperature again remains stationary until the last particle of water has disappeared. There must be, consequently, some relation between the state of aggregation of a substance and that agent which we call heat. It was the heat which liquefied the ice, it was the heat which converted the liquid water into steam or gaseous water. Yet the water, having absorbed considerable heat during the process of melting, shows a temperature of 0° C. (32° F.), and the steam also having absorbed large quantities of heat, shows 100° C. (212° F.), the temperature of boiling water. A certain amount of heat has consequently been lost or at least hidden. What has become of it ? According to our present theory, heat is a result of the motion of molecules. All molecules of any substance are in a constant vibratory motion, and the velocity or amplitude of this molecular motion deter- mines the degree of what we call heat. An increase of heat is equal to an increase of the vibratory motion DIVISIBILITY. of the molecules and a decrease in temperature is caused by slower motion. The transfer of heat is a transfer of the motion of some particles to other particles One of the effects of increased heat is in nearly all cases an increase in volume, or, in other words, all substances expand when heated, and contract on cooling. Another effect of the application of heat is, as we have just learned, the conversion of solids into liquids, and of liquids into gases. We noticed also the apparent loss of heat during this conversion, and can easily account for it now by saying, that a certain amount of vibratory motion or a certain velocity of the molecules (more correctly speaking, perhaps, a certain amplitude of molecular motion) is required to convert solids into liquids and liquids into gases. The molecules of steam vibrate with a much greater velocity than those of water of the same temperature, and the molecules of water move with greater velocity than those of ice of the same temperature. In other words, the different states of aggregation depend on the rapidity of the motion of molecules; and the heat which is necessary to convert solids into liquids and liquids into gases, and which is not indicated by the ther- mometer, is called latent heat. This latent heat may again be converted into free heat (heat capable of being indicated by a thermometer), by reconverting the gas into a liquid, or this latter iuto a solid. In both cases a liberation of heat, which is a transfer of the motion of the molecules upon the surround- ings, will be noticed. Increase of volume by heat. The increase of volume by heat is not alike for all matter. Gases expand more than liquids, liquids more than solids, and of the latter the metals more than most other solid substauces. Whilst the expansion of any two or more different solids or liquids is not alike, gases show a fixed regularity in this respect, namely, all gases without exception expand or contract alike, when the temperature is raised or lowered an equal number of degrees. This expansion or contraction of gases is 0.3665 per cent., or of their volume for every degree centigrade; thus 100 volumes of air become 100.8665 volumes when heated 1 degree C., or 136.65 when heated 100 degrees C. This regularity in the expansion and contraction of gases is expressed in the law of Charles, which says: If the pressure remain constant, the volume of a gas increases regularly as the temperature increases, and decreases as the temperature decreases. If heat be applied to a gas confined in a closed vessel and be thus prevented from expanding, the increase of heat will manifest itself as pressure, which 26 INTRODUCTION. rises with the same regularity as shown for expansion, viz., 0.3665 per cent, for every degree centigrade. Melting and boiling. The temperature at which a solid substance is converted into a liquid and this into a gas, is of a certain fixed degree or point for every substance, and the temperatures at which this conversion takes place are known as melting- (fusing-) and boiling-points. Some forms of matter appear incapable of existing in the three states of aggregation, like water. As yet, we know carbon in the solid state only, and the conversion of some gases, as, for instance, oxygen and hydrogen, into liquids or solids, has been accomplished only recently and in very small quantities. Other substances, again, may assume two, but not the third state. Some substances pass from the solid directly into the gaseous state (ammonium chloride, calomel), and the process of converting a solid into a gas directly, and this back again into a solid, is called sublima- tion. Distillation is the conversion of a liquid into a gas, and the recon- densation of the gas into a liquid. Many liquids, and even some solids, evaporate or assume the gaseous state at nearly all temperatures. Water and ice, mercury, camphor, and many other substances vaporize at temperatures which are far below their regular boiling- points. This fact is to be explained by the assumption that during the rapid vibratory motion of the particles of these masses, some particles are driven from the surface beyond the sphere to which the surrounding molecules exert an attraction, and thus intermingle with the molecules of the surrounding air. This evaporation, which takes place at various temperatures and at the surface only, is not to be confounded with boiling, which is the rapid conversion of a liquid into a gas at a fixed temperature with the phenomena of ebullition, due to the formation of gas in the mass of liquid. Boiling-point may there- fore be defined as the highest point to which any liquid can be heated under the normal pressure of one atmosphere. Thermometers are instruments indicating different temperatures. Use is made in their construction of the change in volume of dif- ferent substances by the action of heat. The most common ther- mometer is the mercury thermometer. This instrument may be easily constructed by partly filling with mercury a glass tube having a bulb at the lower end, and placing it into boiling water. The point to which the mercury rises is marked B. P. (boiling-point), and the tube sealed by fusion of the glass. It is then placed in melting ice, and the point to which the mercury sinks is marked F. P. DIVISIBILITY. 27 (freezing-point). The distance between the boiling- and freezing- points is then divided into 100 degrees in the so-called centigrade thermometer, or into 180 degrees in the Fahrenheit thermometer. The inventor of the latter instrument, Fahrenheit, com- menced counting not from the freezing- point, but 32° below it, which causes the freezing-point to be at 32°, the boiling- point at 180° above it, or at 212°. (Fig. 3.) Molecular motion. Heat is but one of the results of molecular motion; other results are light, actinism, electricity and magnet- ism. When a body is heated the molecules vibrate quicker, and this molecular motion gives rise to heat waves in the assumed surrounding and all- pervading medium called ether ; if the heating be continued to a higher degree, the body begins to give out light, which is due to ether waves of shorter length; and if heated yet higher, it gives out not only dark heat waves and light waves, but also waves of still shorter length, which make no direct impression on our senses, but which are capable of producing chemical changes in certain substances, and are known as actinic waves. Of the character of the molecular motion causing electricity and magnetism we know little, and the various theories which have been advanced in order to explain electrical phenomena are inadequate and insufficient to do so satisfactorily. Fig. 3. Centigrade. Fahrenheit, Thermometric scales. Specific heat. Equal weights of different substances require dif- ferent quantities of heat to raise them to the same temperature. For instance : The same quantity of heat which is sufficient to raise one pound of water from 60° to 70°, will raise the temperature of one pound of olive oil from 60° to 80°, or two pounds of olive oil from 60° to 70°. Olive oil consequently requires only one-half of the heat necessary to raise an equal weight of water the same number of degrees. Specific heat is therefore the heat required to raise a certain weight of a substance a certain number of degrees, compared with the heat required to raise an equal weight of water the same number of degrees. 28 INTRODUCTION. The heat required to raise one pound of water one degree centigrade is usually taken as the unit of comparison. On thus comparing olive oil, we find its specific heat to be If we say the specific heat of mercury is we indicate that equal quantities of heat will be required to raise one pound of water or 32 pounds of mercury one degree centi- grade, or that the heat which raises one pound of water one degree will raise one pound of mercury 32 degrees. 3. GRAVITATION. Action of gravitation. Every particle of matter in the universe attracts every other particle; consequently, all masses attract each other, and this attraction is known as gravitation. The action of gravitation between the thousands of heavenly bodies moving in the universe is to be considered by astronomy, but some of the phenomena caused by the mutual attraction of the substances composing the earth are of importance for our present considerations. Such phenomena caused by gravitation are the falling of substances, the flowing of rivers, the resistance which a substance offers on being lifted or carried. A body thrown up into the air or deprived of its support will fall back upon the earth. In this case the mutual attrac- tion between the earth and the substance has caused its fall. It might appear that in this case the attraction was not mutual, but ex- erted by the earth only; it has been proved, however, by most exact experiments, that there is also an attraction of the falling substance for the earth, but the amount or the force of this attraction is directly proportional to the mass of the bodies, and consequently too insignificant in the above case to be noticed. The law of gravitation, known as Newton's law, may thus be stated : All bodies attract each other with a force directly proportional to their masses and inversely proportional to the squares of their distance apart. Questions.—11. What two kinds of divisibility of matter do we distinguish, and by what actions are they accomplished? 12. Explain the term molecule. 13. Mention one of the facts which prove that a gas consists of particles with intervals between them. 14. State the law of Avogadro. 15. Mention the effects produced by increased velocity of the molecules of a mass. 16. Give an explanation of the expressions—latent heat, free heat, and specific heat. 17. Explain the construction of a mercury thermometer. 18. How many degrees of Fahrenheit are equal to 50° C.? 19. How many degrees of centigrade are equal to 167° F.? 20. What is distillation, and what is sublimation? GRAVITATION. 29 Weight is an expression used to denote the quantity of mutual attraction between the earth and the body weighed. Here, again, the attraction of the substance for the earth is not taken into considera- tion. All our weighing is a comparison with, or measurement by, some standard weight, such as pound, ounce, gramme, etc. Specific weight or specific gravity denotes the weight of a body, as compared with the weight of an equal bulk or equal volume of an- other substance, which is taken as a standard or unit. The word density is frequently used for specific weight, as density means comparative mass. By the density of a body consequently is meant its mass (or quantity of matter) compared with the mass of an equal volume of some body arbitrarily chosen as a standard. The standard or unit adopted for all solids and liquids is water at a temperature of 15° C. = 59° F. Specific weight is generally expressed in numbers which denote how many times the weight of an equal bulk of water is contained in the weight of the substance in question. If we say that mercury has a specific gravity or density of 13.6, or that alcohol has a specific gravity of 0.79, we mean that equal volumes of water, mercury, and alcohol represent weights in the proportion of 1, 13.6, and 0.79, or 100, 1360, and 79. The standard or unit chosen for comparing the specific gravity of gases is either atmospheric air or hydrogen. In order to obtain the specific gravity of any liquid, it is only necessary to weigh equal volumes of water and the liquid to be ex- amined, and then to divide the weight of the liquid by the weight of the water. A second method by which the specific gravity of liquids may be determined is by means of the instruments known as hydrometers, or, if made for some special purposes, as alcoholometers, urinometers, alkalimeters, lactometers, etc. Hydrometers are instruments generally made of glass tubes, having a weight at the lower end to maintain them in an upright position in the fluid to be tested as to specific gravity, and a stem above, bearing a scale. The principle upon which their construction depends is the fact that a solid substance when placed in a liquid heavier than itself displaces a volume of this liquid equal to the whole weight of the dis- placing substance. The hydrometer will consequently sink lower in liquids of lower specific gravity than in heavier ones, as the instru- 30 INTRODUCTION. ment has to displace a larger bulk of liquid in the lighter than in the heavier liquid in order to displace its own weight. Weight of gases. We have so far considered the gravity of solids and liquids only, and the next question will be : Do gases also possess weight, are they also attracted by the earth ? The fact that a gas, when generated or liberated, expands in every direction, might indi- cate that the molecules of a gas have no weight, are not attracted by the earth. A few simple experiments will, however, show that gases, like all other substances, have weight. Thus a flask from which the atmospheric air has been removed will weigh less than the same flask wheu filled with atmospheric air or any other gas. Barometer. A second method, by which may be demonstrated the fact that atmospheric air possesses weight, is by means of the barometer. The atmosphere is that ocean of gas which encircles the eartli with a layer some 50 or 100 miles in thickness, exerting a considerable pressure upon all substances by its weight. The instruments used for measuring that pressure are known as barometers, and the most com- mon form of these is the mercury barometer. It may be constructed by tilling with mercury a glass tube closed at one end (and about three feet long) and then inverting it in a vessel containing mercury, when it will be found that the mercury no longer tills the tube to the top, but only to a height of about 30 inches, leaving a vacuum above. The column of mercury is maintained at this height by the pressure of the atmosphere upon the surface of the mercury in the vessel; a column of mercury about 30 inches high must consequently exert a pressure equal to the pressure of a column of the atmosphere of the same diameter as that of the mercury column. As the weight of a column of mercury, having a base of one square inch and a height of about 30 inches, is equal to about 15 pounds, a column of atmosphere having also a base of one square inch must also weigh 15 pounds. In other words, the atmospheric pressure is equal to about 15 pounds to the square inch, or about one ton to the square foot. This enormous pressure is borne without inconvenience by the animal frame in consequence of the perfect uniformity of the pressure in every direction. A barometer may be constructed of other liquids than mercury, but as the height of the column must always bear an inverse proportion to the density of the liquid used, the length of the tube required must be greater for lighter GRAVITATION. 31 liquids. As water is 13.6 times lighter than mercury, the height of a wrater column to balance the atmospheric pressure is 13.6 times 30 inches or about 34 feet, which would therefore be the height of the column of water required. Changes in the atmospheric pressure. The height of the mercury column in a barometer is not the same at all times, but varies within certain limits. These variations are due to a number of causes dis- turbing the density of the atmosphere, and are chiefly atmospheric currents, temperature, and the amount of moisture contained in the atmosphere. As the height and with it the density of the atmosphere diminishes gradually from the level of the sea upward, the height of the mercury column will be lower in localities situated at an elevation. This diminution of pressure is so constant that the barometer is used for estimating elevations. Influence of pressure on state of aggregation. We have seen that the volume of a substance, and, more especially, of a gas, depends upon pressure and temperature, an increase of pressure or decrease of temperature causing the volume to become smaller. We learned also that liquids may be converted into gases, and that this conversion takes place at a certain fixed temperature called the boiling-point. The point, however, changes with the pressure. An increased pressure will raise, a decreased pressure will lower, the boiling-point. Thus water boils at the normal pressure of one atmosphere at 100° C. (212° F.), but it will boil at a lower temperature on mountains in consequence of the diminished atmospheric pressure. If the pressure be increased, as, for in- stance, in steam-boilers, the boiling-point will be raised. Thu3 the boiling- point of water under a pressure of two atmospheres is at 122° C. (251° F.), of five atmospheres at 153° C. (307° F.), often atmospheres at 180° C. (356° F.) Questions.—21. What is gravitation ? 22. Mention some phenomena caused by gravitation. 23. Give a definition of weight. 24. What is specific weight? 25. Name the substances adopted as standards for the determination of specific gravities of solids, liquids, and gases. 26. What is the use made of hydrometers, and on what principle is their construction based ? 27. Explain construction and use of the mercury barometer. 28. Mention some of the causes which have an influence upon the height of the mercury column in the barometer. 29. What is the atmospheric pressure upon a surface of five square feet ? 30. State the relation between boiling-point, temperature, and pressure. 32 INTRODUCTION. 4. POROSITY. Nature of porosity. We have seen that the molecules of any sub- stance are not in absolute contact, but that there are spaces between them which we call intermolecular spaces, aud the property of matter to have spaces between the particles composing it is known as porosity. In the case of solids, these spaces or pores are sometimes of consid- erable size, visible even to the naked eye, as, for instance, in charcoal, whilst in most cases these spaces cannot be discovered, even by the microscope. That even apparently very dense substances are porous, can be demonstrated by the fact that liquids may be pressed through metallic disks of considerable thickness, that gases may be caused to pass through plates of metal or stone, that solids dissolve in liquids without showing a corresponding increase in volume of the solution thus obtained, and, finally, also by the fact that substances suffer expan- sion or contraction in consequence of increased or diminished heat, or in consequence of mechanical pressure. Surface. In every-day life the expression “ surface ” refers to that part of a substance which is open to our senses, visible and measur- able ; but from a more scieutific point of view, we have also to take into consideration those surfaces which, in consequence of porosity, extend to the interior of matter and are invisible to our eyes and absolutely immeasurable by instruments. Surface-action. Attraction acts differently under different condi- tions, aud, accordingly, we assign different names to it. We call it cohesion when it acts between molecules, gravitation when acting between masses, and surface-action or surface-attraction when the attraction is exerted either by the visible surface or by that surface which pervades the whole interior of matter. The phenomena caused by this surface-action are extremely manifold, and some are of suffi- cient interest to be taken into consideration. Adhesion. Most solid substances, when immersed in water, alcohol, or many other liquids, become moist; immersed in mercury, they remain dry. We explain this fact by saying that the surfaces of most solid substances exert an attraction for the particles of such liquids as water and alcohol to such an extent that these particles adhere to the surface of the solids. Such an attraction, however, does not manifest itself PO ROSITY. 33 for the particles of mercury. This form of surface-attraction by which liquids are caused to adhere to solids is called adhesion. This adhesion may be noticed also between two plates having plane surfaces. A drop of water pressed between these plates will cause them to adhere to each other. The application and use of glue and mucilage, our methods of writing and painting, the welding together of pieces of metal, etc., depend on this kind of surface-action. Capillary attraction. Whilst it is the general rule, that liquids in vessels present a horizontal surface, this rule does not hold good near the sides of the vessel. When the liquids wet the vessel, as in the case of water in a glass vessel, the surface is somewhat concave in consequence of the attraction of the glass surface for the particles of water; on the contrary, when the liquids do not wet the vessel, as in the case of mercury in a glass vessel, the surface is somewhat convex. The smaller the diameter of the vessel holding the liquids, the more concave or convex will the surface be. If a narrow tube is placed in a liquid, this surface-action will be more striking, and it will be found that a liquid wetting the tube will not only have a completely concave surface, but the level of the liquid stands perceptibly higher in the tube than the level of the liquid outside. Substances not wetting the tube will show the reverse action, namely, the surface inside of the tube will be convex, and will be below the level of the liquid outside. The attraction of the surface of tubes for liquids, manifesting itself in the concave shape of the surface and in the elevation of the liquid near the tube, is known as capillary attraction. Capillary elevations and depressions depend upon the diameter of the tube, tem- perature, and the nature of the liquid. The narrower the tube, the higher the elevation or the lower the depression; both are diminished by increased temperature. Capillary elevations and depressions, all other circumstances being equal, are inversely proportional to the diameters of the tubes. Defining the phenomena of capillary attraction more scientifically, we may say that the adhesive force of glass, wood, etc., for water and most other liquids exceeds the cohesive force acting between the mole- cules of these liquids, while in mercury the cohesive force predominates over the adhesive. Surface-attraction of solids for gases. Any dry solid substance, carefully weighed, will, after having been exposed to a higher temper- 34 INTRODUCTION. ature, show a decrease in weight whilst yet warm. Upon cooling, the original weight will be restored. This fact cannot be explained otherwise than that some substance or substances must have been expelled by heat, and that this substance or these substances are reabsorbed on cooling. This is actually the case, and the substances expelled and reab- sorbed are the gaseous constituents of the atmospheric air, chiefly the aqueous vapor. Every solid substance upon our earth condenses upon its surface more or less of the gaseous constituents of the atmosphere. This condensation takes place upon the outer as well as upon the inner surface. The amount of gas absorbed depends upon the nature of the gas as well as upon the nature of the absorbing solid. Some of the so-called porous substances, such as charcoal, generally condense or absorb larger quantities than solids of a more dense and compact structure. Heat, as stated above, counteracts this absorbing power. Surface-attraction of solids for liquids or for solids held in solution. When a mixture of different liquids, or a mixture of different solids dissolved in a liquid, is brought in contact with or filtered through a porous solid substance, such as charcoal or bone-black, it will be found that the surface of the solid substance retains a certain amount of the liquids or of the solids held in solution, and that it retains more of one kind than of another. It is this peculiarity of surface-attraction which is made use of in purifying drinking-water by allowing it to pass through charcoal. Bone-black is similarly used for decolorizing sugar-syrup and other liquids. Absorbing power of liquids. In a similar manner as in the case of solids, liquids also exert an attraction for gases. When a gas is con- densed within the pores or upon the surface of a solid, or when it is taken up and condensed by a liquid, we call this process absorption. This absorbing power of different liquids for different gases varies greatly; it is facilitated by low temperature and high pressure, and counteracted by high temperature and removal of pressure. Thus: One volume of water absorbs at ordinary temperature and pressure about 0.03 volume of oxygen, 1 volume of carbon dioxide, 30 vol- umes of sulphur dioxide, and 800 volumes of ammonia. Diffusion. When a cylindrical glass vessel has been partially filled with water, and alcohol, which is specifically lighter than water, is POROSITY. 35 poured upon it, care being taken to prevent a mixing of the two liquids, so as to form two distinct layers, it wil be found that after a certain lapse of time the two liquids have mixed with each other, particles of water having entered the a coliol and particles of alcohol the water, until a uniform mixture of the two liquids has taken place. Upon repeating the experiment with a layer of water over a column of solution of common salt, it will again be found that the two liquids gradually enter one into the other until a uniform salt solution has been formed. In a similar manner, two or more gases introduced into a vessel or a room will readily mix with each other. This gradual passage of one liquid into another, of a dissolved substance into another liquid, or of one gas into another gas, is called diffusion. Osmose. Dialysis. This diffusion takes place also when two liquids are separated by a porous diaphragm, such as bladder or parchment paper, and it is then called osmose or dialysis. The apparatus used for dialysis is called a dialyzer (Fig. 4), and consists usually of a glass cylinder, open at one end and closed at the other by the membrane to be used as the separating medium. This vessel is placed into another, and the two liquids are introduced into the two vessels. If the inner vessel be filled with a salt solution and the outer one with pure water, it will be found that part of the salt solution passes through the mem- brane into the water, whilst at the same time water passes over to the salt solution. On subjecting different sub- stances to this process of dialysis, it has been found that some substances pass through the membrane with much greater facility or in larger quantities than others, and that some do not pass through at all. As a general rule, crystallizable substances pass through more freely than amorphous substances. Those substances which do uot pass through membranes in the process of dialysis are known as col- loids, those which diffuse rapidly crystalloids. Capillary attraction, or, more generally speaking, surface-attraction, is undoubtedly to some extent the cause of the phenomena of osmose, the surface of the diaphragm exercising an attraction upon the liquids. Fig. 4. Dialyzer. 36 INTRODUCTION. Diffusion of gases. A diffusion similar to that of liquids takes place also when two different gases are separated from each other by some porous substance, such as burned clay, gypsum, and others. It has been found that specifically lighter gases diffuse with greater rapidity than the heavier ones. The quantities of two different gases which diffuse into one another in a given time are, as a general rule, inversely as the square roots of their specific gravities. Oxygen is sixteen times as heavy as hydrogen; when the two gases diffuse, it will be found that four times as much hydrogen has penetrated into the oxygen as of the latter gas into the hydrogen. This regularity in the diffusion of gases is expressed in the Law of Graham, thus: The velocity of the diffusion of any gas is inversely proportional to the square root of its density. Indestructibility. All matter is indestructible—i. e., cannot pos- sibly be destroyed by any means whatever, and this property is known as indestructibility. Form, shape, appearance, properties, etc., of matter may be changed in many different ways, but the matter itself can never be annihilated. Apparently, matter often disappears, as, for instance, when water evaporates or oil burns; but these apparent destructions indicate simply a change in the form of matter; in both cases gases are formed, which become invisible constituents of the atmospheric air, and can, therefore, not be seen for the time being, but may be recondensed or rendered visible in various ways. Not only is matter indestructible, energy also partakes of this property. Energy may be converted from one form into some other form. Motion may be converted into heat, and heat into motion, or this motion into electrical energy and chemical energy. In fact, all the different forms of energy are convertible one into the other without loss. This fact is spoken of as the Law of the conservation of energy. To repeat: The total quantity of matter in the universe is con- stant, and the same is true of energy. Questions.—31. What is porosity? 32. What two meanings may be assigned to the word surface? 33. Mention some phenomena caused by surface-action. 34. Explain the term adhesion. 35. Under what circumstances can capillary attraction be noticed, and how does it manifest itself? 36. Give an explana- tion of the word absorption, and mention some instances of the absorption of gases by solids or liquids. 37. What do we understand by diffusion of gases or liquids? 38. Define the word osmose. 39. Which substances are most apt to dialyze, and which have no such tendency ? 40. What is meant by saying that matter and energy are indestructible? II. PRINCIPLES OF CHEMISTRY. RESULTS OF THE ATTRACTION BETWEEN ATOMS. 5. CHEMICAL DIVISIBILITY. Decomposition by heat. The results of the action of heat upon matter have been stated to be : Increased velocity of the motion of molecules, increase in volume of the substance heated, and in many cases a conversion of solids into liquids and of these into gases. Be- sides these results there frequently may be noticed another, now to be mentioned. Fig. 5. Decomposition of mercuric oxide in A; collection of mercury in B, and of oxygen in C. To illustrate this action of heat, we will select the red oxide of mercury, a solid substance which is insoluble in water, almost taste- less, and of a brick-red color. When this oxide of mercury is placed in a glass tube and heated, it will be found to disappear gradually, and we might assume that it has been converted into a gas from which, upon cooling, the red oxide of mercury would be re-obtained. If the apparatus for heating the oxide of mercury be so constructed that the escaping gases may be collected and cooled, we shall not find the red oxide in our receiver, but in its place a colorless gas, whilst at 38 PRINCIPLES OF CHEMISTRY. the same time globules of metallic mercury will be found deposited in the cooler parts of the apparatus (Fig. 5). The action of heat consequently has in this case produced an effect entirely different from the effects spoken of heretofore. There is no doubt that the first action of the heat upon the oxide of mercury is an increased velocity of the motion of its molecules and simulta- neously an increase of its volume, but afterward a decomposition of the oxide takes place, and two substances are liberated, each different from the oxide. One of these substances is a silvery-white, heavy, liquid metal, the mercury ; the other substance is a colorless, odorless gas, which sup- ports combustion much more freely than does atmospheric air, and is known as oxygen. Elements. We have thus succeeded in proving that red oxide of mercury may be converted or decomposed by the mere action of heat into mercury and oxygen. It is but natural to inquire whether it would be possible further to subdivide the mercury or the oxygen into two or more new substances of different properties. To this question, which lias been experimentally propounded to Nature over and over again, we have but one answer, viz.: Oxygen and mercury are substances incapable of decomposition by any method or means as yet known to us. They resist the powerful influences of electricity and heat, even when raised to the highest attainable degrees of in- tensity, and they issue unchanged from every variety of reaction hitherto devised with the view of resolving them into simpler forms of matter. Therefore we are justified in regarding oxygen and mercury as nou- decomposable or simple substances, in contradistinction to compound or decomposable substances, such as the red oxide of mercury. All substances which cannot by any known means be resolved into simpler forms of matter, are called elements; all substances which may, by one process or another, be subdivided or decomposed in such a manner that new substances with new properties are formed, are called compound substances or compounds. While the number of known compounds exceed many thousands, the number of elements is comparatively small, but sixty seven of these simple substances being known to exist on our earth. And yet this small number of elements, by combining with each other in many different proportions, form all that boundless variety of matter which we see in nature. CHEMICAL DIVISIBILITY. 39 Chemical affinity. There must be some cause which enables or even forces the different elements to unite with each other so as to form compound bodies. There must be, for instance, a cause which enables oxygen and mercury to combine and form a red powder. This cause is to be found in the existence of another form of the general attraction which causes the smallest particles of different elements to unite to form new substances with new properties. This kind of attractive power is called chemical force, or chemical affinity, and bodies possessing this capacity of uniting with each other are said to have an affinity for each other. There is a great difference between chemical attraction and the various forms of attraction spoken of heretofore. Cohesion simply holds together the molecules of the same substance, adhesion acts chiefly between the molecules of solid and liquid substances, gravita- tion acts between masses. But all these forces do not change the nature, the external and internal properties of matter; this is done when chemical force or affinity is operating, when a chemical change takes place. For instance : In a piece of yellow sulphur the molecules are held together by cohesion, and we can counteract this cohesion by mechan- ical subdivision, reducing the sulphur to a tine powder; or by the application of heat we can further subdivide the sulphur, melt, and finally volatilize it; or we can throw a piece of sulphur into the air, when it will fall back upon the earth in consequence of gravitation ; or we can dip it into water, when it becomes moist in consequence of surface-action. Yet in all these cases sulphur remains sulphur. It is entirely different when sulphur enters into chemical combina- tion exerting chemical attraction, for instance, when it burns; this means when it combines with the oxygen of the atmospheric air. In this case a new substance, a disagreeably smelling gas, a compound of oxygen and sulphur, is formed. It is consequently a complete change in the properties of matter which follows the action of true chemical attraction ; we might define affinity to be a force by which elements unite and new substances are. generated. Atoms. Molecules, as stated heretofore, are the smallest particles of matter which can exist. All matter consists of molecules, conse- quently the red oxide of mercury must also consist of molecules. By heating the oxide of mercury, oxygen and mercury are obtained, each of which also must consist of molecules. As the oxide of mercury 40 PRINCIPLES OF CHEMISTRY. consists of molecules, and as these molecules are neither pure oxygen nor pure mercury, we must come to the conclusion that a molecule of the oxide of mercury is composed of a small particle of oxygen and a small particle of mercury. We consequently learn that a molecule of a compound substance is composed of yet smaller particles of ele- ments, and these smallest particles of elements capable of entering into combination are called atoms, while molecules are the smallest particles of matter which are capable of existing in a free state. Having now established the difference between atoms and molecules, we may give a better definition of elements and compounds by saying that an elementary substance is one in which the atoms composing its molecules are alike, while in a compound substance the molecules con- tain atoms of different kinds. Chemistry is the science of affinity, and affinity is the attraction act- ing between atoms and causing them to unite and form molecules. As every chemical change is due to the motion of atoms, chemistry may also be defined as the science of the motion of atoms taking place in consequence of chemical affinity. Also, we may say that chemistry is that branch of science which treats of the composition of substances, the changes in their composition, and the laws governing such changes. The scheme below may help to illustrate the relation of chemistry to some other branches of physical science : GENERAL FORCE OF ATTRACTION acting between— Heavenly bodies or masses. Surfaces. Molecules. Atoms. is termed : Gravitation. Surface-action. Adhesion. Capillary attrac- tion, etc. Cohesion. Chemical affinity. The phenomena caused by these respective actions are considered by: Astronomy or Mechanics. Physics. Physics. ■ Crystallography. Chemistry. Atomic weight. All matter possesses weight; this is true of a mass as well as any part of it, and must consequently be true of the atoms also and of the molecules of which matter consists. It is, of course, impossible to weigh a single atom or a single molecule, yet CHEMICAL DIVISIBILITY. 41 science has formed an opinion in regard to the relative weights of these minute particles. The experiment referred to above may be so conducted as to ascertain the weight of the products of decomposition (viz., the oxygen and the mercury) of a given, previously weighed quantity of oxide of mercury. In doing this, it will be found invari- ably that every 13.5 parts by weight of the oxide of mercury yield upon heating 12.5 parts by weight of mercury and 1 part of oxygen, that we have consequently in 13.5 pounds of oxide 12.5 pounds of mer- cury and 1 pound of oxygen. If we assume that a molecule of the oxide is composed of one atom of mercury and one atom of oxygen, we are justified in saying that a mercury atom is 12.5 times heavier than an oxygen atom. In a manner similar to this, the weights of the atoms of all different elements have been compared with each other, and the element having the lightest atom has been selected as the unit of comparison. The element having the lightest atom is hydrogen, and we say the atomic weight of hydrogen is 1, and compare with this weight the weights of all other elements. In doing this, we find that the atom of oxygen weighs sixteen times as much as the atom of hydrogen, and we conse- quently say the atomic weight of oxygen is 16. We have learned before, from the decomposition of the red oxide of mercury, that the mercury atom is 12.5 times as heavy as that of oxygen. As the atomic weight of this element is 16, the atomic weight of mercury must be 12.5 times 16, or 200. Whilst atomic weight is the weight of the atom of any element as compared to the weight of an atom of hydrogen, molecular weight is the combined weight of the atoms forming the molecule. Thus the molec- ular weight of oxide of mercury is 200 + 16 = 216. Chemical symbols. For reasons to be better understood hereafter, chemists designate each element by a symbol, and the first or first two letters of the Latin name of the element have generally been selected. Thus, the symbol of hydrogen is H, of oxygen O, of mercury Hg (from hydrargyrum), of sulphur S, etc. These symbols designate, moreover, not only the elements, but one atom of these elements. For instance: 0 not only signifies oxygen, but one atom or 16 parts by weight of oxygen; and Hg, one atom or 200 parts by weight of mercury. Chemical formulas. In a similar manner as atoms of elements are represented by symbols, the molecules of a compound substance are 42 PRINCIPLES OF CHEMISTRY. designated, and such a representation of a compound substance by symbols is called its formula. Thus, HgO is the formula of the red oxide of mercury, and it tells at once that it is a substance composed of one atom or 200 parts by weight of mercury, and one atom or 16 parts by weight of oxygen. In the molecule of a compound body there must be at least two atoms, each one of a different element, but there may be in a molecule of a compound more than two atoms belonging to two or more ele- ments. For instance : The composition of water is H,0; this means, a molecule of water contains 2 atoms of hydrogen and 1 atom of oxy- gen. When there is more than one atom of an element in a molecule,, the number of these atoms is designated by placing the figure on the right of the symbol and a little below it, as in H20, whilst 2HO or 20H would designate 2 molecules of a substance containing one atom of hydrogen and one atom of oxygen. 6. LAWS OF CHEMICAL COMBINATION. Law of the constancy of composition. This law, also known as the law of definite proportions, was the first ever recognized in chemical science; it was discovered toward the close of the last century, and may be stated thus : A definite compound always contains the same elements in the same proportion ; or, in other words, All chemical com- pounds are definite in their nature and in their composition. To make this law perfectly understood, the difference between a mechanical mixture and a chemical compound must be pointed out. Two powders, for instance sugar and starch, may be mixed together very intimately in a mortar, so that it seems impossible for the eye to discover more than one body. But in looking at this powder with the aid of a microscope, the particles of sugar as well as those of starch may be easily distinguished. The mixture thus produced is a mechani- cal mixture of molecule clusters. Questions.—41. How does heat act upon the red oxide of mercury ? 42. State the difference between mechanical and chemical divisibility. 43. Define the terms element and compound. 44. How many elements and how many com- pound substances are known ? 45. What is chemical affinity, and how does it differ from other forces ? 46. What is an atom, and how does it differ from a molecule ? 47. What is chemistry ? 48. Give a definition of atomic weight and of molecular weight. 49. The atom of which element has been selected as the unit for comparison of atomic weights? 50. Give an explanation of chemical symbols and formulas. LAWS OF CHEMICAL COMBINATION. 43 It is somewhat different when two substances, for instance two metals, are fused together, or when two gases or two liquids (oxygen and nitrogen, water and alcohol) are mixed together, or when finally a solid is dissolved in a liquid (sugar in water). In these instances no separate particles can be discovered even by the microscope. The mix- tures thus produced are mixtures of molecules. Such mixtures always exhibit properties intermediate between those of their constituents and in regular gradation according to the quantity of each one present. The proportions in which substances may be mixed are variable. In a true chemical compound the proportions of the constituent ele- ments admit of no variation whatever ; it is not formed by the mixing of molecules, but by the combination of atoms into molecules: the properties of a compound thus formed usually differ very widely from those of the combining elements. Powdered iron and powdered sulphur tnay be mixed together in many different proportions. If such a mixture be heated until the sulphur becomes liquid, the two elements, iron and sulphur, combine chemically, but they do so in one pro- portion only, 56 parts by weight of iron combining with 32 parts by weight of sulphur to form 88 parts of sulphide of iron. If the two substances had been mixed together in any other proportion than the one mentioned, and which cor- responds to the atomic weights of both elements, the excess of one will be left- undisturbed and uncombined. Law of multiple proportions. If two elements, A and B, are capa- ble of uniting in several proportions, the quantities of B which com- bine with a fixed quantity of A bear a simple ratio to each other. Thus A may combine with B, or A with 2 B, or A with 3 B, etc. This law was discovered at the beginning of the present century, when it was found that the ratio of carbon to hydrogen in olefiant gas, C2II4, is as 6 to 1, in marsh gas, CII4, as 6 to 2, and that the ratio of carbon to oxygen in carbon monoxide, CO, is as 6 to 8, in carbon dioxide, C02, as 6 to 16. These and similar instances led to the discovery of the law of mul- tiple proportions, and it was this law which led Dalton, in 1804, to the adoption of the atomic theory. In thinking and reasoning about this law, he could find no other explanation than that there must be small particles of definite weight which combine with each other, and to these small particles he gave the name atoms. As a very good example illustrating the law of multiple proportions may be mentioned the five compounds formed by the elements nitrogen and oxygen, which compounds have the composition N20, N202, N203, N204, and N205, respectively. In these compounds we find 16, 2X16, 3X16, 4X16, and 5X16 parts by weight of oxygen in combination with 28 parts by weight of nitrogen. 44 PRINCIPLES OF CHEMISTRY. The law of chemical combination by volume, or the law of Gay- Lussac, may be stated as follows: When two or more gaseous con- stituents combine chemically to form a gaseous compound, the volumes of the individual constituents bear a simple relation to the volume of the product. The law may be divided into two laws, thus : 1. Gases combine by volume iu a simple ratio. 2. The resulting volume of the compound, when in the form of a gas, bears a simple ratio to the volumes of the constituents. For instance : 1 volume of hydrogen combines with 1 volume of chlorine, forming 2 volumes of hydro- chloric acid gas; 2 volumes of hydrogen combine with 1 volume of oxygen, forming 2 volumes of water-vapor; 3 volumes of hydrogen combine with 1 volume of nitrogen, forming 2 volumes of ammonia. If the different combining volumes of the gases mentioned are weighed, it will be found that there exists a simple relation between these volumes and the atomic or molecular weights of the elements. For instance: Equal volumes of hydrogen and chlorine combine, and the weights of these volumes are as 1 : 35.4, which numbers rep- resent also the atomic weights of the two elements. Two volumes of hydrogen combine with one volume of oxygen, and the weights of the volumes are as 1 : 8 or 2 : 16, the latter being the atomic weight of oxygen. 1 Volume Hydrogen Weight=l + 1 Volume Chlorine Weight=35.4 2 Vol umes Hydrochl oric Acid gas. W = 36.4 1 Volume Hydrogen W = 1 + 1 Volume Hydrogen W = 1 + 1 Volume Oxygen W = 16 2 Vol umes Waterj-vapor W =; 18 1 Volume Hydrogen W = 1 + 1 Volume Hydrogen W = 1 + 1 Volume Hydrogen W = 1 + 1 Volume Nitrogen W= 14 2 Volumes Amm ionia gas. W =il7 1 Volume Hydrogen W = 1 + 1 Volume Hydrogen W = 1 + 1 Volume Sulphur W = 32 + 1 Volume Oxygen W = 16 + 2 Volumes Sulphuric acid gas. Weight = 98 1 Volume Oxygen W = 16 1 Volume Oxygen W = 16 + + X Volume Oxygen W = 16 LAWS OP CHEMICAL COMBINATION. 45 The above diagram shows the simple relation which exists between combining volumes, atomic and molecular weights; and that such a relation exists is not surprising, if we remember the law of Avogadro, which has been before stated, and which says that all gases under equal conditions contain the same number of molecules. The weighing of equal volumes of gases consequently is identical with the weighing of equal numbers of molecules. The molecular weight of a substance therefore can be found by weighing this sub- stance in the gaseous state and comparing with it the weight of an equal volume of another gas, the molecular weight of which is known. The gas usually adopted for this comparison is hydrogen. If, for instance, we weigh equal volumes of hydrogen, chlorine, oxygen, hydrochloric acid gas, and steam, we find weights in the pro- portion of 2, 70.8, 32, 36.4 and 18. These numbers express also the molecular weights of these substances; moreover they show that atomic and molecular weights of elements are not identical, but that the latter weight is twice that of the atomic weight, or that the molecules of elements consist of two atoms} One litre of hydrogen at the freezing-point of water and under the ordinary pressure of 15 pounds to the square inch, weighs 0.0896 gramme. This weight of one litre of hydrogen is taken as the unit or standard of comparison for gases, and is called one crith. A litre of oxygen weighs 16 criths, one of chlorine 35.4 criths, one of steam 9 criths, etc. Theory (Law) of equivalents. Quantivalence, or Valence. When one element replaces another element in a compound, the quantities of the two elements are said to be equivalent to each other, and according to the law of equivalents the replacement of elements one by another takes place always in definite proportions. Formerly it was believed that all atoms were equivalent among each other, and, accordingly, atomic weights frequently were designated as equivalent weights This view, however, is not correct, as it is found that one atom of one element frequently displaces two or more atoms of another element. This fact, as well as other considerations, has led to the assumption of the quantivalence of atoms. This property will be understood best by selecting for consideration a few compounds of different elements with hydrogen. I. II. III. IY. HC1 H20 H3N h4c HBr h2s H3As H4Si HI H2Se HsP 1 A few exceptions to this general rule will be mentioned in the proper places. 46 PRINCIPLES OF CHEMISTRY. We see here that Cl, Br, and I combine with H in the proportion of atom for atom; O, S, Se combine with H in the proportion of 2 atoms of hydrogen for 1 atom of the other element; N, As, P com- bine with 3; C and Si with 4 atoms of hydrogen. Moreover, it has been found that the compounds mentioned in column I. are the only ones which can be formed by the union of the elements Cl, Br, and I with H. They invariably combine in this proportion only. Other elements show a similar behavior. For instance, the metal sodium combines with chlorine or bromine in one proportion only, forming the compound NaCl or NaBr. Looking at columns II., III., and IV., we see that the elements mentioned there combine with 2, 3, and 4 atoms of hydrogen, respec- tively. It is evident, therefore, that there must be some peculiarity in the power of attraction of different elements toward other elements, and to this property of the atoms of elements of holding in combina- tion one, two, three, four, or more atoms of other elements the name atomicity, quantivalence, or simply valence, has been given. According to this theory of the valence of atoms, we distinguish univalent, bivalent, trivalent, quadrivalent, quinquivalent, sexivalertt, and septivalent elements. All elements which combine with hydrogen in the proportion of one atom to one atom are univalent, as, for instance, Cl, Br, I, F, and all elements which combine with these in but one proportion, that is, atom with atom, bear the same valence, or are also univalent, as, for instance, Na, K, Ag, etc. Those elements which combine with hydrogen or other univalent elements in the proportion of one atom to two atoms are bivalent, such as O, S, Se. Trivalent and quadrivalent elements are those the atoms of which combine with 3 or 4 atoms of hydrogen, respectively. Figuratively speaking, we may say that the atoms of univalent elements have but one, those of bivalent elements two, of trivalent elements three, of quadrivalent elements four bonds or points of attraction, by means of which they may attach themselves to other atoms. Elementary atoms are often named according to their valence: monads, diads, triads, tetrads, pentads, hexads, and heptads. To indicate the valence of the elements frequently dots or numbers are placed above the chemical symbols, thus IP, Ou, NUi, CiUi or Civ. The bonds are often graphically represented by lines, thus : H—, —0—, —N—, —0— I LAWS OF CHEMICAL COMBINATION. 47 It is needless to say that such representations are merely symbolical, and express the view that atoms have a definite power to combine with others. When atoms combine with one another the bonds are said to be satisfied, and it is graphically expressed thus : H IT H | y H—Cl, II—0—II or 0{ , II—N—II or N—II Xh \h Whilst the valence of some elements is invariably the same under all circumstances, other elements show a different valence (this means a different combining power for other atoms) under different condi- tions. For instance: Phosphorus combines both with 3 and 5 atoms of chlorine, forming the compounds PC13 and PC15. As chlorine is a univalent element, we have to assume that phosphorus has in one case 3, in another case 5 points of attraction. Many similar instances are known, and will be spoken of later. 'file only explanation which we can furnish in regard to the variability of the valence of atoms is the assumption that sometimes one or more of the bonds of an atom unite with other bonds of the same atom. If, for instance, in the quinquivalent phosphorus atom two bonds unite with one another a trivalent atom will remain. It is noticed invariably that the valence of atoms increases or diminishes by two, which could not be otherwise, if the explanation given be correct. Thus chlorine, the valence of which generally is I., may also have a valence equal to III., V., or VII., while sulphur shows a valence either of II., IV., or VI. Atoms whose valence is even, as in case of sulphur, are called artiads; those whose valence is expressed in uneven numbers, as chlorine and phosphorus, are called perissads. While it is now being assumed that most of the elements possess more than one valence, in consequence of the assumed power of bonds in the same atom to saturate one another, in this book will be mentioned chiefly that valence which the element seems to possess predominantly. The doctrine of the valence of atoms has modified our views of the equivalence of atoms. We now say that all atoms of a like valence are equivalent to each other. The atoms of each univalent element are equivalent to each other, and so of the atoms of any other valence, but two atoms of a univalent element are equivalent to one atom of a bivalent element, or two atoms of a bivalent element to one atom of a quadrivalent element, etc. After having explained this valence of atoms, it now may be better understood why the atoms in an element do not exist in a free or 48 PRINCIPLES OF CHEMISTRY. uncombined state, but combine with each other to form molecules. Atoms having the tendency of combining with, or attaching them- selves to other atoms, are bound to exert that attraction, and if they are not combined with atoms of other elements, they combine with each other. For instance : Oxygen gas is not a mass of oxygen atoms, but of oxygen molecules, each molecule being formed by the union of two atoms. 7. DETERMINATION OF ATOMIC AND MOLECULAR WEIGHTS.1 Determination of atomic weights by chemical decomposition. The great difficulties originally encountered in the determination of atomic weights cannot well be described here. Consideration will be given alone to the three principal methods at present in use. These methods depend either on chemical action or on physical properties. One of the chemical methods used for the determination of atomic weights has been stated before in describing the decomposition of the red oxide of mercury by heat. The principle of +his method is the determination of the proportions by weight in which the element, the atomic weight of which is unknown, combines with an element the atomic weight of which is known. For instance : If in decomposing a substance we find it to contain in 72 parts by weight 16 parts by weight of oxygen and 56 parts by weight of another element, we have a right to assume the atomic weight of this second element to be 56, provided, however, that the compound is actually formed by the union of one atom of oxygen and one atom of the other element. These 56 parts by weight might, however, represent 2 or 3 or more Questions.—51. State the law of the constancy of composition. 52. What is the difference between a mixture and a chemical compound? 53. Mention some instances of the production of molecular mixtures. 54. State the law of multiple proportions. 55. What considerations led Dalton to the adoption of the atomic theory? 56. What regularity regarding volume is noticed when gases combine chemically? 57. To what was the term equivalent quantities applied formerly, and what is to be understood by it to-day? 58. Explain the term quantivalence or atomicity. 59. Mention some univalent, bivalent, tri- valent, and quadrivalent elements 60. Suppose a certain volume of hydrogen to weigh 20 grains, how much will an equal volume of oxygen and how much an equal volume of hydrochloric acid gas weigh, provided pressure and tem- perature be the same ? 1 The consideration of Chapter 7 should be postponed until the student has become familiar with chemical phenomena generally. DETERMINATION OF ATOMIC WEIGHTS. atoms. If 56 represented 2 atoms, the atomic weight would be but 28 ; if 4 atoms, 14. As this mode of determination gives no clue to the number of atoms present in the molecule, the results obtained are liable to be incorrect. In fact, the atomic weights of a number of elements had originally been determined incorrectly by using the above or similar methods, and many of these old atomic weights had to be changed (generally doubled) in order to obtain the correct numbers. Thus, in examining water, it was found that it contained 8 parts by weight of oxygen to 1 part of hydrogen, and the conclusion was drawn that the atomic weight of oxygen was 8, and that the molecule of water was formed by the union of one atom of hydrogen and one atom of oxygen. It will be demonstrated below why we assume to- day that the atomic weight of oxygen is 16, and that the molecule of water is composed of 2 atoms of hydrogen and 1 of oxygen. Another chemical method of determining atomic weights is the replacement of hydrogen atoms in a known substauce, by the element the atomic weight of- which -is to be determined. For iustance : Hy- drochloric acid is composed of one atom of chlorine weighing 35.4, and one atom of hydrogen weighing 1, the molecular weight of hy- drochloric acid being 36.4. If in this acid the hydrogen be replaced by some other element, for instance by sodium, we are enabled to determine the atomic weight of sodium by weighing its quantity and that of the liberated hydrogen. Suppose that by the action of 36.4 grammes of hydrochloric acid on sodium, 1 gramme of hydrogen was replaced by 23 grammes of sodium. In that case we would say that the atomic weight of sodium is equal to 23. The difficulty which was alluded to above exists also in this mode of determination of atomic weights, viz., not knowing whether it was actually one atom of sodium that replaced the one part of hydrogen, a doubt is left as to whether or not the determination is correct. Determination of atomic weights by means of specific weights of gases or vapors. It has been stated before that equal volumes of gases contain, under like conditions, the same number of molecules (no matter how few or many the atoms within the molecules may be), and that the molecules of elements contain (in most cases) two atoms. These facts give in themselves the necessary data for the determination of atomic weights. For instance: If a certain volume of hydrogen is found to weigh 2 grammes, and an equal volume of some other gaseous element is 50 PRINCIPLES OF CHEMISTRY. found to weigh 71 grammes, then the atomic weight of the latter element must be 35.5, because 2 and 71 represent the relative weights of the molecules of the two elements. Each molecule being composed of 2 atoms, these molecular weights have to be divided by 2 in order to find the atomic weights, which are, consequently, 1 and 35.5 respectively. In comparing by this method oxygen with hydrogen, it is found that equal volumes of these gases weigh 32 and 2 respectively, that the atomic weight of oxygen is consequently 16, and not 8, as deter- mined by chemical methods. This mode of determining atomic weights may be applied to all elements which are gases or which may without decomposition be con- verted into gas. There are, however, elements which cannot be volatilized, and in this case it becomes necessary to determine the spe- cific gravity of some gaseous compound of the element. The element carbon itself has never been volatilized, but we know many of its volatile compounds, and these may be used in the determination of its atomic weight. Determination of atomic weights by specific heat. Specific heat has been stated to be the quantity of heat required to raise the tem- perature of a given weight of any substance a given number of de- grees, as compared with the quantity of heat required to raise the temperature of the same weight of water the same number of degrees. In comparing atomic weights with the numbers expressing the spe- cific heats, it is found that the higher the atomic weight the lower the specific heat, and the lower the atomic weight the higher the specific heat. This simple relation may thus be expressed : Atomic weights are inversely proportional to the specific heats; or the product of the atomic weight multiplied by the specific heat is a constant quantity for the elements examined. Elements. Specific heats. (Water = 1.) Atomic weights. Product of specific heat X atomic wt ight. Lithium, 0.9408 7 6.59 Sodium, 0.2934 23 6.75 Sulphur, 0 2026 32 6.48 Zinc, 0.0956 65 6.22 Bromine (solid),0.0843 80 6.75 Silver, 0.0570 108 6 16 Bismuth, 0.0308 210 6.48 An examination of this table will show this relation between atomic weight and specific heat, and also that the product of atomic weight multiplied by specific heat is equal to about 6.5. The variations DETERMINATION OF ATOMIC WEIGHTS. 51 noticed in this constant quantity of about 6.5 may be due to errors made in the determinations of the specific heats, and subsequent de- terminations may cause a more absolute agreement. However, the agreement is sufficiently close to justify the deduction of a law which says : The atoms of all elements have exactly the same capacity for heat. This law was first recognized by Du long and Petit in 1819, and is simply a generalization of the facts stated. To show more clearly what is meant by saying that all atoms have the same capacity for heat, we will select three elements to illustrate this law. If we take of lithium 7 grammes, of sulphur 32 grammes, of silver 108 grammes, we have of course in these quantities equal numbers of atoms, because 7, 32, and 108 represent the atomic weights of these elements. If we expose these stated quantities of the three elements to the same action of heat, we shall find that the temperature increases equally for all three substances—that is to say, the same heat will be required to raise 7 grammes of lithium 1°, which is necessary to raise either 32 grammes of sulphur or 108 grammes of silver 1°. The quantity of heat necessary to raise the atom of any element a certain number of degrees is, consequently, the same. As heat is the consequence of motion, the result of the facts stated may also be ex- pressed by saying : It requires the same energy to cause different atoms to vibrate with such a velocity as to acquire the same tempera- ture, no matter whether these atoms be light or heavy. It is evident that these facts give us another means of determining atomic weights, by simply dividing 6.5 by the specific heat of the ele- ment. The specific heat of sulphur, for instance, has been found to be 0.2026. 6.5 divided by this number is 31.6, or nearly 32. Originally the atomic weight of sulphur had been determined by chemical methods to be 16, but its specific heat, as well as other properties, has shown this number to be but one-half of the weight, 32, now adopted. It may be mentioned that elements possess essentially the same specific heat whether they exist in a free state or are in combination ; this fact will, in many cases, be of use in the determination of atomic weights. Determination of molecular weights. From the statements made regarding the determination of atomic weights, it is evident that we may use a number of methods for determining molecular weights, these methods being to some extent analogous to the former. Thus we have methods which are based entirely on chemical analysis 52 PRINCIPLES OF CHEMISTRY. or on chemical changes generally. If, for instance, the analysis of a substance shows of calcium 40 per cent., of carbon 12 per cent., and of oxygen 48 per cent., we have a right to assume that the molecule is made up of 1 atom of calcium, 1 atom of carbon, and 3 atoms of oxygen, as the atomic weights of these elements are 40, 12, and 16 respectively. The molecular weight in this case is 100, and the com- position is expressed by the formula CaC03, but the molecular weight might be 200 and the correct formula Ca2C206. There are actually substances which contain such multiples of atoms, as, for instance, the compounds C2H2 and C6H6, and as their percentage composition is identical, analytical methods are insufficient to indicate the number of atoms contained in these molecules. The second method, based on Avogadro’s law, is applicable to all substances which are or can be converted into gases or vapors without decomposition. Weighing equal volumes of hydrogen and of some other substance in the gaseous state gives at once the data for calcu- lating the molecular weight. (See page 44.) A third method, that of Raoult, is based upon the fact that the freezing-point of a liquid is lowered to the same extent by dissolving in it compounds in quantities proportional to their molecular weights. For example: Water begins to solidify at 32° F. (0° C.), but by dis- solving in it say 4 per cent, of its weight of a salt (the molecular weight of which is known) the freezing-point is lowered, say 1° C. If, then, another salt (the molecular weight of which is not known) be dissolved in water, and it be found that to reduce the freezing-point 1° C. there must be dissolved a quantity equal to 7 per cent, of the weight of the water used—then are the molecular weights of the two salts to each other as is 4 to 7. In regard to this method of Raoult it should be stated that it is applicable only to such substances as do not act chemically upon the Questions.—61. What are the three principal methods used for the deter- mination of atomic weights ? 62. Why are chemical means not always sufficient to determine atomic weights? 63. How can the specific gravity of elements in the gaseous state be used for the determination of atomic weight? 64. Describe a method of the determination of atomic weight by chemical means. 65. State one of the reasons why the atomic weight of oxygen has been changed from 8 to 16. 66. What relation exists between atomic weight and specific heat? 67. State the law of Dulong and Petit. 68. Suppose the specific heat of an element to be 0.1138, what will its atomic weight be? 69. Suppose the specific gravity of an elementary gas to be 14, what will its atomic weight be? 70. Sup- pose 216 grammes of an element replace 2 grammes of hydrogen in 73 grammes of HC1, what will the atomic weight of the element be? DECOMPOSITION OF COMPOUNDS. 53 solvent used, and that the ratio of the lowering of the freezing-point is not the same for all substances, but only for members of the same class of substances. 8. DECOMPOSITION OF COMPOUNDS. GROUPS OF COMPOUNDS. Action of heat upon compounds. All phenomena taking place in nature are, without exception, due to motion. Chemistry considers the motion of atoms, without which no chemical change takes place. The causes for chemical changes are either physical actions (heat, electricity), or the decomposing influence of one substance upon another caused by the atoms rearranging themselves into new bodies, so as to better satisfy their affinities. The decomposing action of heat upon compounds has been mentioned before in connection with the decomposition of red oxide of mercury into mercury and oxygen. Similarly to this process, many other compound substances are decomposed by heat either into elements, or, more frequently, into simpler forms of combination. This means that the molecule of a substance containing, for instance, 10 atoms, is split up into 2, 3, or more molecules, each one containing a portion of the 10 atoms. For instance : A piece of marble, which is carbonate of calcium, or CaC03, is decomposed by heat into oxide of calcium, CaO, and carbon dioxide, C02. That heat has such decomposing influence upon compounds is readily accounted for, if we bear in mind, that increase in heat means increased molec- ular vibration, which most likely weakens the stability of the molecule, and diminishes the attractions of its component atoms. On the other hand, heat will in many cases facilitate chemical combination between two substances, because the increased molecular vibration brings the molecules closer within the sphere of each other’s attraction, thereby facilitating chemical union. For instance: Mercury and oxygen do not act upon each other at ordinary tem- perature, but when heated to a temperature somewhat below the boiling-point of mercury, they combine slowly, forming oxide of mercury. This compound, however, as shown before, readily decomposes into mercury and oxygen when heated to a low red heat. The quantity of heat required for decomposition differs widely according to the nature of the substance. Some substances can be produced only at a temperature below the freezing-point of water, a higher temperature causing their decomposition ; other substances may be decomposed at temperatures between the freezing- and boiling- 54 PRINCIPLES OF CHEMISTRY. points; others again, and to these belong the majority of inorganic compounds, may be raised to red or white heat before decomposition sets in; and still another number of compounds have never yet been decomposed by heat. Theoretically, however, we assume that all compounds may be decomposed by heat, should it be possible to raise it to a sufficiently high degree. Decomposition by electricity. Similarly to heat, also electricity decomposes many substances, provided they are in a liquid or gaseous state. These decompositions are usually accomplished by allowing an electric current to pass through the liquid, or electric sparks to pass through the gas. Thus hydrochloric acid, HC1, may be decomposed into hydrogen and chlorine; water, H20, into hydrogen and oxygen. The act of decomposing a compound by electricity is known as electrolysis, and the substance thus decomposed is termed electrolyte. During the decom- position of substances by electrolysis one of the products of decomposition appears at the negative, the other at the positive pole of the battery. Thus, when water is decomposed, the hydrogen is evolved from the negative, the oxygen from the positive pole. Or, when salts are decomposed, the metal is deposited at the negative pole, and the acid or its decomposition products at the positive pole. It was formerly believed that those elements which in electrolysis appear at the negative pole were charged with positive electricity, and were called electro- positive elements, while those appearing at the positive pole were charged with negative electricity and called electro-negative elements. According to this view the non-metals are electro-negative, while the metals are electro-positive. There is a certain relation between electrical and chemical action, as the quantity of electricity which, for instance, sets free 35.4 grammes of chlorine, will also set free 80 grammes of bromine or 127 grammes of iodine. The figures 35.4, 80, and 127 represent the atomic weights of these elements. Decomposition by light. Another cause of decomposition is, in many cases, the action of light. The art of photography is based upon this kind of decomposition. Many substances, easily affected by light, have to be kept in the dark to prevent them from being decomposed. The phenomena of heat, light, and electricity resemble each other in so far as they are phenomena of motion. Heat is the consequence of the motion of material particles (molecules); light is the consequence of the vibratory motion of the hypothetical medium ether; probably the same is true of electricity. These motions, in being transferred to atoms, have, as shown above, fre- quently the tendency of splitting up the molecules of compound substances. DECOMPOSITION OF COMPOUNDS. 55 Mutual action of substances upon each other. As a general rule, it may be said that no chemical action takes place between two sub- stances, both of which are in the solid state, because the molecules do not come in sufficiently close proximity to exchange their atoms. The free motion of the molecules in liquid or gaseous substances facilitates such a proximity, and consequently chemical action. It is often suffi- cient to have but one of the acting substances in the gaseous or liquid state, while the second one is a solid. By converting two solids into extremely fine powder and mixing them together thoroughly, chemi- cal combination may follow, provided the affinity between them be sufficiently strong. The action of substances upon each other may be represented by the following equations, in which the letters stand for elements or groups of elements : 1. A -f B = AB = direct combination. 2 AB -(- C = AC -|- B = direct decomposition. 3. AB + CD = AC -f- BD = double decomposition. As instances illustrating the above, may be mentioned the following chemical reactions : 1. H + Cl = HCl. The formula here given for the formation of hydrochloric acid is not entirely correct, because the action between hydrogen and chlorine does not take place between free atoms, but between the molecules of the two elements, each molecule containing two atoms. The more correct way of writing the formula would therefore be : Hydrogen. Chlorine. Hydrochloric acid. HH + C1C1 = 2HC1. Or 2H + 2C1 = 2HC1. 2. Hydrochloric acid and sodium form sodium chloride and hydro- gen : The formula more correctly written would be : HC1 + Na = NaCl + H. 2HC1 + 2Na = 2HhC1 + 2H.< 3 HCl + AgN03 = AgCl + HNOs. Hydrochloric acid. Silver Nitrate. Silver Chloride. Nitric acid. This form of decomposition, known as double decomposition or metathesis, is one of the common kinds of chemical changes met with in chemical operations. 56 PRINCIPLES OF CHEMISTRY. All the decompositions mentioned above are caused by the affinity which the atoms of one substance have for atoms of another substance. For instance: The decomposition of hydrochloric acid by sodium may be explained by saying that sodium has a greater affinity for chlorine than for hydrogen, as the latter is expelled by the sodium. No general rule can, however, be given for the intensity of affinity with which the atoms of different elements attract each other, because this attraction differs under different conditions. For instance: Water passed in the form of steam over red-hot iron is decomposed, iron oxide and free hydrogen being formed : Fe + H20 = FeO + 211. This decomposition would indicate that the attraction between iron and oxygen is greater than between hydrogen and oxygen. But in passing free hydrogen over heated oxide of iron the reverse action takes place, water and free iron being formed : FeO + 2H — Fe + H2(>. This reaction would indicate that the affinity between oxygen and hydrogen is greater than between oxygen and iron. As a general rule it may be stated that the quantity of a product formed by chemical action of two substances upon one another, is in- fluenced by the relative proportions of the reacting substances. In the above instance iron decomposes water when the iron is in large excess, while a liberal supply of hydrogen causes the reverse action. As a second instance may be mentioned the decomposition of sodium nitrate by sulphuric acid, with the formation of sodium acid sulphate and free nitric acid. On the other hand, sodium acid sulphate is decomposed by a large excess of nitric acid into sodium nitrate and free sulphuric acid. A consideration of this mass-action, as it is now termed, has led to the establishment of the law, that Chemical action is proportional to the active mass of each substance taking part in the change. While the power of affinity possessed by atoms or compounds does not furnish us with data sufficient to predict all chemical changes, we may lay down a general rule which governs the decomposition of cer- tain compounds and which may be stated thus : When two (or more) substances are brought together in solution, which substances by any rearrangement of the atoms may form a product insoluble in the liquid present, this product will form and separate as a precipitate. As instances of this kind of decomposition may .be mentioned the formation of all the hundreds of insoluble metallic salts, which are DECOMPOSITION OF COMPOUNDS. 57 produced by the action of one salt solution upon another salt solution, the first solution containing a metal which with the acid of the second solution may form an insoluble compound, which is then invariably produced as a precipitate. For instance : Calcium carbonate, CaC03, is insoluble; if we bring together two solutions containing a soluble calcium salt and a soluble carbonate, such as calcium chloride, CaCl2, and sodium carbonate, Na2C03, calcium carbonate is precipitated. A second general rule may be stated thus: When two substances capable of forming a volatile product are brought together, the reaction generally takes place. As instances may be mentioned the liberation of carbon dioxide from any carbonate by the action of an acid, and the liberation of ammonia gas from ammonium compounds by calcium hydroxide. The nascent state. This expression is used of elements at the moment when their atoms leave molecules and have not yet had time to reenter into combination. When in this state the atoms have a much greater energy to combine than after having entered into a combination with other atoms of either the same kind (to form elementary molecules) or of another kind (to form compound mole- cules). White arsenic, As2Oa, is a compound of the metal arsenic with oxygen; if through a solution of this compound hydrogen gas be allowed to pass, no chemical change takes place. If, however, hydro- gen be generated or set free in a solution of white arsenic, then the hydrogen atoms, while in the nascent state, have sufficient energy to combine with both th^elements arsenic and oxygen, forming arsenetted hydrogen, AsH3, and water, H20. Chemical reaction in its broader sense refers to any chemical change, but is used more especially when the intention is to study the nature of the substances decomposed or formed. The expression reagent is applied to those substances used for bringing about such changes. Analysis and synthesis. These expressions refer to two methods of research in chemistry, accomplished by two kinds of reaction, analyti- cal and synthetical. Analysis is that mode of research by which compound substances are broken up into their elements or into simpler forms of combina- tion, and analytical reactions are all chemical processes by which the nature of an element, or of a group of elements, may be recognized. 58 PRINCIPLES OF CHEMISTRY. Synthesis is that method of research by which bodies are made to unite to produce substances more complex. Analytical and synthetical methods, or reactions, frequently blend into one another. This means : A reaction made with the intention of recognizing a substance may at the same time produce some com- pound of interest from a synthetical point of view. Acids. The many compounds formed by the union of elements are so various in their nature, that no system of classification proposed up to the present time can be called perfect. There are, however, a few groups or classes of compounds, the properties of which are so well marked, that a substance belonging to either of them may be easily recognized. These groups are the acids, bases, and neutral sub- stances. Acids are characterized by the following properties : 1. They have (when soluble in water) an acid or sour taste. 2. They change the color of many organic substances, for instance of litmus, from blue to red. 3. They contain hydrogen, which can be replaced by metals, the compound thus formed being a salt. According to the number of hydrogen atoms replaceable by metals, we distinguish monobasic, dibasic, and tribasic acids. Hydrochloric acid, HC1, is a monobasic, sulphuric acid, H2S04, is a dibasic, phosphoric acid, H3P()4, is a tribasic acid. Bases or basic substances show properties which are opposed to those of acids. These properties are : 1. They have (when soluble in water) the taste of lye, or an alka- line taste. 2. They have (when soluble in water) an alkaline reaction, i. e.y they restore the color of organic substances when previously changed by acids, for instance that of litmus, from red to blue. 3. When acted upon by acids, they form salts. For instance : Potas- sium hydroxide is a base; when brought in contact with hydrochloric acid it forms water and the salt potassium chloride: KOH + HC1 = H20 -f KC1. Neutral substances. All substances having neither acid nor basic properties are neutral. Water, for instance, is a neutral substance, having no acid or alkaline taste, and no action on red or blue litmus. Many neutral substances, to some extent even water, appear to possess DECOMPOSITION OF COMPOUNDS. 59 the characteristic properties of both classes, acids and bases ; of neither class, however, to a very great extent. Salts. Salts are acids in which hydrogen has been replaced by metals or by basic radicals; a salt may be formed by the union of an acid and a base (usually with the simultaneous formation of water), or by the action of an acid on a metal (usually with the liberation of hydrogen). For instance: NaOH + HN03 = NaNOa + H20 Sodium hydroxide. Nitric acid. Sodium nitrate. Water, Fe + H2S04 = FeS04 + H2 Iron. Sulphuric acid. Ferrous sulphate. Hydrogen. The process of combining an acid with a base in such a proportion that the acid and alkaline reactions disappear, and a neutral salt is formed, is known as neutralization. According to the number of hydrogen atoms replaced in an acid, we distinguish normal and acid salts. A normal salt is one formed by the replacement of all the replaceable hydrogen atoms of an acid. For instance : Potassium chloride, KC1, potassium sulphate, K2S04, potas- sium phosphate, K3P04. (As monobasic acids have but one atom of hydrogen which can be replaced, they form normal salts only.) Normal salts often have a neutral reaction to litmus, but they may have an acid, or even an alkaline reaction. Acid salts are acids in which there has been replaced only a portion of their replaceable hydrogen atoms. For instance : KHS04, K2HP04, kh2po4. Basic salts are salts containing a higher proportion of a base than is necessary for the formation of a normal salt. Instances are basic mercuric sulphate, IIgS04.2Hg0, basic lead nitrate, Pb2N03.Pb20H. According to modern views basic salts are looked upon as derived from bases by neutralization of part of the hydrogen contained in them. In the base lead hydroxide, Pb20H, the dydrogen atoms may be re- placed by the radical of nitric acid when basic lead nitrate, Pb ; qjj3* is formed. In bismuth hydroxide, Bi(OH)3, one, two, or three hydrogen atoms may be replaced by nitric acid, when the salts BiyQjj^2 and Bi(N03)3 are formed. The first two compounds are basic salts, while the third one .is the normal salt. 60 PRINCIPLES OF CHEMISTRY. Double salts are salts formed by replacement of hydrogen in an acid by more than one metal. For instance: Potassium-sodium sulphate, KNaS04. Residue, radical, or compound radical, are expressions for unsatu- rated groups of atoms known to enter as a whole into different com- pounds, but having no separate existence. For instance : The bivalent oxygen combines with two atoms of the univalent hydrogen, forming the saturated compound H20, water. If we take from this H20 one atom of H, there is left the group of atoms HO (now generally written OH), consisting of an atom of oxygen in which but one point of attrac- tion is actually saturated, the second one not being provided for. This group, OH, is a residue or radical, and is known to enter into many compounds; it is, for instance, a constituent of all the different hydroxides (formerly called hydrates), such as potassium hydroxide, KOH, calcium hydroxide, Ca20H, etc. According to the number of points of attraction left unprovided for in a radical, we distinguish univalent, bivalent, trivalent, and quad- rivalent radicals. Carbon is a quadrivalent element forming with the univalent hydro- gen the saturated compound CH4. By removal of one, two, or three hydrogen atoms the radicals CH3', CH2", CH'", are formed. 9. GENERAL REMARKS REGARDING ELEMENTS. Relative importance of different elements. Of the total number of about sixty-nine elements, comparatively but few (about one-fourth) are of great and general importance for the earth, and the phenomena taking place upon it. These important elements form the greater part of the mass of the solid portion of the earth, and of the water and atmosphere, and of all animal and vegetable matter. Another number of elements are of less importance, because either Questions.—71. What physical actions have a tendency to decompose com- pound substances ? 72. Explain the terms reaction and reagent. 73. Mention some instances of decomposition produced by the action of one substance upon another substance. 74. Why can no general rules be established in regard to the amount of attraction which different elements have for each other? 75. What is the difference between analytical and synthetical methods? 76. Define an acid, and state the general properties of basic and neutral substances. By what means can they be recognized? 77. Distinguish between mono-, di-, and tri-basic acids. 78. What are salts and how are they formed? 79. Define neutral, acid, and double salts. 80. Explain the term radical or residue. GENERAL REMARKS REGARDING ELEMENTS. they are not found in any large quantity, or do not take any active or essential part in the formation of organic matter ; yet they are of interest and importance on account of being used, in their elementary state or in the form of different compounds, in every-day life for various purposes. A third number of elements are found in such minute quantities in nature that they are almost exclusively of scientific interest. Even the existence of some elements, the discovery of which has been claimed, is doubtful. The elements enumerated in column I. are those of great and gen- eral interest; in II. those claiming interest on account of the special use made of them ; in III. those having scientific interest only. I. II. III. Aluminium Antimony Beryllium (Glucinum) Calcium Arsenic Caesium Carbon Barium Cerium Chlorine Bismuth Columbium (Niobium) Hydrogen Boron Didymium Iron Bromine Erbium Magnesium Cadmium Gallium Nitrogen Chromium Germanium Oxygen Cobalt Indium Phosphorus Copper Iridium Potassium Fluorine Lanthanum Silicon Gold Osmium Sodium Iodine Palladium Sulphur Lead Rhodium Lithium Rubidium Manganese Ruthenium Mercury Samarium Molybdenum Scandium Nickel Selenium Platinum Tantalum Silver Tellurium Strontium Terbium Tin Thallium Zinc Thorium Titanium Tungsten Uranium Vanadium Ytterbium Yttrium Zirconium Classification of elements may be based upon either physical or chemical properties, or upon a consideration of both. A natural elas- 62 PRINCIPLES OF CHEMISTRY. sification of all elements is the one dividing them into two groups of metals and non-metals. Metals are all elements which have that peculiar lustre known as metallic lustre ; which are good conductors of heat and electricity; which, in combination with oxygen, form compounds generally show- ing basic properties ; and which are capable of replacing hydrogen in acids, thus forming salts. Non-metals or metalloids are all elements not having the above- mentioned properties. Their oxides in combination with water gen- erally have acid properties. In all other respects the chemical and physical properties of non-metals differ widely. Their number amounts to 14, the other 55 elements being metals. Natural groups of elements. Besides classifying all elements into metals and non-metals, certaiu members of both classes exhibit so much resemblance in their properties, that many of them have been arranged into natural groups. The members of such a natural group frequently show some connection between atomic weights and prop- erties. Chlorine, 35.4 Bromine, 80 Iodine, 126.5 Sulphur, 32 Selenium, 78.8 Tellurium, 125 Lithium, 7 Sodium, 23 Potassium, 39 Calcium, 40 Strontium, 87 Barium, 137 Each three elements mentioned in the above four columns resemble each other in many respects, forming a natural group. The relation between the atomic weights will hardly be suspected by looking at the figures, but will be noticed at once by adding together the atomic weights of the first and last elements and dividing this sum by 2, when the atomic weights (very nearly, at least) of the middle mem- bers of the series are obtained. Thus : 35.4 + 126.5 2 32 + 125 = 2 7±i? = 23. 2 40 + 187_R,fi 2 Mendelejeffs periodic law.1 The relationship between atomic weights and properties lias been used for arranging all elements sys- tematically, in such a manner that the existing relation is clearly pointed out. Of the various schemes proposed, the one arranged by Mendelejeff may be selected as most suitable to show this relation. 1 The consideration of this law should be postponed until the student has become acquainted with the larger number of important elements. GENERAL REMARKS REGARDING ELEMENTS. 63 Looking at Mendelejeff’s table on page 64, it will be seen that all the elements are arranged in the order of their atomic weights, and that the latter increase gradually by only a unit or a few units. More- over, the arrangement is such that eight groups and twelve series are formed. The remarkable features of this classification may thus be stated : Elemeuts which are more or less closely allied in their physical and chemical properties are made to stand together in a group, as may be seen by pointing out a few of the more generally known instances as found in the groups I., II., and VII., the first one containing the alkali metals, the second, the metals of the alkaliue earths, the last the halogens. There is, moreover, to be noticed a periodic repetition in the prop- erties of the elements arranged in the horizontal lines from left to right. Leaving out group VIII. for the present, we find that the power of the elements to combine with oxygen atoms increases regu- larly from the left to the right, whilst the power of the elements to combine with hydrogen atoms increases from the right to left, as may be shown by the following instances : I. II. III. IV. V. VI. VII. Na20 MgO A1203 Si02 P205 S03 C1207 Hydrogen compounds unknown SiH4 PH3 SH., C1H The oxides on the left show strongly basic properties, as illustrated by sodium oxide; these basic properties become weaker in the second, and still weaker in the third group; the oxides of the fourth group show either indifferent, or but slightly acid properties, which latter increase gradually in the fifth, sixth, and seventh groups. While some elements show an exception, it may be stated that most of the elements of group I. are univalent, of II. bivalent, of III. trivalent, of IV. quadrivalent, of V. quinquivalent, of VI. sexivalent, and of VII. septivalent. Properties other than those above mentioned might be enumerated in order to show the regular gradation which exists between the members of the various series, but what has been pointed out will suffice to prove that there exists a regular gradation in the properties of the elements belonging to the same series, and that the same change is repeated in the other series, or that the changes in the 'properties of elements are periodic. It is for this reason that a series of elements is called a period (in reality a small period, in order to distinguish it from a large period, an explanation of which term will be given directly). The 12 series or periods given in the following table show another PRINCIPLES OF CHEMISTRY. Group I. Group IT. Group III. Group IV. Group V. Group VI. Group VII Group VIII. Series _ _ K II4 KH3 K II„ K H _ KsO K 0 R,03 ko2 R2O5 E, O3 k.,o7 K 04? 1 H, 1 — — — — — * 2 Li, 7 Be, 9 B, 11 O, 12 N, 14 0, 16 F, 19 3 Na, 23 Mg, 24 Al, 27 Si, 28 P, 31 S, 32 Cl, 35 4 K, 39 (Ja, 40 Sc, 44 Ti, 48 V, 51 Cr, 52 Mn, 54 Fe, 56. Ni, 58. Co, 59 5 (Cu, 63) Zn, 65 Ga, 70 Ge, 72 As, 75 Se, 79 Br, 80 6 Kb, 85 Sr, 87 Y, 89 Zr, 90 Nb, 94 Mo, 96 - Ku, 101. Kb, 103. Pd, 106 7 Ag, 108 Cd, 112 In, 114 Sn, 119 Sb, 120 Te, 125 I, 127 8 Cs, 133 Ba, 137 La, 133 Ce, 140 Di, 142 - - - - - 9 — — — — Er, 166 — — 10 - - Yb, 173 - Ta, 182 W, 184 - Os, 190. Ir, 192. Pt, 194 11 (Au, 196) Hg, 200 Tl, 204 Pb, 206 Bi, 209 — - 12 - — — Th, 2o2 — U, 239 — — — — 1 The decimals are omitted in giving the atomic weights. Periodic System.1 GENERAL REMARKS REGARDING ELEMENTS. 65 highly characteristic feature, which consists in the fact that the corre- sponding members of the even (2, 4, 6, etc.) periods and of the uneven (3, 5, 7, etc.) periods resemble each other more closely than the mem- bers of the even periods resemble those of the uneven periods. Thus the metals calcium, strontium, and barium, of the even periods, 4, 6, and 8, resemble each other more closely than they resemble the metals magnesium, zinc, and cadmium, of the uneven periods, 3, 5, and 7, the latter metals again resembling each other greatly in many respects. It is for this reason that in the table the elements belonging to one group are not placed exactly underneath each other, but are divided into two lines containing the members of even and uneven periods separately, whereby the elements resembling each other most are made to stand together. In arranging the elements by the method indicated, it was found that the elements mentioned in group VIII. could not be placed in any of the 12 small periods, but that they had to be kept separately in a group by themselves, three of these metals always forming an intermediate series following the even periods 4, 6, and 10. An uneven and even series, together with an intermediate series, form a large period, the number of elements contained in a complete, large period being, therefore, 7 4- 7 + 3 = 17. An apparently objectionable feature is the incompleteness of the table, many places being left blank; but it is this very point which renders the table so highly interesting and valuable. Mendelejeff, in arranging his scheme, claimed that the places left blank belonged to elements not yet discovered, and he predicted not only the existence of these as yet missing elements, but also described their properties. Fortunately his predictions have, in at least three cases, been verified, three of the missing elements having since been discovered, am! named, scandium, gallium, and germanium. These elements not only fitted in the previously blank spaces by virtue of their atomic weights, but their general properties also assigned to them the places which they now occupy. Physical properties of elements. Most elements are, at the ordinary temperature, solid substances, two are liquids (bromine and mercury), five are gases (oxygen, hydrogen, nitrogen, chlorine, and fluorine). Most of the solid elements may be converted into liquids and gases by the action of heat. Some solid elements, however, have so far resisted all attempts to change their state of aggregation, as, for instance, carbon. 66 PRINCIPLES OF CHEMISTRY. Most, if not all, of the solid elements may be obtained in the crys- tallized state ; a few are amorphous and crystallized, or polymorphous. The physical properties of many elements in these different, states differ widely. For instance : Carbon is known crystallized as diamond and graphite, or amorphous as charcoal. The property of elements to assume such different conditions is called allotropic modification. Some of the gaseous elements are also capable of existing in allo- tropic modifications. For instance : Oxygen is known as such and as ozone, the latter differing from the common oxygen both in its physi- cal and chemical properties. The explanation given for this surprising fact, that one and the same element has different properties in certain modifications, is, that either the molecule or the atoms within the molecules are arranged differently. Ozone, for instance, has three atoms of oxygen in the molecules, while the common oxygen molecule contains but two atoms. Most of the elements are tasteless and odorless; a few, however, have a distinct odor and taste, as, for instance, iodine and bromine. Relationship between elements and the compounds formed by their union. The properties of the compounds formed by the combination of elements are so various that it is next to impossible to give any general rule by which they may be indicated. It may be said, how- ever, that nearly all of the gaseous compounds contain at least one gaseous element, and that solid elements, when combining with each other, generally form solid substances, rarely liquids, and never compounds showing the gaseous state at the ordinary temperature. Nomenclature. The chemical nomenclature of compound sub- stances has undergone considerable changes within the last twenty years. These changes were made in conformity with our present or modern views of the constitution of the compounds, but many years may yet pass before a uniform system of nomenclature will be adopted generally. When two elements combine in one proportion only, little difficulty is experienced in the formation of a name, as, for instance, in iodide of potassium or potassium iodide, KI, chloride of sodium or sodium chloride, NaCl. When two elements combine in more than one proportion, the syllables, mono, di, tri, tetra, and penta are frequently used to designate the relative quantity of the elements. For instance : Carbon mon- oxide, CO, carbon dioxide, C02, phosphorus trichloride, PC13, phos- phorus pentachloride, PC15. GENERAL REMARKS REGARDING ELEMENTS. 67 In many cases the syllables ous and ic are used to distinguish the proportions in which two elements combine ; the syllable ous being used for the simpler or lower, the syllable ic for the more complex or higher form of combination. For instance : Phosphorous chloride, PC13, and phosphoric chloride, PC15; ferrotts oxide, FeO, ferric oxide, Fe2Os. The syllables mono and sesqui also are used occasionally to mark this difference, as, for instance, monoxide of iron, FeO, sesgm’oxide of iron, Fe203. When two oxides of the same element ending in ous and ic form acids (by entering in combination with water), the same syllables are used to distinguish these acids. Phosphorous oxide, P2Os, forms phosphorotts acid ; phosphoric oxide, P205, forms phosphoric acid. The salts formed by these acids are distinguished by using the syllables ite and ate. Phosphite of sodium is derived from phosphorous acid, phosphate of sodium from phosphoric acid. Sulphites and sulphates are derived from sulphurous and sulphuric acid, respectively. According to the new nomenclature, the name of the metal precedes that of the acid or acid radical in an acid. For instance, sodium phosphite, instead of phosphite of sodium ; potassium sulphate, instead of sulphate of potassium. The acids themselves are looked upon as hydrogen salts, and are sometimes named accordingly : hydrogen nitrate for nitric acid, hydrogen chloride for hydrochloric acid, etc. When the number of elements and the number of atoms increase in the molecule, the names become in most cases more complicated. The rules applied to the formation of such complicated names will be spoken of later. Writing chemical equations. It has been shown that chemical changes are expressed in chemical equations by means of symbols. These equations are formed by placing the molecules which are to act upon one another, and which are called factors and are connected by the sign +, to the left of the sign of equality, and by placing the molecule or molecules which result from the decomposition, and are called product or products, to the right of the sign of equality, con- necting them also by the + sign if more than one product be formed. Every correct chemical equation is correct mathematically also— i. e, the sum of the atoms as well as that of the molecular weights of the factors equals the sum of the atoms and that of the molecular weights of the products respectively. For instance: Sodium car- 68 PRINCIPLES OF CHEMISTRY. bonate and calcium chloride form calcium carbonate and sodium chloride. Expressed in chemical equation we say : JSTa2C03 -f CaCl2 = CaC03 + 2NaCl. Sodium carbonate and calcium chloride are the factors, calcium car- bonate and sodium chloride the products. Adding together the molecular weights of the factors and those of the products we find equal quantities, as follows : 2JSTa = 46 C == 12 30 = 48 Ca = 40 2C1 = 71 Ca = 40 C = 12 30 = 48 2Na — 46 2C1 = 71 Chemical equations not only are used for representing chemical changes, but also are the starting-point in all the chemical calculations in which the quantities of substances entering into chemical actions, or the quantities of the product formed, are concerned. The above calculation teaches, for instance, that 106 parts by weight of sodium carbonate are acted upon by 111 parts by weight of calcium chloride, and that 100 parts by weight of calcium carbonate and 117 parts by weight of sodium chloride are formed by this action. These data may, of course, be utilized to find how much calcium chloride may be needed for the decomposition of one pound or of any other definite weight of sodium carbonate ; or how much of these two substances may be required to produce one hundred pounds, or any other definite weight of calcium carbonate. While in many cases of chemical decomposition the change which is to take place cauuot be foretold, but has to be studied experimentally, there are other chemical changes which can be predicted with certainty (see Chapter 8, page 56). In the latter case especially there is no difficulty in writing out the change in the form of an equation. In doing this it must he borne in mind that equivalent quantities replace one another; that, for instance, two atoms of a univalent clement are required to replace one atom of a bivalent element, as, for instance, in the case of the decomposition taking place between potassium iodide and mercuric chloride, when two molecules of the first are required to decompose one molecule of the second compound : 106 + 111 = 217 100 + 117 = 217 K — I TT /Cl TT /I | K — Cl K —I + Hg\Cl = Hg\I + K—Cl or 2KI + HgCI2 = Hgr2 + 2KC1. GENERAL REMARKS REGARDING ELEMENTS. 69 Whenever the exchange of atoms takes place between univalent and trivalent elements, three of the first are required for one of the second, as in the case of the action of sodium hydroxide on bismuth chloride : Na — OH /Cl /OH Na — Cl Na —OH + Bi—Cl = Bi—OH + Na — Cl Na —OH \C1 \OH Na — Cl or 3NaOH + BiCl3 = Bi(OH)s + 8NaCl. In the following examples of double composition an exchange takes place between the atoms of metallic elements, or between the metallic elements and the hydrogen. The student, in completing the equations, has also to select the correct quantity, i. e., the correct number of molecules of the factors required for the change. The interrogation marks indicate that more than one atom or one molecule of the substance is needed for the reaction. Na' + H/C1 = Cu//S04 + H/S! = H/S04 -(- K/ (?) = Ra//CI2 + Na/S04 = Ca// + B/Cl (?) = Na/C03 + H/S04 = Fe" + H/S04 = Bi"'(NOs)3 + K/OH (?) = H'Cl + Ag'N03 = A1/"(S04)3 + K/OH (?) = Ca"CI2 + Ag'K03(?)z= A1'"(S04)3 + Ca"(OH).(?)= + Ag'NOj (?) = FeawCle + Ag/N03(?) = How to study chemistry. In studying chemistry, the student is advised to impress upon his memory five points regarding every important element or compound. These points are : 1. Occurrence in nature. (Whether in free or combined state; whether in the air, water, or solid part of the earth.) 2. Mode of 'preparation by artificial means. 3. Physical properties. (State of aggregation and influence of heat upon it; color, odor, taste, solubility, etc.) 4. Chemical properties. (Atomic and molecular weight; valence; amount of attraction toward other elements or compounds; acid, alkaline, or neutral reaction ; reactions by which it may be recognized and distinguished from other substances.) 5. Application and use made of it in every-day life, in the arts, manufactures, or medicine. 7 ft Of the most important elements and compounds, the history of their discovery, and, occasionally, some special points of interest, should be noticed also. All students having the facility for working in a chemical labora- tory are strongly advised to make all those experiments and reactions 70 PRINCIPLES OF CHEMISTRY. which will be mentioned in connection with the different substances to be considered in this book. By adopting this mode of studying chemistry the student will soon acquire a fair knowledge of chemical facts, yet he might know little of the science of chemistry. In order to acquire this latter knowledge he should study not only .facts, but also the relationship existing between them and between the laws governing the phenomena con- nected with these facts. It is by this method only that the science of chemistry can be successfully mastered. Questions.—81. Why are not all the elements of equal importance ? 82. State the physical and chemical properties of metals. 83. How are metals distinguished from non-metals ? 84. What relation often exists between the atomic weights of elements belonging to the same group? 85. Explain the term allotropic modification. 86. Mention some elements capable of existing in allotropic modifications. 87. What relation exists between the properties of elements and the properties of the compounds formed by their union? 88- In which cases are the syllables mono-, di-, tri-, tetra-, and penta- used in chemical nomenclature? 89. What use is made of the syllables ous and ic, ite and ate, in distinguishing compounds from each other? 90. What are the principal features of the periodic law ? III. NON-METALS AND THEIR COMBINATIONS. The total number of the non-metals is fourteen; two of them, selenium and tellurium, are of so little importance that they will be but briefly considered in this book. Symbols, atomic weights, and derivation of names. Boron, B = 10.9. From borax, the substance from which boron was first obtained. Bromine, Br = 79.8. From the Greek (3pupog (bromos), stench, in allusion to the intolerable odor. Carbon, C = 12. From the Latin carbo, coal, which is chiefly carbon. Chlorine, Cl = 35.4. From the Greek p6g (chloros), green, in allusion to its green color. Fluorine, F = 19. From fluorspar, the mineral fluoride of calcium, used as flux (fluo, to flow). Hydrogen, H = 1. From the Greek vdup (hudor), water, and jEvvau (gennao), to generate. Iodine, I — 126.5. From the Greek lov (ion), violet, referring to the color of its vapors. Nitrogen, N = 14. From the Greek virpov (nitron), nitre, and yevvau (gen- nao), to generate. Oxygen, O = 16. From the Greek otjvg (oxus), acid, and -yEvvau (gennao), to generate. Phosphorus, P = 31. From the Greek (j>ug (phos), light, and Epsiv (pherein), to bear. Silicon, Si = 28.3. From the Latin silex, flint, or silica, the oxide of silicon, Sulphur, S = 32. From sal, salt, and nvp (pur), fire, referring to the com- bustible properties of sulphur. 72 NON-METALS AND THEIR COMBINATIONS. Under ordinary conditions the non-metals show the following states: State of aggregation. Gases. Liquids. Solids. Hydrogen, Oxygen, Nitrogen, Are converted into liquids with difficulty. Bromine, 63° C. B. P. Phosphorus, Iodine, F. P. 44° C. B. P. 280° C. 175 Sulphur, 107 Ill 400 Chlorine, Fluorine, Easily liquefied. ? Carbon, - Boron, Silicon, - Infusible. Occurrence in nature. a. In a free or combined state. Carbon in coal, organic matter, carbon dioxide, carbonates. Nitrogen in air, ammonia, nitrates, organic matter. Oxygen in air, water, organic matter, most minerals. Sulphur chiefly as sulphates and sulphides. Boron in boric acid and borax. Bromine in salt wells and sea-water as bromide of magnesium, etc Chlorine as chloride of sodium in sea-water, etc. Fluorine as fluoride of calcium, fluorspar. Hydrogen in water and organic matter. Iodine as iodides in sea-water. Phosphorus as phosphate of calcium, iron, etc., in bones. Silicon as silicic acid or silica, and in silicates. b. In combination only. Time of discovery. Sulphur, Carbon, Long known in the elementary state; recognized as elements in the latter part of the eighteenth century. Phosphorus, 1669, by Brandt, of Germany. Chlorine, 1770, by Scheele, of Sweden. . Nitrogen, 1772, by Rutherford, of England. Oxygen, 1774, by Priestley, of England, and Scheele, of Sweden. Hydrogen, 1781, by Cavendish, of England. Boron, 1808, by Gay-Lussac, of France. Fluorine, 1810, by Ampere, of France. Iodine, 1812, by Courtois, of France. Silicon, 1823, by Berzelius, of Sweden. Valence. Univalent. Bivalent. Trivalent or quinquivalent. Quadrivalent. Hydrogen, Oxygen, Nitrogen, Carbon, Chlorine, Sulphur. Boron, Silicon. Bromine, Phosphorus. Iodine, Fluorine. OXYGEN. 10. OXYGEN. OH = 16, (15.96). History. Oxygen was discovered in the year 1774 by Priestley, in England, and Scheele, in Sweden, independently of each other ; its true nature was soon afterward recognized by Lavoisier, of France, who gave it the name oxygen, from the two Greek words, bgvg (oxus), acid, and yewau (gennao), to produce or generate. Oxygen means, conse- quently, generator of acids. Occurrence in nature. There is no other element on our earth present in so large a quantity as oxygen. It has been calculated that not less than about one-third, possibly as much as 45 per cent., of the total weight of our earth is made up of oxygen; it is found in a free or uncombined state in the atmosphere, of which it forms about one- fifth of the weight. Water contains eight-ninths of its weight of oxygen, and most of the rocks and different mineral constituents of our earth contain oxygen in quantities varying from 30 to 50 per cent.; finally, it is found as one of the common constituents of most animal and vegetable matters. If the unknown interior of our earth should be similar in composition to the solid crust of mineral constituents which have been analyzed, then the sub- joined table will give approximately the proportions of those elements present in the largest quantity. Oxygen . . .45 parts. Silicon . . . 28 “ Aluminium . . 8 “ Iron . . . . 6 “ Calcium . . .4 parts. Magnesium . . 2 f‘ Sodium . . . 2 “ Potassium . . 2 “ Preparation. The oxides of the so-called noble metals (gold, silver, mercury, platinum) are by heat into, the metal and oxygen : HgO = Hg + O; Ag20 =2Ag + 0. A more economical method of obtaining oxygen is the decomposi- tion of potassium chlorate, KC103, into potassium chloride, KC1, and oxygen by application of heat: KClOo = KC1 + 8 0. While the above formula represents the final result of the decomposition, it takes place actually in two stages. At first potassium chlorate gives up but one-third of its total oxygen, forming potassium chloride and perchlorate, KC104, thus: 2KC103 = KC104 + KC1 + 2 0. 74 NON-METALS AND THEIR COMBINATIONS. This part of the decomposition takes place at a comparatively low temper- ature; after it is complete, the temperature rises considerably and the decom- position of the perchlorate begins: KC104 = KC1 + 40. If the potassium chlorate be mixed with 30-50 per cent, of man- ganese dioxide, and this mixture be heated, the liberation of oxygen takes place with greater facility and at a lower temperature than by heating potassium chlorate alone. Apparently, the manganese dioxide takes no active part in the decomposition, as its total amount is found in an unaltered condition after all potassium chlorate has been decom- posed by heat. A satisfactory explanation regarding this action of manganese dioxide is yet wanting. A third method is to heat to redness, in an iron vessel, manganese dioxide (Mn02), which suffers then- a partial decomposition : 3Mn02 = Mn304 + 20. In this case there is liberated but one-third of the total amount of oxygen present, while two-thirds remain in combination with the manganese. Other methods of obtaining oxygen are: Decomposition of water by elec- tricity, heating of chromates, nitrates, barium dioxide, sulphuric acid, and other substances, which evolve a portion of the oxygen contained in the molecules. Heating a concentrated solution of bleaching powder with a small quantity of a cobalt salt (cobaltous chloride) furnishes a liberal supply of oxygen, the calcium hypochlorite of the bleaching powder being decomposed into calcium chloride and oxygen: Oxygen may be obtained at the ordinary temperature by adding water to a mixture of powdered potassium ferricyanide and barium dioxide, and also by the decomposition of potassium permanganate and hydrogen dioxide in the presence of dilute sulphuric acid. Ca(C10)2 = CaCl2 + 2 0. Experiment 1. Generate oxygen by heating a small quantity (about 5 grammes) of potassium chlorate in a dry flask of about 100 c.c. capacity, to which, by means of a perforated cork, a bent glass tube has been attached, which leads under the surface of water contained in a dish. (Fig. 6.) Collect the gas by placing over the delivery-tube large test-tubes (or other suitable ves- sels) filled with water. Notice that a strip of wood, a wax candle, or any other substance which burns in air, burns with greater energy in oxygen, and that an extinguished taper, on which a spark yet remains, is rekindled when placed in oxygen gas. Notice, also, the physical properties of the gas. How many c.c. of oxygen can be obtained from 5 grammes of potassium chlorate? 1000 c.c. of oxygen weigh 1.43 grammes. OXYGEN. 75 The quantity of oxygen liberated from a given quantity of a substance may be easily calculated from the atomic and molecular weights of the substance or substances suffering decomposition. For instance: 100 pounds of oxygen may be obtained from bow many pounds of potassium chlorate, or from bow many pounds of manganese dioxide? The molecular weight of potassium chlorate is found by adding together the weights of 1 atom of potassium = 39 + 1 atom of chlorine = 35.4 + 3 atoms of oxygen = 48; total = 122.4. Every 122.4 parts by weight of potassium Fig. 6. chlorate liberate the weight of 3 atoms, or 48 parts by weight of oxygen. If 48 are obtained from 122.4, 100 are obtained from 255. Apparatus for generating oxygen. 48 : 122.4 : : 100 : x x — 255. In a similar manner, it will be found that 813.7 pounds of manganese dioxide are necessary to produce 100 pounds of oxygen. MnO, = 54.8 + 32 = 86.8. 3 Mn02 = 3 X 86.8 = 260.4. Every 260.4 parts furnish 2X16 = 32 parts of oxygen. 32 : 260.4 : : 100 : x ar = 813.7. Physical properties. Oxygen is a colorless, inodorous, tasteless gas; up to a few years ago it was looked upon as a permanent or stable gas, as all attempts to liquefy or solidify it had failed. Lately, however, these efforts have been successful, and oxygen has been converted (though in very small quantities) into a colorless liquid by the applica- tion of a pressure of 470 atmospheres at a temperature of—130° C. (—202° F.). Oxygen is but sparingly soluble in water (about 3 volumes in 100 76 NON-METALS AND THEIR COMBINATIONS. at common temperature). A litre of oxygen under 760 mm. pressure, and at the temperature 0° C. (32° F.), weighs 1.4298 grammes. Chemical properties. The principal feature of oxygen is its great affinity for almost all other elements, both metals and non-metals; with nearly all of which it combines in a direct manner. The more important elements with which oxygen does not combine directly are: Cl, Br, I, F, Au, Ag, and Pt; but even with these it combines indirectly, excepting F. The act of combination between other substances and oxygen is called oxidation, and the products formed, oxides. The large number of oxides are divided usually into three groups, and distinguished as basic oxides (sodium oxide, Na20, calcium oxide, CaO), neutral oxides (water, H20, manganese dioxide, Mn02, lead dioxide, Pb02), and acid-forming or acidic oxides, also called anhydrides (carbon dioxide, C02, sulphur trioxide, S03). Whenever the heat generated by oxida- tion (or by any other chemical action) is sufficient to cause the emis- sion of light, the process is called combustion. Oxygen is the chief supporter of all the ordinary phenomena of combustion. Substances which burn in atmospheric air burn with greater facility in pure oxygen. This property is taken advantage of to recognize and dis- tinguish oxygen from most other gases. Processes of oxidation evolv- ing no light are called slow combustion. An instance of slow combus- tion is the combustion of the different organic substances in the living animal, the oxygen being supplied by respiration. For a process of oxidation it is not absolutely necessary that free oxygen be present. Many substances contain oxygen in such a form of combination that they part with it easily when brought in contact with substances having a greater affinity for it. Such substances are called oxidizing agents, as, for instance, nitric acid, potassium chlorate, potassium permanganate, etc. In all combustions we have at least two substances acting chemically upon one another, which substances are generally spoken of as combustible bodies and supporters of combustion. Illuminating gas is a combustible substance, and oxygen a supporter of combustion; but these terms are only relatively correct, since oxygen may be caused to burn in illuminating gas, whereby it is made to assume the position of a combustible substance, whilst illuminating gas is the supporter of combustion. While some substances, such as iron and phosphorus, undergo slow combus- tion at the ordinary temperature, there is a certain degree of temperature, characteristic of each substance, at which it inflames. This point is known as kindling temperature, and varies widely in different substances. Zinc ethyl ignites at the ordinary temperature, phosphorus at 50° C. (122° F.), sulphur at OXYGEN. 77 about 450° C. (842° F.), carbon at a red heat, and iron at a white heat. The heat produced by the combustion is generally higher than the kindling tem- perature, and it is for this reason that a substance continues to burn until it is consumed, provided the supply of oxygen be not cut off, and the temperature be not through some cause lowered below the kindling temperature. The total amount of heat evolved during the combustion of a substance is the same as that generated by the same substance when undergoing slow com- bustion, but the intensity depends upon the time required for the oxidation. A piece of iron may require years to combine with oxygen, and it may be burned up in a few minutes; yet the total heat generated in both cases is the same, though we can notice and measure it in the first instance by most deli- cate instruments only, while in the second it is very intense. Ozone is an allotropic modification of oxygen, which is formed when non-lnminous electric discharges pass through atmospheric air or through oxygen ; when phosphorus, partially covered with water, is exposed to air, and also during a number of chemical decomposi- tions. Ozone differs from ordinary oxygen by possessing a peculiar odor, by being an even stronger oxidizing agent than common oxygen, by liberating iodine from potassium iodide, etc. This latter action may be used for demonstrating the presence of ozone by suspending in the gas a paper moistened with a solution of potassium iodide and starch. The iodine, liberated by the ozone, forms with starch a dark- blue compound. Theoretically, we assume that ozone contains three, common oxygen but two, atoms in the molecule, which is substan- tiated by the fact that three volumes suffer a condensation to two vol- umes when converted into ozone, which would indicate that three molecules of oxygen furnish two molecules of ozone, thus : or 302 = 20;j °=° A 0=0 = °~° °=° 6—b Ozone occurs in small quantities in country air, but is rarely noticed in cities, where it is decomposed too quickly by the impurities of the atmospheric air. It has been assumed that ozone acts advantageously, as it has a tendency to destroy matters which are unwholesome. Too little, however, is known of the subject to justify a positive opinion in regard to it. Questions.—91. By whom and at what time was oxygen discovered? 92. How is oxygen found in nature ? 93. Mention three processes by which -oxygen may be obtained. 94. How much oxygen may be obtained from 490 78 NON-METALS AND THEIR COMBINATIONS. 11. HYDROGEN. II = 1. History. Hydrogen was obtained by Paracelsus in the 16th cen- tury; its elementary nature was recognized by Cavendish, in 1781- The name is derived from i>6II20. Potassium chlorate crystallizes in plates of a pearly lustre; it is soluble in 17 parts of cold, and 2 parts of boiling water. It is even a stronger oxidizing agent than potassium nitrate, for which reason care must be taken in mixing it with organic matter or other deoxi- dizing agents, or with strong acids, which will liberate chloric acid. When heated by itself, it is decomposed into potassium chloride and oxygen. Potassium sulphate, Potassii sulphas, K2S04 = 174. Obtained by the decomposition of potassium chloride, nitrate, or carbonate, by sulphuric acid : 2KC1 + H2S04 = 2HC1 + K2S04; K2C03 + H2S04 = H20 + co2 + K2S04. Potassium sulphate exists in small quantities in plants, and in nearly all animal tissues and fluids, nlore abundantly in urine. Potassium hydrogen sulphate, bisulphate, or potassium acid sulphate, may be obtained by the action of one molecule of potassium chloride upon one molecule of sulphuric acid : KC1 + II2S04 = HC1 -I- KIISO,4. Potassium sulphite, Potassii sulphis, K2S032H20 = 194. Obtained by the decomposition of potassium carbonate by sulphurous acid : K2C03 + H2S03 = H20 + C02 + k2so3. Potassa Sulphurata, U. S. P. {Sulphurated, potassa, Sulpliuret of potash, Hepar mlphuris). A mixture of potassium sulphide, polysulphide, and thiosulphate. It is made by heating a mixture of one part of sulphur and two parts of potas- 138 METALS AND THEIR COMBINATIONS. sium carbonate in a covered crucible, and pouring the fused mass on a marble slab: 3K2C08 + 8S = K2S203 + 2K2S3 + 3C02. The freshly prepared substance has a liver-brown color, turning gradually to greenish yellow; it is very apt to absorb water and oxygen, both the sulphide and hyposulphite becoming oxidized, and finally converted into sulphates. Potassium hypophosphite, Potassii hypophosphis, KPH202 = 104, may be obtained by decomposing a solution of calcium hypophos- phite by potassium carbonate: Ca2PH202 + K2C03 = 2KPH202 + CaC03. The filtered solution is evaporated at a very gentle heat, stirring constantly from the time it begins to thicken until a dry, granular salt is obtained, which is soluble in 0.6 part of cold and 0.3 part of boiliug water. Potassium iodide, Potassii iodidum, KI = 165.5 (Iodide of potas- sium), is made by the addition of iodine to a solution of potassium hydroxide until the dark-brown color no longer disappears : Iodide and iodate of potassium are formed, and may be separated by crystallization. A better method, however, is to boil to dryness the liquid containing both salts, and to heat the mass after having mixed it with some charcoal, in a crucible, when the iodate is con- verted into iodide : 6K0H + 61 = 5KI + KI03 + 3H20. KI03 + 3C = KI + 3C0. Experiment 17. Add to a solution of about 3 grammes of potassium hydroxide in about 25 c.c. of water (or to the solution obtained by making Experiment 16) iodine until the brown color no longer disappears. (How much iodine will be needed for 3 grammes of KOH?) Evaporate the resulting solution (What does this solution contain now?) to dryness, mix the powdered mass with about 10 per cent, of powdered charcoal and heat the mixture in a crucible until slight deflagration has taken place. Dissolve the fluid mass in hot water, filter and set aside for crystallization ; if too much water has been used for dissolving, the liquid must be concentrated by evaporation. Potassium iodide forms colorless, cubical crystals, which are soluble in 0.5 part of boiling and 0.8 part of cold water, also soluble in 18 parts of alcohol, and in 2.5 parts of glycerin. When heated it fuses, and at a bright-red heat is volatilized without decomposition. Potassium bromide, Potassii bromidum, KBr = 118.8 (Bromide of potassium), may be obtained in a manner analogous to that given for POTASSIUM. 139 potassium iodide, by the action of bromine upon potassium hydrox- ide, etc. Or it may be made by the decomposition of a solution of ferrous bromide by potassium carbonate : FeBr2 + K2C03 = 2KBr + FeC03. Ferrous carbonate is precipitated, whilst potassium bromide remains in solution, from which it is obtained by crystallization. Potassium salts of interest, which have not yet been mentioned, will be con- sidered under the head of their respective acids. Some of these salts are potas- sium chromate and permanganate, and the salts formed from organic acids, such as potassium tartrate, acetate, etc. Analytical reactions. (Potassium chloride, KC1, or nitrate, KN03, may be used.) 1. To a solution of potassium chloride, or to any salt of potas- sium, after a few drops of hydrochloric acid have been mixed with it, add platinic chloride and some alcohol : a yellow crystalline pre- cipitate falls, which is a double chloride of platinum and potassium, PtCl42KCl. 2KC1 + PtCl4 = PtCl42KCl; 2KN03 + 2HC1 + PtCI4 = PtCl42KCl + 2HN03. The last formula shows the necessity of adding hydrochloric acid, which is not required in case potassium chloride is used. The ad- dition of alcohol facilitates the precipitation of the double chloride of potassium and platinum, because it is less soluble in alcohol than in water. 2. To a neutral or slightly acid solution of a potassium salt add sodium cobaltic nitrite : a yellow precipitate of potassium cobaltic nitrite, (KN02)6.Co2(N02)6 T H20, is produced. (The reaction is not influenced by the presence of alkaline earths, earths, or metals of the iron-group.) 3. Add to a concentrated solution of a neutral potassium salt a freshly prepared strong solution of tartaric acid : a white precipitate of potassium acid tartrate, KHC4H406, is slowly formed. Addition of alcohol facilitates precipitation. 4. Potassium compounds color violet the flame of a Bunsen burner or of alcohol. The presence of sodium, which colors the flame in- tensely yellow, interferes with this test, as it masks the violet caused by potassium. The difficulty may be overcome by observing the flame through a blue glass or through a thin vessel filled with a solu- 140 METALS AND THEIR COMBINATIONS. tiou of indigo. The yellow light is absorbed by the blue medium, while the violet light passes through and can be recognized. 5. All compounds of potassium are white (unless the acid has a coloring effect), soluble in water, and not volatile at a low red heat. 20. SODIUM (NATRIUM). Nai = 23. Occurrence in nature. Sodium is found very widely diffused in small quantities through all soils. It occurs in large quantities in combination with chlorine, as rock-salt, or common salt, which forms considerable deposits in some regions, or is dissolved in spring waters, and is by them carried to the rivers, and finally to the ocean, which contains immense quantities of sodium chloride. It is found, also, as nitrate, silicate, etc. Sodium hydroxide, Sodium hydrate, Soda, Na0H = 40, may be obtained by the processes mentioned for potassium hydroxide. Sodium chloride, Sodii chloridum, NaCl = 58.4 {Chloride of sodium, Common salt). This is the most important of all sodium compounds, and also is the material from which the other compounds are directly or indirectly obtained. Common table-salt frequently coutaius small quantities of calcium and magnesium chlorides, the presence of which causes absorption of moisture, as these compounds are hygroscopic, whilst pure sodium chloride is not. In the animal system, sodium chloride is found in all parts, it being of great importance in aiding the absorption of albuminoid substances and the phenomena of osmose; also by furnishing, through decompo- sition, the hydrochloric acid of the gastric juice. Questions.—181. How is potassium found in nature, and from what sources is the chief supply of potassium salts obtained? 182. What color have the salts of the alkali metals, and which are insoluble? 183. Mention two pro- cesses for making potassium hydroxide, and what are its properties? 184. Show by symbols the conversion of carbonate into bicarbonate of potassium. 185. Explain the principle of the manufacture of potassium nitrate, and what is the office of the latter in gunpowder ? 186. How is potassium chlorate made, and what are its properties ? 187. Give the processes for manufacturing iodide and bromide of potassium, both in words and symbols. 188. State the com- position of potassium sulphate and sulphite, flow can they be obtained? 189. What is sulphuret of potash? 190. Mention tests for potassium com- pounds. SODIU M. 141 Sodium chloride is soluble in about 2.8 parts of water, at all tem- peratures ; it crystallizes in cubes. Sodium carbonate, Sodii carbonas, Na2C03.10H20 = 286 (Carbonate of sodium, Washing soda, Sal sodee). This compound is, of all alka- line substances, the one manufactured in the largest quantities, being used in the fabrication of many highly important articles, as, for in- stance, soap, glass, etc. Sodium carbonate is made, according to Leblanc’s process, from the chloride by first converting it into sulphate (salt-cake) by the action of sulphuric acid : The escaping vapors of hydrochloric acid are absorbed in water, and this liquid aeid is used largely in the manufacture of bleaching-pow- der. The sodium sulphate is mixed with coal and limestone (calcium carbonate) and the mixture heated in furnaces, when decomposition takes place, calcium sulphide, sodium carbonate, and carbonic oxide being formed : 2NaCl + H2S04 = 2HC1 + Na2S04. Na2S04 + 40 + CaC03 = CaS -f Na2C03 + 4C0. The resulting mass, known as black-ash, is washed with water, which dissolves the sodium carbonate, whilst calcium sulphide enters into combination with calcium oxide, thus forming an insoluble double compound of oxy-sulphide of calcium. The liquid obtained by washing the black-ash, when evaporated to dryness, yields crude sodium carbonate, or u soda ash'’ ; when this is dissolved and crystallized it takes up ten molecules of water, forming the ordinary “soda.” Sodium carbonate is manufactured also by the so-called ammonia process, or the Solvay process. This depends on the decomposition of sodium chloride by ammonium bicarbonate under pressure, when sodium bicarbonate and ammonium chloride are formed, thus : NaCl + NH4HC03 = NH4C1 + NaHC03. The sodium acid carbonate, thus obtained, is converted into carbo- nate by heating: 2NaHC03 = Na2C03 + II20 + C02. The carbon dioxide obtained by this action is caused to act upon ammonia, liberated from ?he ammonium chloride, obtained as one of the products in the first reaction. Ammonium carbonate is thus regenerated and used in a subsequent operation for the decomposition of common salt. 142 METALS AND THEIR COMBINATIONS. Sodium carbonate has strong alkaline properties; it is soluble in 1.6 parts of water at ordinary temperature, and in much less water at higher temperatures; the crystals lose water on exposure to the air, falling into a white powder ; heat facilitates the expulsion of the water of crystallization, and is applied in making the dried sodium carbon- ate, Sodii carbonas exsiccatus of the U. S. P. Sodium bicarbonate, Sodii bicarbonas, NaHC03 = 84 {Bicarbonate of sodium). Obtained, as stated in the previous paragraph, by the am- monia-soda process. It can also be made by passing carbon dioxide over sodium carbonate from which the larger portion of water of crys- tallization has been expelled: Na2C02 -f H20 + C02 = 2NhHC03. It is a white powder, having a cooling, mildly saline taste, and a slightly alkaline reaction. Soluble in 12 parts of cold water, and in- soluble in alcohol. It is decomposed by heat or by hot water into sodium carbonate, water, and carbon dioxide. Sodium sulphate, Sodii sulphas, Na2S0410H20 = 322 {Sulphate of sodium, Glauber’s salt). Made, as mentioned above, by the action of sulphuric acid on sodium chloride, dissolving the salt thus obtained in water, and crystallizing. Large, colorless, transparent crystals, rapidly efflorescing on exposure to air. Soluble in 2.8 parts of water at 15° C. (59° F.), in 0.25 part at 33° C. (91° F.), and in 0.47 part of boiling water. Experiment 18. Dissolve about 10 grammes of crystallized sodium carbonate in 10 c.c. of hot water, add to this solution dilute sulphuric acid until all effer- vescence ceases, and the reaction on litmus-paper is exactly neutral. Evaporate to about 20 c,c., and set aside for crystallization. Explain the action taking place, and state how much H2S04, and how much of the diluted sulphuric acid, U. S. P., are needed for the decomposition of 10 grammes of crystallized sodium carbonate. Sodium sulphite, Sodii sulphis, Na2S03.7H20 = 252. Sodium bisul- phite, Sodii bisulphis, NaHS03 = 104. By saturating a cold solution of sodium carbonate with sulphur dioxide, sodium bisulphite is formed, and separates in opaque crystals : Na2C03 + 2S02 + II20 = 2NaHS03 + C02. If to tlie sodium bisulphite thus obtained a quantity of sodium car- bonate be added, equal to that first employed, the normal salt is formed : 2NaHS0, + Na2COs = 2Na2SOs + H20 + C02. SODIUM. 143 Sodium thiosulphate, Sodium hyposulphite, Sodii hyposulphis, Na2S203.5H20 — 248. Made by digesting a solution of sodium sulphite with powdered sulphur, when combination slowly takes place : Na2S03 + S = Na2S203. It is used under the name of “hypo,” in photography to dissolve chloride, bromide, or iodide of silver. Disodium hydrogen phosphate, Sodii phosphas, Na2HP04.12H20 = 358 (Phosphate of sodium), is made from calcium phosphate by the action of sulphuric acid, which removes two-thirds of the calcium, forming calcium sulphate, while acid phosphate of calcium is formed and remains in solution : C»32P04 + 2H2S04 = 2CaS04 + CaII42P04. The solution is filtered and sodium carbonate added, when calcium phosphate is precipitated, phosphate of sodium, carbon dioxide, and water being formed: CaH4(P04)2 + Na2CO, = CaHP04 + Ha0 + C02 + Na2HP04. The filtered and evaporated solution yields crystals of phosphate of sodium, which have a slightly alkaline reaction. Experiment 19. Mix thoroughly 30 grammes of bone-ash with 10 c.c. of sul- phuric acid, let stand for some hours, add 20 c.c. of water, and again set aside for some hours. Mix with 40 c.c. of water, heat to the boiling-point, and filter. The residue on the filter is chiefly calcium sulphate. To the hot filtrate of calcium acid phosphate add concentrated solution of sodium carbonate until a precipitate ceases to form and the liquid is faintly alkaline, filter, evaporate, and let crystallize. When sodium phosphate is heated to a low red heat it loses water, and is converted into pyrophosphate, which, dissolved in hot water, and crystallized, forms the sodium pyrophosphate, Na4P207.10H20, of the U. S. P. The normal sodium phosphate, Na3P04, is known also, but it is not a very stable compound, being acted upon even by the moisture and carbon dioxide of the air, with the formation of sodium carbonate and disodium hydrogen phos- phate, thus: 2Na3P04 -f H20 + C02 = 2Na2IIP04 + Na2CO:(. Sodium nitrate, Sodii nitras, NaN03= 85 (Nitrate of sodium, Chili saltpetre, Cubic nitre). Found in nature, and is purified by crystal- 144 METALS AND THEIR COMBINATIONS. lization. The crystals are transparent, deliquescent, and readily soluble. Sodiumborate, Sodii boras, Na2B407 + 10H20 = 382.2 (Borax). This salt occurs in Clear Lake, Nevada, and in several lakes in Asia. It is manufactured by adding sodium carbonate to the boric acid found in Tuscany, Italy. It forms colorless, transparent crystals, but is sold mostly in the form of a white powder. It is slighly efflorescent, is soluble in 16 parts of cold, and in 0.5 part of boiling water; insolu- ble in alcohol, but soluble in one part of glycerin at 80° C. (176° F.). When heated, borax puffs up, loses water of crystallization, and at red heat it melts, forming a colorless liquid which, on cooling, solidi- fies to a transparent mass, known as fused borax, or borax glass. Molten borax has the power to combine with metallic oxides, forming double borates, some of which have a characteristic color, for which reason borax is used in blow-pipe analysis. Borax has antiseptic properties, preventing the decomposition of some organic substances. Other sodium salts which are officinal are sodium hypophosphite, NaPH202 -f H20; bromide, NaBr; iodide, Nal; chlorate, NaC103. These salts may be obtained by processes analogous to those given for the corresponding potassium compounds. (Sodium chloride, NaCl, may be used.) Tests for sodium. 1. As all salts of sodium are soluble in water, we cannot precipi- tate this metal in the form of a compound by any of the common reagents. (Potassium antimoniate precipitates neutral solution of so- dium salts, but this test is not reliable.) 2. The chief reaction for sodium is the flame-test, compounds of sodium imparting to a colorless flame yellow color, which is very intense, brilliant, and luminous. A crystal of potassium dichromate appears colorless, and a paper coated with red mercuric iodide appears white, when illuminated by the yellow sodium flame. (The spectro- scope shows a characteristic yellow line.) 3. Sodium compounds are white and are not volatile at or below a red heat. Lithium Li = 7. Found in nature in combination with silicic acid in a few rare minerals or as a chloride in some spring waters. Of inorganic salts, the bromide and carbonate are officinal. Hydroxide, carbonate, and phosphate of lithium are much less soluble than the corresponding compounds of potassium AMMONIUM. 145 and sodium. Sodium phosphate added to a strong solution of a lithium salt produces, on boiling, a white precipitate of lithium phosphate, Li3P04. Lithium compounds color the flame a beautiful crimson or carmine-red. 21. AMMONIUM. NH4i = 18. General remarks. The salts of ammonium show so much resem- blance, both in their physical and chemical properties, to those of the alkali-metals, that they may be studied most conveniently at this place. The compound radical NH4 acts in these ammonium salts very much like one atom of an alkali-metal, and, therefore, frequently has been looked upon as a compound metal. The physical metallic prop- erties (lustre, etc.) of ammonium cannot be fully demonstrated, as it is not capable of existing in a separate or free state. There is known, however, an alloy of ammonium and mercury, which may be obtained by dissolving potassium in mercury, and adding to the potassium- amalgam thus formed, a strong solution of ammonium chloride, when potassium chloride and ammonium-amalgam are formed. The latter is a soft, spongy, metallic-looking substance, which readily decomposes into mercury, ammonia, and hydrogen : HgK + NH4C1 = KC1 + NH4Hg; NII4Hg = NH3 + H + Hg. The source of all ammonium compounds is ammonia NH3, or am- monium hydroxide, NH4OH, both of which have been considered heretofore. Ammonium chloride, Ammonii chloridum, NH4C1 = 53.4 (Chloride of ammonium, Sal-ammoniac). Obtained by saturating the “ ammonia- cal liquor” of the gas-works with hydrochloric acid, evaporating to dryness, and purifying the crude article by sublimation. Questions.—191. What is the composition of common salt; how is it found in nature, and what is it used for? 192. Describe Leblanc’s and the Solvay process for manufacturing sodium carbonate on a large scale. 193. How much water is in 100 pounds of the crystallized sodium carbonate? 194. What is Glauber’s salt, and how is it made? 195. State the composition of disodium hydrogen phosphate, and how is it prepared from calcium phosphate? 196. What difference exists between sodium carbonate and bicarbonate, both in regard to physical and chemical properties? 197. Give the composition of sodium hyposulphite; what is it used for? 198. Which sodium salts are soluble, and which are insoluble? 199. How does sodium and how does lithium color the flame? 200. Which lithium salts are officinal ? 146 METALS AND THEIR COMBINATIONS. Pure ammonium chloride either is a white, crystalline powder, or occurs in the form of long, fibrous crystals, which are tough and flexible; it has a cooling, saline taste; is soluble in 3 parts of cold, and in 1 part of boiling-water; and like all ammonium compounds, is completely volatilized by heat. Experiment 20. To 10 c.c. of water of ammonia add hydrochloric acid until the solution is neutral to test paper. Evaporate to dryness and use the salt for the analytical reactions mentioned below. How many c.c. of 32 per cent, hydrochloric acid are required to saturate 10 c.c. of 10 per cent, ammonia water ? Ammonium carbonate, Ammonii carbonas, NH4HC03.NH4NH2C02 = 157 [Carbonate of ammonium). Commercial ammonium carbonate is not the normal salt, but, as shown by the above formula, a combi- nation of acid ammonium carbonate with ammonium carbamate. It is obtained by sublimation of a mixture of ammonium chloride and calcium carbonate, when calcium chloride is formed, ammonia gas and water escape, and ammonium carbonate condenses in the cooler part of the apparatus : 2CaC03 + 4NH4C1 == NH4HC03.NH4NH2C02 + 2CaCI2 + H20 + NH3. Ammonium carbonate thus obtained forms white, translucent masses, losing both ammonia and carbon dioxide on exposure to the air, becoming opaque, and finally converted into a white powder of acid ammonium carbonate: NH4HC03 NH4NH2C02 = NH4HC03 + 2NH3 + co2. When commercial ammonium carbonate is dissolved in water, the carbamate unites with one molecule of water, forming normal ammonium carbonate: NH4NH2C02 + II20 = (NH4),C03. A solution of the common ammonium carbonate in water is, consequently, a liquid containing both acid and normal carbonate of ammonium; by the addi- tion of some ammonia water the acid carbonate is converted into the normal salt. The solution thus obtained is used frequently as a reagent. The Aromatic spirit of ammonia (sal volatile) is a solution of normal ammo- nium carbonate in diluted alcohol to which some essential oils have been added. Ammonium sulphate, (NH4)2S04, Ammonium nitrate, NH4N03, and Ammonium phosphate, (NH4)2HP04, may be obtained by the addition of the respective acids to ammonia water or ammonium carbonate: H2S04 + 2NH4OJI = (NH4)2S04 + 2H20. i-ino3 + nii4oii = nh4no3 + h2o. H3P04 -f 2NII4OH = (NII4)2HP04+ 2H20. H2S04 + (NH4)2C03 = (NH4)2S04 + H20 + co2. AMMONIUM. 147 Ammonium iodide, Ammonii iodidum, NH4I, and Ammonium bro- mide, Ammonii bromidum, NH4Br, may be obtained by mixing together strong solutions of potassium iodide (or bromide) and ammonium sul- phate, and adding alcohol, which precipitates the potassium sulphate formed ; by evaporation of the solution the ammonium iodide (or bromide) is obtained : 2KI + (NH4)2S04 = 2NH4I + K2S04; 2KBr + (NH4)2S04 = 2NH4Br + K2S04. Another mode of preparing these compounds is by the decomposi- tion of ferrous bromide (or iodide) by ammonium hydroxide: FeBr2 + 2NH.0H = 2NII4Br + Fe(OH)2. Ammonium iodide is the principal constituent of the Decolorized tincture of iodine. Ammonium hydrogen sulphide, NH4SH (Ammonium hydrosulphide, Ammonium sulphydrate). Obtained by passing hydrogen sulphide through water of ammonia until this is saturated : h2s + nh4oii = nh4sh + h2o. The solution thus obtained is, when recently prepared, a colorless liquid, having the odor of both ammonia and of hydrogen sulphide; when exposed to the air it soon assumes a yellow color. By the addition of ammonia water it is converted into ammonium sulphide, (NH4)2S: Both substances, the ammonium hydrogen sulphide and ammonium sul- phide, are valuable reagents, frequently used for precipitation of certain heavy metals, or for dissolving certain metallic sulphides. NH4SH + NH4OH z= (NH4)2S + H20. Analytical reactions. 1. All compounds of ammonium are volatilized by heating to a low red heat. 2. All compounds of ammonium evolve ammonia gas when heated with hydroxide of calcium, potassium, or sodium. The ammonia may be recognized by its odor, or by its action on paper moistened with solution of cupric sulphate, which is thereby colored dark-blue, or by causing the appearance of dense white fumes of ammonium chloride, upon holding a glass rod, moistened with hydrochloric acid, in the gas. 3. Add to solution of ammonium salt some platinic chloride, a few (Ammonium chloride, NH4C1, may be used.) 148 METALS AND THEIR COMBINATIONS. drops of hydrochloric acid, and some alcohol; a yellow precipitate of ammonium platinic chloride, (NH4Cl)2PtCl4, is produced. See ex- planation of the corresponding potassium reaction on page 139. 4. The addition of sodium cobaltic nitrite causes in neutral or acid solutions a yellow precipitate of ammonium cobaltic nitrite, (NH4N02)6.Co2.(N02)6. 5. Ammonium salts are colorless, and (almost all) soluble in water. Traces of ammonium compounds may be detected by alkaline mercuric- potassium iodide (Nessler’s solution), which causes a reddish-brown precipitate or coloration. Summary of analytical characters of the alkali-metals. Potassium. Sodium. Lithium. Ammonium. Sodium cobaltic nitrite . . Platinic chloride .... Sodium bitartrate . . . Sodium phosphate . . . Sodium hydroxide . . . Action of heat .... Yellow pre- cipitate. Yellow pre- cipitate White pre- cipitate. Yellow pre- cipitate. Yellow pre- cipitate. White pre- cipitate. White preci- pitate in cone solution on boiling. Ammonia gas. Volatile. Fusible. V iolet. Fusible. Yellow. Fusible. Crimson. 22. MAGNESIUM. Mgii == 24.3. General remarks. Magnesium occupies a position intermediate between the alkali metals and the alkaline earths, with which latter it Questions.—201. What is ammonium, and why is it classed with the alkali- metals? 202. Is ammonium known in a separate state? 203. What is ammo- nium-amalgam, how is it obtained, and what are its properties? 204. What is the source of ammonium compounds? 205. State the composition, mode of preparation, and properties of sal-ammoniac. 206. How is ammonium car- bonate manufactured, and what difference exists between the solid article and its solution ? 207. State the composition of ammonium sulphide and of am- monium hydrogen sulphide; how are they made, and what are they used for? 208. By what process may ammonium sulphate, nitrate, and phosphate be obtained from ammonium hydroxide or ammonium carbonate, and what chemical change takes place? 209. How does heat act upon ammonium com- pounds? 210. Give analytical reactions for ammonium salts. MAGNESIUM. 149 was formerly classed as a member. To some extent it resembles also the heavy metal zinc, with which it has in common, the volatility of the chloride, the solubility of the sulphate, and the isomorphism of several of its compounds with the analogously constituted compounds of zinc. Occurrence in nature. Magnesium is widely diffused in nature, and several of its compounds are found in large quantities. It occurs as chloride and sulphate in many spring waters and in the salt-mines of Stassfurt; as carbonate in the mineral magnesite ; as double carbonate of magnesium and calcium in the mineral dolomite (magnesian-lime- stone), which forms entire mountains; as silicate of magnesium in the minerals serpentine, meerschaum, talc, asbestos, soapstone, etc. Metallic magnesium may be obtained by the decomposition of mag- nesium chloride by sodium : MgCl2 + 2Na = 2NaCl + Mg. Magnesium is an almost silver-white metal, losing its lustre rapidly in moist air by oxidation of the surface. It decomposes hot water with liberation of hydrogen : Mg + 2H20 = 2H + Mg(OH)2. When heated to a red heat it burns with a brilliant bluish-white light forming magnesium oxide. Magnesium carbonate, Magnesii carbonas. Approximately: 4(MgC03).Mg(0H)25H20 (Carbonate of magnesium, Magnesia alba, Light magnesia). The normal magnesium carbonate, MgCOa, is found in nature, but the officinal preparation is a mixture of carbonate, hydrox- ide, and water. It is obtained by boiling a solution of magnesium sulphate with solution of sodium carbonate, when the carbonate is precipitated, some carbon dioxide evolved, and sodium sulphate remains in solution : 6MgS<>4 + 5Nf>2C03 4. 6H20 = 4(MgC0,).Mg(0H)2.6H20 + 5Na2S044 C02. * By filtering, washing, and drying the precipitate, it is obtained in the form of a white, light powder; if, however, the above-mentioned solutions are mixed, evaporated to dryness, and the sodium sulphate removed by washing, the magnesium carbonate is left in a more dense condition, and is then known as heavy magnesia. Experiment 21. Dissolve 10 grammes of magnesium sulphate in hot water and add a concentrated solution of sodium carbonate until no more precipitate 150 METALS AND THEIR COMBINATIONS. is formed. Collect the precipitated magnesium carbonate on a filter and dry it at a low temperature. (How much crystallized sodium carbonate is needed for the decomposition of 10 grammes of crystallized magnesium sulphate?) Notice that the dried precipitate evolves carbon dioxide when heated with acids. Magnesium oxide, Magnesia, MgO =40.3 (Calcined magnesia, Light magnesia), is obtained by heating light magnesium carbonate in a crucible to a full red heat, when all carbon dioxide and water are expelled: 4(MgC03).Mg(0H)2.5H20 = 5MgO + 4C02 + 6H20. It is a very light, amorphous, white, almost tasteless powder, which absorbs moisture and carbon dioxide gradually from the air ; iu contact with water it forms the hydroxide Mg(OH)2, which is almost insoluble in water, requiring of the latter over 50,000 parts for solution. Milk of magnesia is the hydroxide suspended in water (1 part in about 15). The heavy magnesia, magnesia ponderosa of the U. S. P., differs from the common or light magnesia, not in its chemical composition, but merely in its physical condition, being a white, dense powder obtained by heating the heavy magnesium carbonate. Experiment 22. Place 1 gramme of magnesium carbonate, obtained in per- forming Experiment 21, into a weighed crucible and heat to redness, or until by further heating no more loss in weight ensues. Treat the residue with dilute hydrochloric acid and notice that no evolution of carbon dioxide takes place. What is the calculated loss in weight of magnesium carbonate when converted into oxide, and how does this correspond with the actual loss deter- mined by the experiment? Magnesium sulphate, Magnesii sulphas, MgS04.7H20 = 246.3 (Sul- phate of Magnesium, Epsom salt), is obtained from spring waters, from the mineral Kieserite, MgS04.H20, and by decomposition of the native carbonate by sulphuric acid : MgC03 + H2S04 = MgS04 + C02 + H20. It forms colorless crystals, which have a cooling, saline, and bitter taste, a neutral reaction, and are easily soluble in water. Magnesium sulphite, Magnesii sulphis, MgS036H20 = 212.3, may be obtained by adding sulphurous acid to magnesium carbonate : MgC03 + H2S03 = MgS03 + C02 + H20. CALCIUM. 151 Analytical reactions. (Magnesium sulphate, MgS04, may be used.) 1. Add to a magnesium solution potassium or sodium carbonate and heat: a white precipitate of basic magnesium carbonate, 4MgC03. Mg(OII)2, is produced. 2. Add to a magnesium solution ammonium carbonate (or ammo- nium hydroxide): part of the magnesium will be precipitated as carbonate (or hydroxide). These precipitates, however, are soluble in ammonium chloride and many other ammonium salts : if these latter had been added previously to the magnesium solution, ammonium carbonate (or hydroxide) would cause no precipitation. (The dissolv- ing action of the ammonium chloride is due to the tendency of mag- nesium to form double salts with ammonium salts.) 3. To solution of magnesium add a solution containing sodium O O phosphate, ammonium chloride, and ammonia : a white crystalline precipitate of magnesium ammonium phosphate, MgNH4P04, is pro- duced, which is somewhat soluble in water, but almost insoluble in water containing some ammonia. 4. Salts of magnesium are white and soluble, except the carbonate, phosphate, and arseniate; the oxide and hydroxide also are insoluble; the latter is precipitated by sodium or potassium hydroxide. 23. CALCIUM. Ca.ii = 40 (39.91). General remarks regarding the metals of the alkaline earths. The three metals, calcium, barium, and stroutium, form the second group of light metals. Similar to the alkali-metals, they decompose water at the ordinary temperature with liberation of hydrogen ; their separation in the elementary state is even more difficult than that of the alkali-metals. They differ from the latter by forming insoluble carbonates and Questions.—211. How is magnesium found in nature? 212. By what pro- cess is metallic magnesium obtained? 213. Give the physical and chemical properties of magnesium. 214. State two methods by which magnesium oxide can be obtained. 215. What is calcined magnesia? 216. State the composition and properties of the officinal magnesium carbonate, and how it is made. 217. What is Epsom salt, and how is it obtained? 218. Which compounds of mag- nesium are insoluble? 219. Give tests for magnesium compounds. 220. How can the presence of magnesium be demonstrated in a mixture of magnesium sulphate and sodium sulphate? 152 METALS AND THEIR COMBINATIONS. phosphates (those of the alkalies are soluble), from the earths by their soluble hydroxides (those of the earths are insoluble), and from all heavy metals by the solubility of their sulphides (those of heavy metals are insoluble). The sulphates are either insoluble (barium) or spar- ingly soluble (strontium and calcium). The hydroxides and carbon- ates are decomposed by heat, water or carbon dioxide being expelled and the oxides formed. In case of calcium carbonate this decompo- sition takes place easily, while the carbonates of barium and strontium require a much higher temperature. They are bivalent elements. Occurrence in nature. Calcium is one of the most abundantly occurring elements. As carbonate (CaC03) it is found in the form of calc-spar, limestone, chalk, marble, shells of eggs and mollusca, etc., or, as acid carbonate, dissolved in water. The sulphate is found as gypsum or alabaster (CaS042H20); the phosphate (Ca32P04) in the different phosphatic rocks (apatite, etc.); the fluoride (CaF2) as fluor- spar ; the chloride (CaCl2) in some waters, and the silicate in many rocks. It enters the vegetable and animal system in various forms of combination, chiefly, however, as phosphate and snlphate. Calcium oxide, Lime, Calx, CaO = 56 (Oxide of calcium, Quick- lime, Burned lime), is obtained on a large scale by the common process of lime-burning, which is the heating of limestone or any other calcium carbonate to about 800° C. (1472° F.), in the so-called lime-kilns. On a small scale decomposition may be accomplished in a suitable crucible over a blowpipe flame : CaC03 = CaO + C02. The pieces of oxide thus formed retain the shape and size of the carbonate used for decomposition. Lime is a white, odorless, amorphous, infusible substance, of alkaline taste and reaction; exposed to the air it gradually absorbs moisture and carbon dioxide, the mixture thus formed being known as air- slaked lime. Lime occupies among bases a position similar to that of sulphuric acid among acids, and is used directly or indirectly in many branches qf chemical manufacture. Calcium hydroxide, Calcium hydrate, Ca(OH)2 (Slaked lime). When water is sprinkled upon pieces of calcium oxide, the two substances combine chemically, liberating much heat; the pieces swell up, and CALCIUM. 153 are converted gradually into a dry, white powder, which is the slaked lime. When this is mixed with water, the so-called milk of lime is formed. Lime-water, Liquor calcis. This is a saturated solution of calcium hydroxide in water: 10,000 parts of the latter dissolving about 15 parts of hydroxide. In making lime-water, 1 part of calcium oxide is slaked and stirred for about half an hour with 30 parts of water. The mixture is then allowed to settle, and the liquid, containing besides calcium hydroxide the salts of the alkali-metals which may have been present in the lime, is decanted and thrown away. To the calcium hydroxide left, and thus purified, 300 parts of water are added and occasionally shaken in a well-stoppered bottle, from which the clear liquid may be poured off for use. Lime-water is a colorless, odorless liquid, having a feebly caustic taste, and an alkaline reaction. When heated to boiling it becomes turbid by precipitation of calcium hydroxide (or perhaps oxide), which re-dissolves when the liquid is cooled. Carbon dioxide causes a pre- cipitation of calcium carbonate. Experiment 23. Make lime-water according to directions given above. Calcium carbonate, Calcii carbonas praecipitatus, CaC03 = 100 (Carbonate of calcium). Precipitated calcium carbonate is obtained as a white, tasteless, neutral, impalpable powder by mixing solutions of calcium chloride and sodium carbonate : CaCl2 + Na2COs = 2NaCl + CaC03. Experiment 24. Add to about 10 grammes of marble (calcium carbonate) in small pieces, hydrochloric acid as long as effervescence takes place; filter the solution of calcium chloride thus obtained and add to it solution of sodium carbonate as long as a precipitate is formed, collect the precipitate on a filter, wash and dry it. Calcium sulphate, CaS04 = 136 (Anhydrous sulphate of calcium, Plaster-of-Paris, Calcined plaster). It has been mentioned above that the mineral gypsum is native calcium sulphate in combination with 2 molecules of water of crystallization. By heating to about 115° C. (239° F.) this water is expelled, and the anhydrous sulphate formed. It readily recombines with water, becoming a hard mass, for which reason it is used for making moulds and casts, and in surgery. For the latter purpose plaster is often mixed with alum and gelatin before adding the water, this mixture being preferred on account of 154 METALS AND THEIR COMBINATIONS. forming a harder, less porous mass, with a smooth surface that can be washed with water containing disinfecting agents. Tricalcium phosphate, Calcii phosphas prsecipitatus, Ca3(P04)2 = 310 (.Precipitated phosphate of calcium, Phosphate of lime, Bone-phos- phate). By dissolving bone-ash (bone from which all organic matter has been expelled by heat) in hydrochloric acid, and precipitating the solution with ammonia water there is obtained calcium phosphate, which contains traces of calcium fluoride and magnesium phosphate. A pure article is made by precipitating a solution of calcium chloride by sodium phosphate and ammonia. 2Na2HP04 + 3CaCl2 + 2NII4OH = Ca3(P04)a + 4NaCl + 2NH4C1 + 2H20. It is a white, tasteless, amorphous powder, insoluble in cold water, soluble in hydrochloric or nitric acids. Superphosphate, or acid phosphate of lime. Among the inorganic substances which serve as plant-food, calcium phosphate is a highly important one. As this compound is found usually in very small quantities as a constituent of the soil, and as this small quantity is soon removed by the various crops taken from a cultivated soil, it becomes necessary to replace it in order to enable the plant to grow and to form seeds. For this purpose the various phosphatic rocks (chiefly calcium phosphate) are converted into commercial fertilizers, which is accomplished by the addi- tion of sulphuric acid to the ground rock. The sulphuric acid removes from the tricalcium phosphate one or two atoms of calcium, forming mono- or dicalcium phosphate and calcium sulphate. The mixture of these substances, containing also the impurities originally present in the phosphatic rocks, is sold as acid phosphate or superphosphate. Bone-black and bone-ash. Phosphates enter the animal system in the various kinds of food, and are to be found in every tissue and fluid, but most abundantly in the bones and teeth. Bones contain about 30 per cent, of organic and 70 per cent, of inorganic matter, most of which is tricalcium phosphate. When bones are burned until all the organic matter has been destroyed and volatilized, the resulting product is known as bone-ash. If, however, the bones are subjected to the process of destructive distillation (heating with exclusion of air), the organic matter suffers decomposition, many volatile products escape, and most of the non-volatile carbon remains mixed with the inorganic portion of the bones, which substance is known as bone-black or animal charcoal. Calcium hypophosphite, Calcii hypophosphis, Ca(PH202)2 = 170. Obtained by heating pieces of phosphorus with milk of lime until CALCIU M. 155 phosphoretted hydrogen ceases to escape. From the filtered liquid the excess of lime is removed by carbon dioxide, and the clear liquid evaporated to dryness. (Great care must be taken during the whole of the operation, which is somewhat dangerous on account of the in- flammable and explosive nature of the compounds.) 8P + 6H20 4- 3[Ca(OH)2] = 3[Ca(PH202)2] + 2PH3. Calcium chloride, Calcii chloridum, CaCl2 = 110.8, and Calcium bromide, Calcii bromidum, CaBr2 = 199.6, may both be obtained by dissolving calcium carbonate in hydrochloric acid or hydrobromic acid, until the acids are neutralized. Both salts are highly deliques- cent. Chlorinated lime, Calx chlorata (Bleaching-powder, incorrectly called Chloride of lime). This is chiefly a mixture (according to some, a compound) of calcium chloride with calcium hypochlorite, and is manufactured on a very large scale by the action of chlorine upon calcium hydroxide : 2Ca(0H), + 4C1 = 2H20 + Ca(C10)2 + CaCl2. Bleaching-powder is a white powder, having a feeble chlorine-like odor; exposed to the air it becomes damp from absorption of moist- ure, undergoing decomposition at the same time; with dilute acids it evolves chlorine, of which it should contain not less than 30 per cent, in available form. The action of hydrochloric acid takes place thus : Calcium hydroxide. Chlorinated lime. Ca(C10)2 + 2HC1 = CaCl2 + 2HC10; 2HC10 + 2HC1 = 2H20 + 401. Bleaching-powder is a powerful disinfecting and bleaching agent. Sulphurated lime, Calx sulphurata, is a mixture of calcium sulphide and sulphate, obtained by beating in a crucible a mixture of equal parts of sulphur aud calcium oxide. Analytical reactions. (Calcium chloride, CaCl2, may be used.) 1. Add to solution of a calcium salt, the carbonate of either potas- sium, sodium, or ammonium : a white precipitate of calcium carbon- ate, CaC03, is produced. 156 METALS AND THEIR COMBINATIONS. 2. Add sodium phosphate to neutral solution of calcium: a white precipitate of calcium phosphate, CaHP04, is produced. 3. Add ammonium (or potassium) oxalate to a calcium solution : a white precipitate of calcium oxalate, CaC204, is produced, which is insoluble in acetic, soluble in hydrochloric acid. 4. Sulphuric acid or soluble sulphates produce a white precipitate of calcium sulphate, CaS04, in concentrated, but not in dilute solutions of calcium. 5. Add potassium or sodium hydroxide : a white precipitate of cal- cium hydroxide, Ca20H, is produced in concentrated but not in diluted solutions. Ammonia water gives no precipitate. 6. Calcium compounds impart a reddish-yellow color to the flame. Strontium, SrH = 87.3. Found in a few localities in the minerals strontianite, SrC03, and celestite, SrS04. Its compounds resemble those of calcium and barium. The oxide, SrO, cannot be obtained easily by heating the carbonate, as this is much more stable than calcium car- bonate. It may, however, be readily prepared by heating the nitrate. The hydroxide, Sr20H, is formed when the oxide is brought in con- tact with water; it is more soluble than calcium hydroxide. Strontium nitrate, Sr(N03)2, Strontium chloride, SrCl2, and Strontium bromide, SrBr2, may be obtained by dissolving the carbonate in the respective acids. The nitrate is used extensively for pyroteehnical purposes, as strontium imparts a beautiful red color to flames. A number of strontium salts are now also used medicinally. Analytical reactions. (Strontium nitrate, Sr2N03, may be used.) 1. The reactions of strontium with soluble carbonates, oxalates, and phosphates are analogous to those of calcium. 2. Add calcium sulphate : a white precipitate of strontium sulphate, SrS04, is formed after a few minutes. 3. Add sulphuric acid or a soluble sulphate: a white precipitate forms at once in concentrated, after a while in dilute solutions. 4. Add potassium chromate : a pale-yellow precipitate of strontium chromate, SrCr04, is formed, which is soluble in acetic acid and in hydrochloric acid. (Potassium dichromate causes no precipitation.) 5. Strontium compounds color the flame beautifully red. CALCIUM. 157 Barium, Bau = 136.9. Occurs in nature chiefly as sidphate in barite or heavy spar, BaS04, but also as carbonate in witherite, BaC03. Barium and its compounds reserhble closely those of calcium and strontium. A Barium chloride, BaCl2 + 2H20, is prepared by dissolving the carbonate in hydrochloric acid. It crystallizes in prismatic plates, and is used as a valuable reagent. Barium dioxide or peroxide, Ba02, is made by heating the oxide to a dark-red heat in the air or in oxygen. When heated above the tem- perature at which it is formed decomposition into oxide and oxygen takes place. This power to absorb oxygen from air and to give it up again at a higher temperature has been used as a method of preparing oxygen on the large scale. Unfortunately, the barium oxide cannot be used for an unlimited number of operations, as it loses the power to absorb oxygen after it has been heated a number of times. The use made of barium dioxide in preparing hydrogen dioxide has been mentioned before. Barium oxide, BaO, is made by heating barium nitrate, Ba2N(>3, which itself is made by dissolving the carbonate in nitric acid. Barium salts are poisonous; antidotes are sodium and magnesium sulphate. Analytical reactions. (Barium chloride, BaCl2, may be used.) 1. The reactions of strontium with soluble carbonates, oxalates, and phosphates are analogous to those of calcium solutions. 2. Add sulphuric acid or soluble sulphates : a white precipitate of barium sulphate, BaS04, is produced immediately, even in dilute solu- tions. The precipitate is insoluble in all diluted acids. 3. Add calcium sulphate: a white precipitate, insoluble in all diluted acids, is formed immediately. 4. Add potassium chromate or dichromate : a pale-yellow precipi- tate of barium chromate, BaCr()t, is formed, which is soluble in hydrochloric acid. 5. Barium compounds color the flame yellowish-green. 158 METALS AND THEIR COMBINATIONS. Summary of analytical characters of the alkaline earth-metals. Magnesium. Calcium. Strontium. Barium. Potassium dichromate . . Potassium chromate. . . Calcium sulphate.... Yellow pre- cipitate. White pre- cipitate form- ing slowly. Yellow pre- cipitate. Yellow pre- cipitate. White pre- cipitate form- ing at once. Ammonium carbonate . . White pre- White pre- White pre- White pre- Ammonium hydroxide cipitate solu- ble in NH4C1. White pre- cipitate. cipitate. cipitate. cipitate. Ammonium oxalate . . . No precipi- White pre- White pre- W lute pre- tate unless cipitate in di- cipitate in cipitate in very con- centrated. lute solution. strong solution. strong solution. Sodium phosphate . . . White pre- White pre- White pre- White pre- cipitate. cipitate. cipitate. cipitate. Flame color Yellowish- red. lied. Yellowish- green. 24. ALUMINIUM [Aluminum). Aliii 27 (27.04). Aluminium is the representative of the metals of the earths proper ; all other members of this class are found in nature in very small quan- tities, and are chiefly of scientific interest, with the exception of cerium, which furnishes au officinal preparation. Occurrence in nature. Aluminium is found almost exclusively in the solid mineral portion of the earth; rarely more than traces of Questions.—221. Which metals form the group of the alkaline earths, and in what respect do their compounds differ from those of the alkali-metals? 222. How is calcium found in nature? 223. What is burned lime; from what, and by what process is it made, and how does water act on it? 224. What is lime-water; how is it made, and what are its properties? 225. Mention some varieties of calcium carbonate as found in nature, and how is it obtained by an artificial process from the chloride? 226. What is Plaster-of-Paris, and what is gypsum ; what are they used for? 227. State composition and mode of manufacturing bleaching-powder; what are its properties, and how do acids act upon it? 228. What is bone-black, bone-ash, acid phosphate, and precipitated tricalcium phosphate? How are they made? 229. Give tests for barium, calcium, and strontium; how can they be distinguished from each other? 230. Which compounds of barium and strontium are of interest, and what are they used for? ALUMINIUM. 159 aluminium compounds are found dissolved in water, and the occur- rence of aluminium in either the vegetable or animal organism seems to be purely accidental. By far the largest quantity of aluminium is found in combination with silicic acid in the various silicated rocks forming the greater mass of our earth, such as feldspar, slate, basalt, granite, mica, hornblende, etc., or in the various modifications of clay formed by their decompo- sition. The minerals known as corundum, ruby, sapphire, and emery, are aluminium oxide in a crystallized state, and more or less colored by traces of other substances. Metallic aluminium may be obtained by the decomposition of alu- minium chloride by metallic sodium : A]2ci6 + 6Na = 6NaCl + 2A1. It is now manufactured by the electrolysis of aluminium and sodium fluoride. Aluminium is an almost silver-white metal of a very low specific gravity (2.67); it is capable of assuming a high polish, and for this reason is used for ornamental articles; it is very strong, yet malleable, and does not change in dry or moist air. Some of the alloys of aluminium are now used in the arts, as, for instance, aluminium-bronze, an alloy resembling gold and composed of 10 parts of aluminium with 90 of copper. Aluminium is trivalent, and shows, like a number of other elements (iron, chromium, etc.), the peculiarity that the double atom Al2vi acts as a single sexivalent atom. Alum is the general name for a group of isomorphous salts, com- posed of one molecule of the sulphate of a univalent metal in combi- nation with one molecule of the sulphate of a trivalent metal, combined in crystallizing with 24 molecules of water. The general formula of an alum is consequently M‘2S04M2iii(S04)3.24H20. M* rep- resents in this case a univalent, MiU a-trivalent metal. Alums known are, for instance : Potassium-aluminium sulphate, K2S04, A12(S04)3.24II20. Ammonium-aluminium sulphate, (N'II4)2S04, A12(S04)3.24H20. Potassium-chromium sulphate, K2S04, Cr2(S04)3 241120. Ammonium-ferric sulphate, (NH4)2S04, Fe2(S04)3 24H20. The officinal alum, alumen, is the potassium alum, a white substance crystallizing in large octahedrons, soluble in 10 parts of cold and 0.3 160 METALS AND THEIR COMBINATIONS. part of boiling water; this solution has an acid reaction and a sweetish astringent taste. Alum is manufactured on a large scale by decomposing certain kinds of aluminium silicates by sulphuric acid, when aluminium sul- phate is formed, to the solution of which potassium or ammonium sulphate is added, when, on evaporation, potassium or ammonium alum crystallizes. Dried alum, Alumen exsiccatum, K2S04.A12(S04)3 = 516. This is common alum, from which the water of crystallization has been ex- pelled by heat. It is a white powder dissolving very slowly in cold, but quickly in boiling water. Aluminium hydroxide, Aluminium hydrate, Alumini hydras, Al2 (0H)6 —156. Obtained by adding water of ammonia or solution of sodium carbonate to solution of alum, when aluminium hydroxide is precipitated in the form of a highly gelatinous substance, which, after being well washed, is dried at a temperature not exceeding 40° C. (104° F.). K2S04.A12(S04)3 + 6NH4OH = K2S04 + 3[(NH4)2S04] + AI2(OH)6; K2S04A12(S04)s + 3Na2C03 + 3H20 = K2S04 + 3Na2S04 + 3C02 + Al2(OH)6. The usual decomposition between a soluble carbonate and any soluble salt (provided decomposition takes place at all) is the formation of an insoluble carbonate; according to this rule, the addition of a soluble carbonate to alum should produce aluminium carbonate. The basic properties of aluminium oxide, however, are so weak that it is not capable of uniting with so weak an acid as carbonic acid, and it is for this reason that the decomposition takes place as shown by the above formula, with liberation of carbon dioxide, whilst the hydroxide is formed. (Other metals, the oxides of which have weak basic properties, show similar reactions, as, for instance, chromium, and iron in the ferric salts.) The weak basic properties of aluminium are shown also by the fact that alu- minium sulphate, chloride, and nitrate, and even alum itself, have an acid reaction, while the corresponding salts of the alkalies or alkaline earths are neutral. Aluminium hydroxide shows considerable surface-attraction toward many substances, which property is made use of in the art of dyeing, where the hy- droxide is used for retaining coloring matter upon the cotton-fibre. Practically this is accomplished by precipitating aluminium hydroxide from solutions containing coloring matter, which latter is carried down and precipitated upon the fibre by the aluminium hydroxide; or by impregnating the articles to be dyed with this compound and placing them in the colored solutions. Experiment 25. Dissolve 10 grammes of sodium carbonate in 150 c.c. of water, heat it to boiling, and add to it, with constant stirring, a hot solution, made by dissolving 11 grammes of alum in 150 c.c. of water. Wash the pre- ALUMINIUM. 161 cipitate first by decantation, and then upon a filter, until the washings are not rendered turbid by barium chloride. Dry a portion of the precipitate at a low temperature, and use as aluminium hydroxide. Mix a small quantity of the wet precipitate with a decoction of logwood (made by boiling about 0.2 grammes of logwood with 50 c.c. of water), agitate for a few minutes, and filter. Notice that the red color of the solution has entirely disappeared, or nearly so, in consequence of the great surface-attraction of the aluminium hydroxide for coloring matter. Aluminium oxide, A1203= 102 [Alumina), is obtained as a white, tasteless powder either by burning the metal or by expelling the water from the hydroxide by heat: Al2(OH)6 = A1203 + 3H20. Aluminium sulphate, Alumini sulphas, A12(S04)3.16H20 — 630. A white crystalline powder, soluble in about its weight of water; ob- tained by dissolving the oxide or hydroxide in sulphuric acid. Al2(OH)6 + 3H2S04 = A12(S04)3 + 6H20. Aluminium chloride, A12C16 = 267. This compound is of interest on account of being the salt from which the metal was formerly obtained. Most chlorides may be obtained by dissolving the metal, its oxide, hydroxide, or carbonate in hydrochloric acid. Accordingly aluminium chloride may be obtained in solution : Al2(OH)6 + 6HC1 = A12C16 + 6H20. On evaporating the solution to dryness, however, and heating the dry mass further with the view of expelling all water, decomposition takes place, hydrochloric acid escapes, and aluminium oxide is left: A12CI6 + 3H20 = A1203 + 6HC1. Aluminium chloride, consequently, cannot be obtained in a pure state (free from water) by this process, but it may be made by expos- ing!: to the action of chlorine a heated mixture of aluminium oxide © and carbon. Neither carbon nor chlorine alone causes any decompo- sition of the aluminium oxide, but by the united efforts of these two substances decomposition is accomplished : A1,03 + 8C + 6C1 = 3C0 + A]2C]6. Clay is the name applied to a large class of mineral substances, dif- fering considerably in composition, but possessing in common the two characteristic features of plasticity and the predominance of aluminium silicate in combination with water. The various kinds of clay have been formed in the course of time from such double silicates as feldspar and others, by a process which is partly of a 162 METALS AND THEIR COMBINATIONS. mechanical, partly of a chemical nature, and consists chiefly in the disintegra- tion of rocks and a removal of potassium and sodium by the chemical action of carbonic acid, water, and other agents. The various kinds of clay are used in the manufacture of bricks, earthenware, stoneware, porcelain, etc. The process of burning these substances accom- plishes the hardening by expelling water which is present in the clay. Pure clay is white; the red color of the common varieties is due to the presence of ferric oxide. For china or porcelain, clay is used containing silicates of the alkalies which, in burning, melt, causing the production of a more homoge- neous mass, while in common earthenware the pores, produced by expelling the moisture, remain unfilled. Glass is similar in composition to the better varieties of porcelain. All varieties of glass are mixtures of fusible, insoluble silicates, made by fusing silicic acid (white sand) with different metallic oxides or carbonates, the silicic acid combining chemically with the metals. Sodium and calcium are the chief metals in common glass, though potassium, lead, and others also are frequently used. Color is im- parted to the glass by the addition of certain metallic oxides, which have a coloriug effect, as, for instance, manganese violet, cobalt blue, chromium, green, etc. Ultramarine is a beautiful blue substance, found in nature as the mineral ‘‘lapis lazuli,” which was highly valued by artists as a color before the dis- covery of the artificial process for manufacturing it. Ultramarine is now manufactured on a very large scale by heating a mix- ture of clay, sodium sulphate and carbonate, sulphur, and charcoal in large crucibles, when decomposition takes place and the beautiful blue compound is obtained. As neither of the substances used in the manufacture has a ten- dency to form colored compounds, the formation of this blue ultramarine is rather surprising, and the true chemical constitution of it is yet unknown. Ultramarine is insoluble in water and is decomposed by acids with libera- tion of hydrogen sulphide, which shows the presence of sodium sulphide. A green ultramarine is now also manufactured. Analytical reactions. (A solution of aluminium sulphate, A12(S04)3, or of aluminium chloride, A12C16, may be used.) 1. To solution of an aluminium salt add potassium or sodium hydroxide: a white gelatinous precipitate of aluminium hydroxide, Al2(OH)6, is produced, which is soluble in excess of the alkali. 2. To aluminium solution add ammonium hydroxide : the same precipitate as above is obtained, but it is insoluble in an excess of the reagent. ALUMINIUM. 163 3. The carbonates of ammonium, sodium, or potassium produce the same precipitate with liberation of carbon dioxide. (See explana- tion above.) 4. Ammonium sulphide produces the same precipitate with libera- tion of hydrogen sulphide: A12C16 + 3(NH4)2S + 6H20 = Al2(OH)6 + 6NH4C1 + 3H2S. 5. Sodium phosphate produces a precipitate of aluminium phos- phate, soluble in acids. Cerium, Ce = 141. This element occurs in nature sparingly in a few rare minerals, chiefly as silicate in cerite. In its general deportment cerium resem- bles aluminium. Cerous solutions give with either ammonium sulphide, or ammonium and sodium hydroxide, a white precipitate of cerous hydroxide, Ce2(OH)6. Ammonium oxalate forms a white precipitate of cerous oxalate, Ce2(C204)39H20, which is the only officinal cerium preparation. Cerium oxa- late is a white, granular powder, insoluble in water and alcohol, but soluble in hydrochloric acid. Exposed to a red heat it is converted into yellow oxide of cerium. If this oxide, or the residue obtained by heating any cerium salt to red heat, is dissolved in concentrated sulphuric acid, and a crystal of strych- nine added, a deep blue color appears, which changes first to purple and then to red. Summary of analytical characters of the earth-metals and chromium. Aluminium,. Cerium. Chromium. Ammonium sulphide . White precipitate. White precipitate. Green precipitate. Potassium hydroxide . White precipitate. Soluble in KOH. Not re-precipitated by boiling. White precipitate. Insoluble in KOH. Green precipitate. Soluble in KOH. Be-precipitated by boiling. Ammonia water. . . White precipitate. White precipitate. Green precipitate. Ammonium carbonate White precipitate. White precipitate. Green precipitate. Questions.—231. Mention some varieties of crystallized aluminium oxide found in nature and some silicates containing it. 232. Give the general formula of an alum, and mention some alums. 233. Which alum is officinal, how is it made, what are its properties, and what is it used for? 234. What is dried alum, and how does it differ from common alum? 235. How is alu- minium chloride made, and how is the metal obtained from it? 236. State the properties of aluminium. 237. What change takes place when ammonium hydroxide, and what change when sodium carbonate is added to a solution of alum? 238. What is the composition of earthenware, porcelain, and glass; how and from what materials are they manufactured? 239. What is ultra- marine? 240. Give tests for aluminium compounds. 164 METALS AND THEIR COMBINATIONS. 25. IRON. (Ferrum.) Feii = 55.9 (55.88). General remarks regarding the metals of the iron group. The six metals (Fe, Co, Ni, Mu, Cr, Zn) belonging to this group are distin- guished by forming sulphides (chromium excepted) which are insoluble in water, but soluble in dilute mineral acids; they are, consequently, not precipitated from their neutral or acid solutions by hydrosul- phuric acid, but by ammonium sulphide as sulphides (chromium as hydroxide); their oxides, hydroxides, carbonates, phosphates, and sulphides are insoluble; their chlorides, iodides, bromides, sulphates, and nitrates are soluble in water. With the exception of zinc, these metals are magnetic; they decom- pose water at a red heat, the oxide being formed and hydrogen liber- ated ; in dilute hydrochloric or sulphuric acid, they dissolve with formation of chlorides or sulphates, respectively, and liberation of hydrogen. With the exception of zinc, which is bivalent, the metals of the iron group are bivalent in some compounds, trivalent in others, and form a number of oxides, the higher of which show, in some cases, decidedly acid properties, as, for instance, chromic or manganic oxides. The trivalence of the elements mentioned is now assumed to be due to the combining of two quadrivalent atoms of these elements. It is for this reason that we find in ferric, manganic, or chromic com- pounds always a double atom of these elements exerting a valence of six. The constitution of ferric chloride, Fe2Cl6, and ferric oxide, Fe203, may be graphically represented thus : /Cl Fe — Cl I XC1 I /Cl Fe_Cl \C1 F<° I >0 F< Occurrence in nature. Among all the heavy metals, iron is both the most useful and the most widely and abundantly diffused in nature. It is found, though usually in but small quantities, in nearly all forms of rock, clay, sand, and earth; its presence in these being indicated generally by their color (red, reddish-brown, or yellowish- red), as iron is the most common of all natural, inorganic coloring agents. It is found also, though in small quantities, in plants, and in somewhat larger proportions in the animal system, chiefly in the blood. In the metallic state iron is scarcely ever found, except in the IRON. 165 meteorites or metallic masses which fall occasionally upon our earth out of space. The chief compounds of iron found i n nature are : Hematite, ferric oxide, Fe203. Magnetic iron ore, ferrous-ferric oxide, Fe0.Fe203. Spathic iron ore, ferrous carbonate, FeCOa Iron pyrites, bisulphide of iron, FeS2. The carbonate and sulphate are found sometimes in spring waters, which, when containing considerable quantities of iron, are called chalybeate waters. Finally, iron is a constituent of some organic substances which are of importance in the animal system. Manufacture of iron. There is no other metal manufactured in such immense quantities as iron, the use of which in thousands of different tools, machines, and appliances is highly characteristic of our present age, Iron is manufactured from the above-named oxides or the carbonate by heating them with coke and limestone in large blast furnaces, which have a somewhat cylindrical shape, and are constantly fed from above with a mixture of the substances named, while hot air is forced into the furnace through suitable apertures near its hearth. The chemical change which takes place in the upper and less heated part of the furnace is a deoxidation of the iron oxide by the carbon : Fe203 + 3C = 3C0 + 2Fe. The heat necessary for this decomposition and fusion of the re- duced iron is produced by the combustion of the fuel, maintained by the oxygen of the air blown into the furnace. At the same time the lime and other bases combine with the silica contained in the ore, forming a fusible glass, called cinder or slag. The iron and slag col- lect at the bottom of the furnace, where they separate by gravity, and are run off every few hours. Iron thus obtained is known as cast-iron, or jpig-iron, and is not pure, but always contains, besides silicon (also sulphur, phosphorus, and various metals), a quantity of carbon varying from 2 to 5 per cent. It is the quantity of this carbon and its condition which im- parts to the different kinds of iron different properties. Steel contains from 0.16 to 2 per cent., wrought- or bar-iron but very small quanti- fies, of carbon. Wrought-iron is made from cast-iron by the process known as puddling, which is a burning-out of the carbon by oxida- tion, accomplished by agitating the molten mass in the presence of an oxidizing flame. Steel is made either from cast-iron by partially 166 METALS AND THEIR COMBINATIONS. removing the carbon, or from wrought-iron by recombining it with carbon—i. e., by agitating together molten wrought- and cast-iron in proper proportions. Properties. The high position which iron occupies among the useful metals is due to a combination of valuable properties not found in any other metal. Although possessing nearly twice as great a tenacity or strength as any of the other metals commonly used in the metallic state, it is yet one of the lightest, its specific gravity being about 7.7. Though being when cold the least yielding or malleable of the metals in common use, its ductility when heated is such that it admits of being rolled into the thinnest sheets and drawn into the finest wire, the strength of which is so great that a wire of one- tenth of an inch in diameter is capable of sustaining 700 pounds. Finally, iron is, with the exception of platinum, the least fusible of all the useful metals. Iron is little affected by dry air, but is readily acted upon by moist air, when ferric oxide and ferric hydroxide (rust) are formed. Iron forms two series of compounds, distinguished as ferrous and ferric compounds; in the former, iron is bivalent, in the latter, appa- rently trivalent, because, as shown above, the double atom exerts a valence of six. Almost all ferrous compounds show a tendency to pass into ferric compounds when exposed to the air, or more readily when treated with oxidizing agents, such as nitric acid, chlorine, etc. As the reactions of iron in ferrous and ferric compounds differ con- siderably, they must be studied separately. Reduced iron, Ferrum reductum. This is metallic iron, obtained as a very fine, grayish-black powder by passing hydrogen gas (puri- fied and dried by passing it through sulphuric acid) over ferric oxide, heated in a glass tube : Fe2Os + 6H = 3H20 + 2Fe. The officinal article should have at least 80 per cent, of metallic iron. Ferrous oxide, FeO (Monoxide or suboxide of iron). This com- pound is little kuown in the separate state, as it has (like most ferrous compounds) a great tendency to absorb oxygen from the air. The ferrous hydroxide, Fe(OH)2, may be obtained by the addition of any alkaline hydroxide to the solution of any ferrous salt, when a white precipitate is produced which rapidly turns bluish-green, dark-gray, black, and finally brown, in consequence of absorption of oxygen (see Plate I., 2): IRON. 167 FeS04 + 2NH4OH = (NH4)2S04 + Fe(OH)2; 2Fe(OH)2 + 0 + H20 = Fe2(OH)6. The precipitation of ferrous hydroxide is not complete, some iron always remaining in solution. Ferrous oxide is a stroug base, uuiting with acids to form salts, which have usually a pale green color. Ferric oxide, Fe203. A reddish-brown powder, which may be obtained by heating ferric hydroxide to expel water : Fe2(OH)6 = Fe203 -f- 3H20. It is a feeble base; its salts show usually a brown color. Ferric hydroxide, Ferric hydrate, Ferri oxidum hydratum, Fe2(OH)6 = 213.8 (Hydrated oxide of iron, Per- or sesqui-oxide, Ped oxide of iron), is obtained by precipitation of ferric sulphate or ferric chloride by ammonium or sodium hydroxide (see Plate I., 3) : Precipitation is complete, no iron remaining in solution as in the case of ferrous salts. Ferric hydroxide is a reddish-brown powder, sometimes used as an antidote in arsenic poisoning; for this purpose it is not used in the dry state, but after having been freshly precipitated aud washed, it is mixed with water, and this mixture used. Recently precipitated and consequently highly divided ferric hydroxide combines more readily with arsenious acid than the hydroxide which has been kept some time, or which has been dried, and thereby assumed a denser con- dition. Hydrated oxide of iron with magnesia, U. S. P., is a mixture made by adding magnesia to a solution of ferric sulphate, when magnesium sulphate and ferric hydroxide are formed ; the two substances are not separated from each other, the mixture being intended for immediate administration as an antidote in cases of arsenic poisoning. Fe2(S04)3 + 6NH4OH = 3[(NH4)2S04] + Te2(OH)6. Ferrous-ferric oxide, Fe0.Fe203 (.Magnetic oxide). This compound, which shows strong magnetic properties, has been mentioned above as one of the iron ores, and is known as loadstone. It has a metallic lustre and iron-black color, and is produced artificially by the com- bustion of iron in oxygen, or in the hydrated state by the addition of ammonium hydroxide to a mixture of solutions of ferrous and ferric salts. 168 METALS AND THEIR COMBINATIONS. Iron trioxide, Fe03. Not known in a separate state, but in com- bination with alkalies. In these compounds, called ferrates, Fe03 acts as an acid oxide, analogous to chromium trioxide, Cr03, in chromates. The composition of potassium ferrate is K2Fe04. Ferrous Chloride, FeCl2 (.Protochloride of iron), is obtained as a pale-green solution by dissolving iron in hydrochloric acid : Fe + 2HCI = FeCI2 + 2H. By evaporation of the solution, the dry salt may be obtained. The solution and salt absorb oxygen very readily : 3FeCl2 + O = FeO + Fe2CI6 Ferric chloride, ferrous, and afterward ferric oxide, are formed. Ferric chloride, Ferri chloridum, Fe2Cl6.12H20 = 540.2 (Chloride, sesqui-chloride, or perchloride of iron), is obtained by adding to the solution of ferrous chloride (obtained as mentioned above) hydro- chloric and nitric acids in sufficient quantities, and applying heat until complete oxidation has taken place. The nitric acid oxidizes the hydrogen of the hydrochloric acid to water, while the chlorine combines with the ferrous chloride, nitrogen dioxide being formed also: 6FeCl2 + 2HN03 + 6HC1 = 3Fe2Cl6 + 4H20 + 2N0. By sufficient evaporation of the solution, ferric chloride is obtained as a crystalline mass of an orange-yellow color; it is very deli- quescent, has an acid reaction, and a strongly styptic taste. The water of crystallization cannot be expelled by heat, because heat decomposes the salt, free hydrochloric acid and ferric oxide being formed. Experiment 26. Dissolve by the aid of heat 1 gramme of fine iron wire in about 4 c.c. of hydrochloric acid, previously diluted with 2 c.c. of water. Filter the warm solution of ferrous chloride, mix it with 2 c.c. of hydrochloric acid, and add to it slowly and gradually about 0.6 c.c. of nitric acid. Evap- orate in a fume chamber as long as red vapors escape; then test a few drops with potassium ferricyanide, which should not give a blue precipitate; if it does, the solution has to be heated with a little more nitric acid until the con- version into ferric chloride is complete and the potassium ferricyanide produces no precipitate. Ferric chloride thus obtained may be mixed with 4 c.c. of hot water and set aside, when it forms a solid mass of Fe2Cl6.12H20. How much FeCl2, how much Fe2Cl6, and how much Fe2Cl6.12H20 can be obtained from 1 gramme of iron ? IRON. 169 Solution of chloride of iron, Liquor ferri chloridi, U. S. P. This is a solution in water, containing 37.8 per cent, of the anhydrous ferric chloride. It is a reddish-brown liquid of specific gravity 1.405, having the taste aud reaction of the dry salt. This solution, mixed with about 2 parts of alcohol and left standing in a closed vessel for several months, forms the tincture of chloride of iron, Tinctura ferri chloridi, IT. S. P. Dialyzed iron is an aqueous solution of about 5 per cent, of ferric hydroxide with some ferric chloride. It is made by slowly adding to a solution of ferric chloride, ammonium hydroxide as long as the precipitate of ferric hydroxide formed is redissolved in the -ferric chloride solution, on shaking violently. The clear solution thus obtained is placed in a dialyzer floating in water, which latter is renewed every day until it shows no reaction with silver nitrate. The ammonium chloride passes through the membrane of the dialyzer into the water, while all iron as hydroxide with some chloride is left in solution. The combination of an oxide or hydroxide with a normal salt is called usually a basic salt or oxy-salt; dialyzed iron is a highly basic oxychloride of iron. Ferrous iodide, Fel2. When water is poured upon a mixture of metallic iron (fine wire is best) and iodine, the two elements combine directly, forming a pale-green solution of ferrous iodide, from which the salt may be obtained by evaporation. As it is oxidized and decomposed easily by the action of the air, an officinal preparation, the saccharated iodide of iron, U. S. P., is made by adding about 30 parts of sugar of milk to 20 parts of ferrous iodide; the sugar pre- vents, to some extent, rapid oxidation. Experiment 27. Cover some fine iron wire with water, heat gently, and add iodine in fragments as long as the red color of iodine disappears. Notice that the iron is dissolved gradually, the result of the reaction being the formation of a pale-green solution of ferrous iodide. Ferrous bromide, FeBr2. Made analogously to ferrous iodide, by the action of bromine on metallic iron. Ferrous suphide, FeS. Easily obtained as a black, brittle mass, by heating iron filings with sulphur, when the elements combine. It is used chiefly for liberating hydrogen sulphide, by the addition of sul- 170 METALS AND THEIR COMBINATIONS. phuric acid. Iron combines with sulphur in several proportions ; some of these iron sulphides are found in nature. Ferrous sulphate, Ferri sulphas, FeS04.7H20 = 277.9 [Sulphate of iron, Green vitriol, Copperas). Obtained by dissolving iron in dilute sulphuric acid, evaporating, and crystallizing : Fe + H2S04 = 2H + FeS04. Also obtained as a by-product in some branches of chemical indus- try, and by heap-roasting of the native iron sulphide : Ferrous sulphate crystallizes in large, bluish-green prisms; it is soluble in water, insoluble in alcohol. Exposed to the air, it loses water of crystallization, and absorbs oxygen. The dried ferrous sulphate, U. S. P., is made by expelling from 5 to 6 molecules of water by heating to 150° C. (302° F.); the granu- lated (precipitated) ferrous sulphate is made by pouring a strong aque- ous solution of ferrous sulphate, slightly acidulated with sulphuric acid, into alcohol, when ferrous sulphate separates as a crystalline powder, which is washed and dried. FeS2 + 60 = FeS04 + S02. Ferric sulphate, Fe2(S04)3. The solution of this salt, Liquor ferri tersulphatis, Solution of tersulphate of iron, U. S. P., is made by add- ing sulphuric and nitric acids to a solution of ferrous sulphate, and heating: 6FeS04 + 3H2S04 + 2HN03 = 3[Fe2(S04)3] + 2N0 + 4H20. The action of nitric acid is similar to that described above under ferric chloride. The hydrogen of the sulphuric acid is oxidized, and the radical S04 unites with the ferrous sulphate, nitrogen dioxide being liberated. Solution of ferric sulphate is used in the preparation of Ammonio- ferric sulphate, Fern et ammonii sulphas, (NH4)2S04.Fe2(S04)3.24H20 (iron alum or ammonio-ferric alum), which is made by mixing solu- tion of ferric sulphate with ammonium sulphate and crystallizing. Solution of subsulphate of iron, Liquor ferri subsulphatis (MonseTs solution). This is a solution similar to the preceding, but contains less sulphuric acid, and is, therefore, looked upon as a basic ferric sulphate, of the doubtful composition Fe405S04. The color of the tersulphate of iron solution is reddish-brown; that of MonsePs solu- tion is ruby-red. IRON. 171 Ferric nitrate, Fe2(N03)6. A 6 per cent, solution of this salt is officinal, under the name Solution of nitrate of iron, Liquor ferri nitratis, U. S. P., and is made by dissolving ferric hydroxide in nitric acid : Fe2(OH)6 + 6HN03 = 6H20 + Fe2(N03)6. It is an amber-colored, or reddish, acid liquid. Ferrous carbonate, FeC03. Occurs in nature; may be obtained by mixing solutions of ferrous sulphate and sodium carbonate or bicarbonate: FeS04 + Na2C03 = Na2S04 + FeC03. The precipitate is first nearly white, but soon assumes a gray color from oxidation. The saccharated carbonate of iron, U. S. P., is made by mixing the washed precipitate with sugar, and drying. The sugar prevents, to some extent, rapid oxidation. Ferric carbonate does not exist, the affinity between the feeble ferric oxide and the weak carbonic acid not being sufficient to unite them chemically. Ferrous phosphate, Fe3(P04)2. When sodium phosphate is added to solution of ferrous sulphate, a precipitate of the composition FeHP04 is formed : If, however, sodium acetate is added, a precipitate of the composition Fe3(P04)2 is formed: Na2HP04 + FeS04. = FeHP04 + Na2S04. 3FeS04 + 2Na2HP04 = Fe3(P04)2 + 2Na2S04 + H2S04. The sulphuric acid liberated, as shown in this formula, decomposes the sodium acetate, forming sodium sulphate and free acetic acid. Ferrous phosphate is a slate-colored powder, absorbing oxygen readily, becoming darker in color. Ferric phosphate, FeP04, may be obtained from ferric chloride solution by precipitation with an alkali phosphate. The soluble ferric phosphate of the U. S. P. is a scale compound. (See index.) Eerric hypophosphite, Eerri hypophosphis, Fe2(H2P02)6 = 501.8 (Hypophosphite of iron). Made by dissolving ferric hydroxide in hypophosphorous acid, and evaporating. It is a grayish-white powder, insoluble in water, soluble in hydrochloric acid. Questions.—241. Which metals belong to the “ iron group,” and what are their general properties? 242. How is iron found in nature, and what com- pounds are used in its manufacture ? 243. Describe the process for manufac- 172 METALS AND TREIR COMBINATIONS. Analytical reactions. 1. Ammonium sul- --'- V . . ■ •• •* Ferrous salts. (Use FeS04.) Black precipitate of ferrous Ferric salts. (Use Fe2Cl6.) Black precipitate of ferrous sul- phide. sulphide (Plate I., 1). phide mixed with sulphur. 2. Hydrosulphuric FeS04 -f- (NH4)2S = (NH4)2S04 + FeS. No change. Fe2Cl6 + 3[(NH4)2S] = 6NII4C1 + 2FeS + S. Ferric salts are converted into acid. 3. Ammonium, so- White precipitate of ferrous ferrous salts with precipita- tion of sulphur. , Fe2Cl6 + II2S = 2FeCl2 + 2IIC1 + S. Beddish-brown precipitate of dium, or potas- hydroxide soon turning ferric hydroxide. Precipita- sium hydroxide. green, black, and brown. tion is complete (Plate 1,3). 4. Ammonium, so- Precipitation not complete (Plate I., 2). FeCl2 + 2NaOH = 2NaCl + Fe(OH)2. White precipitate of ferrous Fe2Cl6 + 6(NH4OH) = 6NH4C1+ Fe2(OH)6. Beddish-brown precipitate of dium, or potas- carbonate, soon turning ferric hydroxide, with libera- sium carbonate. darker. tion of carbon dioxide (Plate 5. Alkali phos- FeCl2 + Na2C03 = 2NaCl + FeC03. Almost white precipitate, soon 1, 3). Fe2Cl6+3Na2C03 +3H20= 6NaCl + Fe2(0H)6+3C02. A yellowish-wliite precipitate phates or arseni- turning darker. is produced. ates. 6. Potassium ferro- Almost white precipitate, soon Dark-blue precipitate of ferric cyanide. turning blue by absorption ferrocyanide, or Prussian blue. K4Fe(CN)6. of oxygen (Plate I., 4). Decomposed by alkalies; in- 7. Potassium ferri- Blue precipitate of ferrous ferri- soluble in acids (Plate I., 5). 2Fe2Cl6 + 3[K4Fe(CN)6] = 12KC1 + Fe43[Fe(CN)6]. No precipitate is produced, but cyanide. cyanide, or Turnbull’s blue. the liquid is darkened to a K6Fe2(CN)12. 3FeCJ2 + K6Fe2(CN)12 = greenish brown hue. 8. Tannic acid. 6KC1 + Fe3Fe2(CN)I2 No change, provided oxidation A dark greenish-black precipi- 9. Potassium sul- of the ferrous salt has not taken place. As above. tate of ferric tannate is pro- duced (Plate I., 7). Deep blood-red precipitate of phocyanate. ICONS. ferric sulphocyanate (Plate I., 6). PLATE I. IRON. ZINC. 1 Ferrous sulphide, precipitated from ferrous or ferric solutions by am- monium sulphide. [Page 172.] Ferrous hydroxide passing into ferric hydroxide. Ferrous solutions precipitated by alkaline hydroxides. [ Pages 1(1(1, 172 ] 2 3 Ferric hydroxide, precipitated from ferric solutions by alkaline hy- droxides. [Pages 1(17, 172.] 4 Ferrous solutions, precipitated by potassium ferro-cyanide. [Page 172.] Ferric solutions, precipitated by potassium ferro-cyanide or Ferrous solutions precipitated by potassium ferri-cyanide. [Page 172.] 5 6 Ferric solutions, precipitated by alkaline 8ulpho-cyanat.es. [Page 172 ] Ferric solutions, precipitated by tannic acid. [Page 172.] 7 Zinc solutions, precipitated by either ammonium sulphide or by alka- line hydrates, carbonates, phosphates, ferro-cyanides. etc. [Page 182.] 8 MANGANESE — CHROMIUM—COBALT—NICKEL. 173 26. MANGANESE—CHROMIUM—COBALT—NICKEL. Manganese, Mn = 54.8. Manganese is found either as dioxide (Black oxide of manganese, pyrolusite), Mn02, or as sesquioxide, Mn2Oa. In small quantities it is a constituent of many minerals. Metallic manganese resembles iron in its physical and chemical properties, and may be obtained by reducing the carbonate with char- coal. Manganese is darker in color than iron, considerably harder, and somewhat more easily oxidized. Oxides of manganese. Four well-defined compounds of manganese with oxygen are known in the separate state, and two others only in combination with other elements. These oxides are : Manganous oxide (monoxide or protoxide), MnO. Manganous-manganic oxide, Mn0.Mn203 = Mn304. Manganic oxide (sesquioxide), Mn203. Manganese dioxide (binoxide, peroxide, black oxide), Mn02. Manganic acid, Permanganic acid,. Not known in a separate state, H20 -f- MnOs. H20 + Mn207. Manganous oxide is a greenish-gray powder, obtainable by heating the carbonate ; or as a nearly white hydroxide by precipitating a man- ganous salt by sodium hydroxide. It is a strong base, saturating acids completely, and forming salts which have generally a rose color or a pale reddish tint. Manganese dioxide, Mangani oxidum nigrum, Mn02 = 86.8. This is by far the most important compound of manganese, as it is largely used for generating chlorine : It a heavy, grayish-black, crystalline mineral, liberating oxygen when heated to redness : Mn02 + 4HC1 = MnCl2 + 2H20 + 2C1. 3Mn02 = Md304 + 20. turing iron on a large scale, and state the difference between cast-iron, wrought-iron, steel, and reduced iron. 244. State the composition and mode of preparation of ferrous and ferric hydroxides. What are their properties ? 245. Describe in words and chemical symbols the process for making ferric chloride. What is tincture of chloride of iron ? 246. How are ferrous iodide and bromide made? 247. State the properties of ferrous sulphate. Under what other names is it known, and how is it made? 248. What change takes place when soluble carbonates are added to soluble ferrous and ferric salts? 249. Mention agents by which ferrous compounds may be converted into ferric compounds, and these into ferrous compounds. Explain the chemical changes taking place. 250. Mention tests for ferrous and ferric compounds. 174 METALS AND THEIR COMBINATIONS. The officinal article should contain at least 66 per cent, of Mn02. Manganous sulphate, Mangani sulphas, MnS04 4H,0 = 222.8, may be obtained by dissolving the oxide or dioxide in sulphuric acid ; in the latter case oxygen is evolved : As manganese dioxide generally contains iron oxide, the solution contains sulphates of both metals. By evaporating to dryness and strongly igniting, the iron salt is decomposed. The ignited mass is now lixiviated with water, and the filtered solution evaporated for crystal- lization. It is an almost colorless, or pale rose-colored substance, isomor- plious with the sulphates of magnesium and zinc; it is easily soluble in water. JVIn02 -|- H2S04 = MnS04 -(- H20 -|- O. Potassium permanganate, Potassii permanganas, KMn04 = 157.8 [Permanganate of potassium). Whenever a compound (any oxide or salt) of manganese is fused with alkali carbonates (or hydroxides) and alkali nitrates (or chlorates) the manganese is converted into manganic acid, which combines with the alkali, forming potassium (or sodium) manganate : 3Mn02 + 3K2C03 + KC10S = 3K2Mn04 + 8C()2 + KC1. The fused mass has a dark-green color, and when dissolved in water gives a dark emerald-green solution, from which, by evaporation, green crystals of potassium manganate may be obtained. The green solution is decomposed easily by any acid (or even by water in large quantity) into a red solution of potassium permanganate and a precipitate of manganese dioxide. 3K2Mn04 + 2H2S04 = Mn02 + 2K2S04 + 2KMn04 + 2H20. By evaporation and crystallization potassium permanganate is ob- tained in slender, prismatic crystals, of a deep purple-violet color, and a somewhat metallic lustre. The solution in water lias a deep purple, or, when highly diluted, a pink color (Plate II., 1). It is a powerful oxidizing agent, and an excellent disinfectant, both properties being due to the facility with which a portion of the oxygen is given off to any substance which has affinity for it. If the oxidation takes place in the an acid, a lower oxide of manganese is formed, which separatea $s. an insoluble substance. If an acid is present, both the potassium and manganese combine with it, forming salts, thus: MANGANESE — CHROMIUM — COBALT — NICKEL. 175 2(KMn04) + 6HC1 + a; = 2KC1 + 2MnCl2 + 3H20 + xOs. x represents here any substance capable of combining with oxygen while in solution. Experiment 28. Heat in a porcelain crucible a mixture of 2 grammes man- ganese dioxide, 2 grammes potassium hydroxide, and 1 gramme potassium chlorate, until the fused mass has turned dark-green. Dissolve the cooled mass with water, filter the green solution of potassium manganate, and pass carbon dioxide through it until it has assumed a purple color, showing that the con- version into permanganate is complete. Notice that the acidified solution is readily decolorized by ferrous salts and other deoxidizing agents. Analytical reactions. 1. Ammonium sulphide produces a yellowish-pink or flesh-colored precipitate of hydrated manganous sulphide, MnS.H20, soluble in acetic and in mineral acids (Plate II., 2): (Manganous sulphate, MnS04, may be used.) 2. Ammonium (or sodium) hydroxide produces a white precipitate of manganous hydroxide, which soon darkens by absorption of oxygen (Plate II., 3) and dissolves in oxalic acid with a rose-red color. MnSO, + (NH4)2S = (NH4)2S04 + MnS. MnCI2 + 2NH4OH = 2NH4C1 + Mn(OH)2. 3. Sodium (or potassium) carbonate produces a nearly white pre- cipitate of manganous carbonate: MnS04 + Na20O3 = JSa2S04 + MnC03. 4. Any compound of manganese heated on platinum foil with a mixture of sodium carbonate and nitrate forms a bluish-green mass, giving a green solution in water, which turns red on addition of an acid. (See explanation above.) 5. Manganese compounds fused with borax on a platinum wire give a violet color to the borax bead. 6. Traces of manganese may be detected by boiling with dilute nitric acid and red lead, when the solution acquires a reddish-purple color due to the formation of permanganic acid. Chromium, Cr = 52. Found in nature almost exclusively as chro- mite, or chrome-iron ore, FeO.Cr2Oa, a mineral analogous in compo- sition to magnetic iron ore, FeO.Fe2Oa. The name chromium, from the Greek xp&pa (chroma), color, was given to this metal on account of the beautiful colors of its different compounds, none of which is 176 METALS AND THEIR COMBINATIONS. colorless. Chromium forms two basic oxides, CrO and Cr203, and an acid oxide, Cr()3, the combinations and reactions of which have to be studied separately. While chromium is closely allied to aluminium and iron on one side, it also shows a resemblance to sulphur, as indi- cated by the trioxide, Cr03, and the acid, H2Cr04, which are analogous to S03 and H2S04. Moreover, the barium and lead salts of chromic and sulphuric acids are both insoluble. Potassium dichromate, Potassii bichromas, K2Cr207 = 294 (Bichro- mate or red chromate of potassium). This salt is by far the most im- portant of all combinations of chromium, and is the source from which they are obtained. Potassium dichromate is manufactured on a large scale by exposing a mixture of the finely ground chrome-iron ore with potassium car- bonate and calcium hydroxide to the heat of an oxidizing flame in a reverberatory furnace, when both constituents of the ore become oxid- ized, ferric oxide and chromic acid being formed, the latter combining with the potassium, forming normal potassium chromate, K2Cr04: By treating the heated mass with water a yellow solution of potas- sium chromate is obtained, which, upon the addition of sulphuric acid, is decomposed into potassium dichromate and potassium sulphate: 2(Fe0Cr203) + 4K2C03 + 70 = Fe203 + 4C02 + 4(K2Cr04). The two salts may be separated by crystallization. Potassium di- chromate forms large, orange-red, transparent crystals, which are easily soluble in water; heated by itself oxygen is evolved, heated with hydrochloric acid chlorine is liberated, heated with organic matter or reducing agents these are oxidized. To explain the constitution of dichromates we have to assume that chromic anhydride, Cr03, is capable of forming two acids : 2(K2Cr04) + H2S04 = K2Cr207 + K2S04 + II20. Cr03 + H20 = H2Cr 04 = Chromic acid. 2CrOs -f- H20 = H2Cr207 = Dichromic acid. Chromium trioxide, Aeidum Chromicum, Cr03 = 100 (Chromic acid, Chromic anhydride), is prepared by adding sulphuric acid to a satu- rated solution of potassium dichromate, when chromium trioxide separates in crystals : K2Cr207 + H2S04 = K2S04 + H20 + 2Cr03. Thus prepared, it forms deep crimson-red, needle-shaped crystals, which are deliquescent, and very soluble in water; it is powerfully MANGANESE—CHROMIUM — COBALT NICKEL. 177 corrosive, and one of the strongest oxidizing agents; the solution in water has strong acid properties; it combines with metallic oxides, forming chromates and dichromates. Experiment 29. Dissolve a few grammes of potassium dichromate in water and add to 4 volumes of the cold saturated solution 5 volumes of strong sulphuric acid ; chromium trioxide separates on cooling. Collect the crystals on asbes- tos, wash them with a little nitric acid, and dry them by passing warm dry air through a tube in which they have been placed for this purpose. Chromic oxide, Cr203 (Sesquioxide of,chromium), is obtained by heat- ing potassium dichromate with sulphur, when potassium sulphate and chromic oxide are formed : K2Cr207 + S = K2S04 + Cr203. By washing the heated mass with water, the chromic oxide is left as a green powder, which is insoluble in water and in acids; it is a basic oxide combining with acids to form salts; it is used as a green color, especially in the manufacture of painted glass and porcelain. Chromic hydroxide, Cr2(OH)6. A solution of potassium dichromate may be deoxidized by the action of hydrosulpliuric acid, sulphurous acid, alcohol, or any other deoxidizing agent, in the presence of sul- phuric or hydrochloric acid : K2Cr207 + 4H2S04 + 3H2S = K2S04 + 7H20 + 3S + Cr2(S04)3. As shown by this formula, the sulphates of potassium and chro- mium are formed and remain in solution, while sulphur is precipi- tated, the hydrogen of the hydrosulphuric acid having been oxidized and couverted into water. By adding ammonium hydroxide to the solution thus obtained, chromic hydroxide is precipitated as a bluisli-green gelatinous sub- stance : Cr2(SOi)3 + 6NH4OH = 8[(NH4)2S04] + Cr2(OH)6. By dissolving this hydroxide in the different acids, the various salts, such as chloride, Cr2Cl6, sulphate, etc., are obtained. Chromic sulphate, similar to aluminium sulphate, combines with potassium or ammonium sulphate and water, forming chrome alum, K2S04.Cr2 (S04)324H20 ; it is a purple salt, and is isomorphous with other alums. 178 METALS AND THEIR COMBINATIONS. Analytical reactions. a. Of chromic acid or chromates. (Use potassium chromate, K2Cr04.) 1. Hydrogen sulphide added to an acidified solution of a chromate, changes the red color into green with precipitation of sulphur. The solution now contains chromium in the basic form. (See explanation above.) (Plate II., 4.) The conversion of chromic acid into oxide is more readily accomplished by heating the chromic solution with alco- hol and hydrochloric acid ; the alcohol is partly oxidized, being con- verted into aldehyde. 2. Soluble lead salts produce a yellow precipitate of lead chromate (chrome yellow), PbCr04, insoluble in acetic, soluble in hydrochloric acid and in sodium hydroxide (Plate II., 6) : K2Cr04 + Pb(N03)2 = PbCr04 + 2KN03. 3. Barium chloride produces a pale-yellow precipitate of barium chromate, BaCr04: 4. Silver nitrate produces a dark-red precipitate of silver chromate, Ag2Cr04 (Plate II., 8): K2Cr04 + BaCl2 = BaCrO, + 2KC1. 2AgN03 + K2Cr04 = 2KNOs + Ag2Cr04. 5. Mercurous nitrate produces a red precipitate of mercurous chro- mate, Hg2Cr04. b. Of salts of chromium. (Use chrome-alum or chromic chloride, Cr2Cl6.) 6. To chromic chloride or sulphate add ammonium hydroxide or ammonium sulphide : in both cases the green hydroxide of chromium, Cr2(OH)6, is precipitated (Plate II., 5): 7. Potassium or sodium hydroxide causes a similar green precipi- tate of chromic hydroxide, which is soluble in an excess of the reagent, but is re-precipitated on boiling for a few minutes. Cr2C)6 + 3[(NH4)2S] + 6H20 = 6NH4C1 + 3H2S + Cr2(OH)6. c. Of chromium in any form. 8. Compounds of chromium, when mixed with sodium (or potas- sium) carbonate and nitrate, give, when heated upon platinum foil, a yellow mass of the alkali chromate. PLATE XI. MANGANESE. CHROMIUM. 1 Potassiu m permanganate solu- tion, more or less saturated, Borax- head colored by Manganese. [Pages 174, 175 ] 2 Manganous sulphide precipi- tated from manganous solutions hy ammonium sulphide. [Page 175.] 3 Manganous hydroxide passing into the higher oxides. Manganous solutions precipitated by alkaline hy- droxides. [Page 175.] 4 Potassium dichromate solution deoxidized by reducing agents. [Page 178.] 5 Chromic hydroxide precipitated from chromic solutions by alkaline hydrates or by ammonium sulphide. [ Pages 177, 178 ] 6 Lead chromate precipitated from soluble chromates by lead ace- tate. [Pages 178.186.j 7 Silver chromate precipitated from neutral chromates hy silver ni- trate. f Pages 178, 195.] 8 Mercurous chromate precipi- tated from neutral chromates by mer- curous solutions. [Pagr«178.] ZINC. 179 9. Compounds of chromium impart a green color to the borax bead. Cobalt and Nickel, Co = 58.7, Ni = 58.6. These two metals show much resemblance to each other in their chemical and physical properties, and occur in nature associated with each other as sulphides or arsenides. Both metals are nearly silver-white ; the salts of cobalt show generally a red, those of nickel a green color. The solutions of both metals give a black pre- cipitate of the respective sulphides on the addition of ammonium sulphide. Ammonium hydroxide produces in solutions of cobalt a blue, in solutions of nickel a green precipitate of the hydroxides, both of which are soluble in an excess of the reagent; potassium or sodium hydroxide produces similar pre- cipitates, which are insoluble in an excess. Cobalt is used chiefly when in a state of combination (for coloring glass blue); nickel when in the metallic state. (German silver is an alloy of nickel, copper, and zinc.) 27. ZINC. Znii = 65.1. Occurrence in nature. Zinc is found chiefly either as sulphide (zinc-blende), ZnS, or as carbonate (calamine), ZnCOs; also it occurs in combination with silicic acid as silicate and with oxygen as the red oxide. Metallic Zinc is obtained by heating in retorts the oxide or carbon- ate mixed with charcoal, when decomposition takes place. The liber- ated metal is vaporized, and distils into suitable receivers, where it solidifies. Zinc is a bluish-white metal, which slowly tarnishes in the air, becoming coated with a film of oxide and carbonate; it has a crys- talline structure and is, under ordinary circumstances, brittle; when heated to about 130°-150° C. (260°-302° F.) it is malleable, and may be rolled or hammered without fracture. Zinc thus treated retains this malleability when cold; the sheet-zinc of commerce is thus made. When zinc is further heated to about 200° C. (392° F.), Questions.—251. How is manganese found in nature ? 252. Mention the different oxides of manganese. What is the binoxide used for? 253. What is the color of manganese salts, of manganates, and of permanganates ? 254. How is potassium permanganate made; what are its properties, and what is it used for? 255. Give tests for manganese. 256. State composition and prop- erties of potassium dichromate. 257. How is chromium trioxide made; what are its properties; what is it used for; and under what other name is it known ? 258. By what process may chromium sesquioxide be converted into chromates ? 259. What is the composition of the oxide and hydroxide of chromium, and how are they made? 260. Mention tests for chromates and chromium salts. 180 METALS AND THEIR COMBINATIONS. it loses its malleability and becomes so brittle that it may be pow- dered ; at 410° C. (760° F.) it fuses, and at a bright-red heat it boils, volatilizes, and, if air be not excluded, burns with a splendid greenish- white light, generating the oxide. Zinc is used by itself in the metallic state or fused together with other metals (German silver and brass contain it); galvanized iron is iron coated with metallic zinc. Zinc is a bivalent metal, forming but one oxide and one series of salts, all of which have a white color. Zinc oxide, Zinci oxidum, ZnO — 81.1 (Oxide of zinc, Flores zinci, Zinc-white), may be obtained by burning the metal, but if made for medicinal purposes by heating the carbonate, when carbon dioxide and water escape and the oxide is left: It is a soft, pale-yellowish or nearly white, tasteless powder, insoluble in water, soluble in acids; when strongly heated it turns yellow, but on cooling resumes the white color. Zinc hydroxide, Zn20H, is obtained by precipitating zinc salts with the hydroxide of sodium or ammonium ; the precipitate, however, is soluble in an excess of either of the alkali hydroxides. 3[Zn(0H)2].2ZnC03 = 5ZnO + 2C02 + 3H20. Zinc chloride, Zinci chloridum, ZnCl2 = 135 9 (Chloride oj zinc). Made by dissolving zinc or zinc carbonate in hydrochloric acid aud evaporating the solution to dryness : Zn + 2HC1 = ZnCl2 + 2H. It is met with either as a white, crystalline powder, or in white opaque pieces ; it is very deliquescent and easily soluble in water and alcohol; it combines readily with albuminoid substances; it fuses at about 115°C. (239° F.), and is volatilized, with partial decomposition, at a higher temperature. Zinc bromide, Zinci bromidum, ZnBr2 = 224.7 (.Bromide of zinc). Obtained analogously to the chloride by dissolving zinc in hydrobromic acid : Zn + 2HBr = ZnBr2 + 2H. A white powder, resembling the chloride in its properties. Zinc iodide, Zinci iodidum, Znl2 = 318.1 (Iodide of zinc). The two elements zinc and iodine combine readily when heated with water; ZINC. 181 the colorless solution when evaporated to dryness yields a powder whose physical properties resemble those of the chloride. Zinc carbonate, Zinci carbonas prsecipitatus, 2(ZnC03).3[Zn(0H2)] = 547.5 (.Precipitated Carbonate of zinc). Solutions of equal quantities of zinc sulphate and sodium carbonate are mixed and boiled, when a white precipitate is formed, which is a mixture of the carbonate and hydroxide of zinc, corresponding more or less to the formula given above. "5ZnS04 + 5Nh2C03 + 3H20 = 3C02 + 5Na2S04 + 2(ZnC03).3[Zn(OH)2]. Precipitated zinc carbonate is a white, impalpable powder, odorless and tasteless, insoluble in water, soluble in acids. Zinc sulphate, Zinci sulphas, ZnS04.7H20 = 287.1 [Sulphate of zinc, White vitriol), is obtained by dissolving zinc in dilute sulphuric acid : H2S04 + *H20 + Zn = ZnS04 + a;H20 + 2H. If zinc be added to strong sulphuric acid, no decomposition takes place : no sufficient explanation has as yet been given for this fact. Zinc sulphate forms small, colorless crystals, which are isomorphous with magnesium sulphate; it is easily soluble in water. Experiment 30. Use the liquid obtained when performing Experiment 2, or, if not left, dissolve a few grammes of metallic zinc in dilute sulphuric acid, filter the solution, evaporate sufficiently, and set aside for crystallization. Use the zinc sulphate thus obtained for the analytical reactions. State the quantity of dilute sulphuric acid required for dissolving 5 grammes of zinc, and the quantity of crystallized zinc sulphate which may be obtained. Zinc phosphide, Zinci phosphidum, Zn3P2 = 257.3 (Phosphide of zinc). The two elements zinc and phosphorus combine readily when the latter is thrown upon melted zinc, forming a grayish-black powder, or minutely crystalliue, friable fragments, having a metallic lustre on the fractured surface. Antidotes. Soluble zinc salts (sulphate, chloride) have a poisonous effect. If the poison have not produced vomiting, this should be induced. Milk, white of egg, or, still better, some substance containing tannic acid (with which zinc forms an insoluble compound) should be given. 182 METALS AND THEIR COMBINATIONS. Analytical reactions. 1. Add to solution of a zinc salt ammonium sulphide: a white precipitate of zinc sulphide, ZnS, is produced. (Zinc sulphide is the only white insoluble sulphide.) (Zinc sulphate, ZnS04, may be used.) ZnS04 + (NH4)2S = (NII4)2S04 + ZnS. 2. From neutral zinc solutions, or from solutions containing free acetic acid, hydrogen sulphide precipitates white zinc sulphide. 3. Add ammonium, sodium, or potassium hydroxide : a white pre- cipitate of zinc hydroxide, Zn(OH)2, is produced, soluble in excess of the reagent, with the formation of zincates, such as Zn (OK)2. 4. Soluble carbonates and phosphates give white precipitates in neutral solutions of zinc. 5. Potassium ferrocyanide gives a white precipitate of zinc ferro- cyanide. (This test may be used to distinguish compounds of zinc from those of magnesium or aluminium.) 6. Zinc is the only heavy metal whose compounds are all colorless. The oxide, carbonate, phosphate, and ferrocyanide are insoluble; the chloride, nitrate, and sulphate soluble. Cadmium, Cd = 111.5. Found in nature associated (though in very small quantities) with the various ores of zinc, with which metal is has in common a number of physical and chemical properties. Cadmium ditfers from zinc by forming a yellow sulphide (with hydrosulphuric acid), soluble in diluted acids. Cadmium and its compounds are of little interest here; the yellow sulphide is used as a pigment, the sulphate and iodide sometimes for medicinal purposes. Questions.—261. How is zinc found in nature, and by what process is it obtained? 262. Mention the properties of metallic zinc, and what is it used for? 263. Mention two processes for making zinc oxide. 264. How does heat act on zinc oxide ? 265. Show by chemical symbols the action of hydrochloric and sulphuric acids on zinc. 266. State the properties of chloride and of sulphate of zinc. 267. What is white vitriol ? 268. Explain the formation of precipitated zinc carbonate, and state its composition. 269. Mention tests for zinc compounds. 270. How many pounds of crystallized zinc sulphate may be obtained from 21.7 pounds of metallic zinc? ZINC. Ferrous salts. Ferric salts. Manganese. Zinc. Cobalt. Nickel. Ammonium sulphide Black precipi- tate. Black precipi- tate. Flesh-colored precipitate. White precipi- tate. Black precipi- tate. Black precipitate. Ammonia water Dirty green pre- cipitate. Eeddish-brown precipitate. White precipi- tate. White precipi- tate. Blue precipitate. Green precipitate. In excess of reagent Insoluble. Insoluble. Soluble. Soluble Soluble. Soluble. Sodium hydroxide . Dirty green pre- cipitate. Eeddish-brown precipitate. White precipi- tate. White precipi- tate. Blue precipitate. Green precipitate. In excess of reagent Insoluble. Insoluble. Insoluble. Soluble. Insoluble. Insoluble. Sodium carbonate White precipi- tate, turning dark. Eeddish-brown precipitate. White precipi- tate, turning darker. White precipi- tate. Blue precipitate. Green precipitate. Ammonium carbonate White precipi- tate, darkens. Insoluble Eeddish-brown precipitate. White precipi- tate, darkens. White precipi- tate Blue precipitate. Green precipitate. In excess of reagent Insoluble. Insoluble. Soluble. Soluble. Soluble. Potassium ferrocyanide Pale blue pre- Dark blue pre- White precipi- White precipi- Grayish-green Greenish-white cipitate turn- ing darker. cipitate. tate tate. precipitate. precipitate. Potassium ferricyanide Dark blue pre- No precipitate, Pale brown Pale brownish- Deep brown-red Y ellowish- brown Potassium sulphocyanate . Borax bead in oxidizing flame cipitate. greenish-brown color. Dark red color. Dark yellow to red, while hot. precipitate. Yiolet. yellow precipi- tate. precipitate. Color much in- tensified. Blue. precipitate. Green color, slightly intensi- fied. Eed, while hot. Summary of analytical characters of metals of the iron group. 184 METALS AND THEIR COMBINATIONS. 28. LEAD-COPPER—BISMUTH. General remarks regarding the metals of the lead group. The six metals belonging to this group (Pb, Cu, Bi, Ag, Hg, and Cd) are dis- tinguished by forming sulphides which are insoluble in water, insoluble in dilute mineral acids, insoluble in ammonium sulphide ; consequently they are precipitated from neutral, alkaline, or acid solutions by hydrogen sulphide or ammonium sulphide. The metals themselves do not decompose water at any temperature, and are not acted upon by dilute sulphuric acid ; heated with strong sulphuric acid, most of these metals are converted into sulphates with liberation of sulphur dioxide; nitric acid converts all of them into nitrates with liberation of nitrogen dioxide. The oxides, iodides, sulphides, carbonates, phosphates, and a few of the chlorides and sulphates of these metals are insoluble; all the nitrates, and most of the chlorides and sulphates are soluble. In regard to valence, they show no uniformity whatever, silver being univalent, copper, cadmium, and mercury bivalent, bismuth trivalent, and lead either bivalent or quadrivalent. Lead, Pbu = 206.4 {Plumbum). This metal is obtained almost exclusively from the native sulphide of lead, called galena, PbS, by roasting until it is converted into oxide, and smelting this with coke in a blast furnace. Lead owes its usefulness in the metallic state chiefly to its softness, fusibility, aud resistance to acids, which properties are of advantage in using it for tubes or pipes, or in constructing vessels to hold or manu- facture sulphuric acid. Lead is a constituent of many alloys, as, for instance, of type-metal, solder, britanuia metal, shot, etc. Experiment 31. Dissolve 1 gramme of lead acetate or lead nitrate in about 200 c.c. of water, suspend in the centre of the solution a piece of metallic zinc and set aside. Notice that metallic lead is deposited slowly upon the zinc in a crystalline condition, whilst zinc passes into solution, which may be verified by analytical methods. The chemical change taking place is this: Pb(N03)2 + Zn = Zn(N08)2 + Pb. The formation of the crystallized lead is called generally a lead-tree. Lead oxide, Plumbi oxidum, PbO = 222.4 (Oxide of Lead, litharge). Obtained by exposing melted lead to a current of air, when the metal is gradually oxidized with the formation of a yellow powder, known LEAD—COPPER—BISMUTH. 185 as massicot; at a higher temperature this fuses, forming reddish-yellow crystalline scales, known as litharge ; by heating still further in contact with air, a portion of the oxide is converted into dioxide (or peroxide), Pb02, and a red powder is formed, known as red lead (or minium), which probably is a mixture (or combination) of oxide and peroxide of lead, Pb022Pb0. Oxide of lead is used in the manufacture of lead salts, lead plaster, glass, paints, etc. Nitric acid when heated with red lead combines with the oxide, while the dioxide is left as a dark-brown powder, which, on heating with hydrochloric acid, evolves chlorine (similar to manganese dioxide). Lead nitrate, Plumbi nitras, Pb(N03)2 = 330.4 [Nitrate of lead). Obtained by dissolving the oxide in nitric acid : PbO + 2IINO3 = II20 + Pb(N03)2 Lead nitrate is the only salt of lead (with a mineral acid) which is easily soluble in water; it has a white color, and a sweetish, astringent, and afterward metallic taste. Lead carbonate, Plumbi carbonas, 2(PbC03).Pb(0H)2 — 773.2 (Car- bonate of lead, White-lead). This compound may be obtained by precipitation of lead nitrate with sodium carbonate, but is manu- factured on a large scale directly from lead, by exposing it to the simultaneous action of air, carbon dioxide, and vapors of acetic acid. The latter combines with the lead, forming a basic acetate, which is converted into the carbonate (almost as soon as produced) by the carbon dioxide present. The action of acetic acid on lead or lead oxide will be considered in connection with acetic acid. Lead carbonate is a heavy, white, insoluble, tasteless powder; the white-lead of commerce frequently is found adulterated with barium sulphate. Lead iodide, Plumbi iodidnm, Pbl2 = 459.4 (Iodide of lead). Made by adding solution of potassium iodide to lead nitrate (Plate III., 6) : Pb2N03 + 2KI = 2KN03 + Pbl2. It is a heavy, bright citron-yellow, almost insoluble powder, which may be distinguished from lead chromate by its solubility in ammo- nium chloride solution on boiling, lead chromate being insoluble in this solution. 186 METALS AND THEIR COMBINATIONS. Poisonous properties and antidotes. Compounds of lead are directly poisonous, and it happens, not infrequently, that water passing through leaden pipes or collected in leaden tanks becomes contaminated with lead. Water free from air and salts scarcely acts on lead; but if it contain air, oxide of lead is formed, which is either dissolved by the water or is decomposed by the nitrates or chlorides present in the water, the soluble nitrate or chloride of lead being formed. If the water contains carbonates and sulphates, however, these will form insoluble compounds, producing a film or coating over the lead, preventing further contact with the water. Rain water, in consequence of its containing atmospheric constituents, and no sulphates, acts as a solvent on lead pipe; spring and river waters generally do not. Water containing lead will show a dark color on passing hydrogen sulphide through it; if the quantity present be very small, the water should be evaporated to A or even of its original volume before applying the test. The constant handling of lead compounds is one of the causes of lead poisoning (painters’ colic). As an antidote, magnesium sulphate should be used, which forms with lead an insoluble sulphate; the purgative action of magnesia is also useful. (In lead works workmen often drink water containing a little sulphuric acid.) Analytical reactions. (Lead acetate or lead nitrate, Pb(N03)2 may be used.) 1. To a solution of a lead salt add hydrogen sulphide or ammo- nium sulphide: a black precipitate of lead sulphide is produced (Plate III., 1): Pb(N03)2 + H2S = 2IiN03 + PbS. 2. Add sulphuric acid or soluble sulphate: a white precipitate of lead sulphate is formed : Pb(N03)2 + Nh2S04 = 2NaN03 + PbS04. 3. Add hydrochloric acid or a soluble chloride : a white precipitate of lead chloride, PbCl2, is produced, which dissolves on heating or on the addition of much water, as lead chloride is not entirely soluble. For the same reason, the precipitate is not formed when the solutions used are highly dilute. 4. Other reagents which give precipitates with lead solutions are : Potassium chromate, producing yellow lead chromate (chrome yellow). (Plate II., 6.) Potassium iodide, producing yellow lead iodide. (Plate III., 6.) Alkali carbonates, producing white basic lead carbonate. Alkali phosphates, producing white lead phosphate. Copper, Cu11 = 63.2 (Cuprum). Found in nature sometimes in the metallic state—generally, however, combined with sulphur or oxygen. LEA D—COPPER—BISMUTH. 187 The commonest copper-ore is Copper pyrites, a double sulphide of copper and iron, CuFeS2 or Cu2S.Fe2S3, having the color and lustre of brass or gold. Other ores are: Copper glance, cuprous sulphide., having a dark-gray color and the composition Cu2S; malachite, a beautiful green mineral, being a carbonate and hydroxide of copper, CuC03.Cn(0H)2. Cuprous and cupric oxide also are found occasion- ally. Copper is obtained from the oxide by reducing it with coke; sulphides previously are converted into oxide by roasting. Copper is the only metal showing a distinct red color; it is so malleable that, of the metals in common use, only gold and silver surpass it in that respect; it is one of the best conductors of heat and electricity, it does not change in dry air, but becomes covered with a film of green subcarbonate when exposed to moist air. Copper frequently is used in the manufacture of alloys, of which the more important are : Copper. Zinc. Tin. NicTcel. Brass . . 64 36 / ••• German silver . 51 31 18 Bell-metal . . 78 22 Bronze . 80 4 16 Gun-metal . . 90 10 Copper frequently is alloyed with gold and silver. Copper is a bivalent element, forming two oxides and two series of salts, distinguished as cuprous and cupric compounds ; the cuprous salts are here of but little interest. Cupric oxide, CuO (Oxide or monoxide of copper). Heated to red- ness, copper becomes covered with a black scale, which is cupric oxide; it is obtained also by heating cupric nitrate or carbonate, both com- pounds being decomposed with formation of the oxide; finally, it may be made by adding sodium or potassium hydroxide to the solution of a cupric salt, when a bulky, pale-blue precipitate of cupric hydroxide, Cu(OH)2, is formed, which, upon boiling, is decomposed into water and cupric oxide, a heavy dark-brown powder (Plate III., 2): CuS04 + 2K0H = K2S04 + Cu(OH)„; Cuprous oxide, Cu20 (Red oxide or suboxide of copper). When cupric oxide is heated with metallic copper, charcoal, or organic matter, the cupric oxide is decomposed, and cuprous oxide is formed. (Excess of carbon or organic matter reduces the oxide to metallic copper.) Cu(OH)2 = H20 + CuO. CuO -(- Cu = Cu20 ; 2CuO + 0 = Cu20 + CO. 188 METALS AND THEIR COMBINATIONS. Some organic substances, especially grape-sugar, decompose strong alkaline solutions of cupric sulphate with precipitation of cuprous oxide, which is a red, insoluble powder. Cupric sulphate, Cupri sulphas, CuS04.5H20 = 249.2 (Sulphate of copper, Blue vitriol, Blue-stone). This is the most important compound of copper. It is manufactured on a large scale, either from copper pyrites, or by dissolving cupric oxide in sulphuric acid, evaporating and crystallizing the solution : CuO + H2S04 = CuS04 + h2o. Cupric sulphate forms large, deep-blue crystals, which are easily soluble in water, and have a nauseous, metallic taste. By heating it to about 230° C. (446° F.) all water of crystallization is expelled, and the anhydrous cupric sulphate formed, which is a nearly white powder. By further heating this is decomposed, sulphuric and sulphurous oxides are evolved, and cupric oxide is left. Experiment 32. Boil about 5 grammes of fine copper wire with 15 c.c. of concentrated sulphuric acid until the action ceases and most of the copper is dissolved. Dilute with about 15 c.c. of hot water, filter, and set aside for crystallization. State the exact quantities of copper and H2S04 required to make 100 pounds of crystallized cupric sulphate. Cupric carbonate is obtained by the addition of sodium carbonate to solution of cupric sulphate, when a bluish-green precipitate is formed, which is cupric carbonate with hydroxide (Plate III., 4); by dissolv- ing this in the various acids, the different cupric salts are obtained. Ammonio-copper compounds. A number of compounds are known which are either double salts of ammonia and copper, or are derived from ammonium salts and contain copper. Thus, cupric chloride forms with ammonia the compounds: CuC122NH3, CuCl24NH3, and CuC126NH3. Cupric sulphate forms in like manner, cupri-diammo- nium sulphate, CuS042NH3, or (N2H6Cu)S04, which is a deep azure- blue compound capable of forming large crystals (Plate III., 3). It is this formation of soluble ammonio-copper compounds which prevents ammonium hydroxide from precipitating cupric hydroxide from its salts. Poisonous properties and antidotes. The use of copper for culinary vessels is frequently the cause of poisoning by this metal. A perfectly clean surface of metallic copper is not affected by any of the substances used in the preparation of food, but as the metal is very apt to become covered with a film XII. COPPER. LEAD. BISMUTH. Cupric sulphide or lead sul- phide precipitated from solutions of copper or lead by hydrogen sulphide. [Pages 186,189.] 1 Cupric hydroxide passing into cupric oxide. Cupric solutions precipitated hy potassium hydroxide and boiling. \_Pages 187, 189. 2 3 Amtnonio-cupric compounds obtained by adding ammonium hy- droxide to cupric solutions. [Page 188.] Cupric carbonate precipitated from cupric solutions by sodium car- bonate. | Page 189. j 4 5 Cupric ferro-cynnide precipi tated from cupric solutions by potas sium ferro-cyanide. [Page 189.] Lead, iodide precipitated from lead solutions by soluble iodides. [Pages 185,186.] 6 7 Lend .solutions with soluble chlorides, sulphates or carbonates. Bismuth solutions with alkaline hy- droxides or carbonates. [Pages 186, 191.] 8 Jiisrnuth .sulphide precipitated l'rom solutions of bismuth by hydro- gen sulphide. [Page 191. LEAD—COPPER — BISMUTH. 189 of oxide when exposed to the air, and as the oxide is easily dissolved by the combined action of water, carbonic or other acids, such as are found in vine- gar, the juice of fruits, or rancid fats, the use of copper for culinary vessels is always dangerous. Actual adulterations of food with compounds of copper have been detected. In cases of poisoning by copper the stomach-pump should be used, vomiting induced, and albumen (white of egg) administered, which forms an insoluble compound with copper. Reduced iron, or a very dilute solution of potassium ferrocyanide, may be of use as antidotes. Analytical reactions. (Cupric sulphate, CuS04, may be used.) 1. Add to solution of copper, hydrogen sulphide or ammonium sul- phide : a black precipitate of cupric sulphide is formed. (Plate III., 1): CuS04 + H2S = H2S04 + CuS. 2. Add sodium or potassium hydroxide: a bluish precipitate of cupric hydroxide, Cu(OH)2, is formed which is converted into dark- brown cupric oxide, CuO, by boiling. (See equation above.) (Plate III., 2.) 3. Add ammonium hydroxide in excess: a deep azure-blue solution is produced, containing an ammonio-copper compound. (See explana- tion above.) (Plate III., 3.) 4. Add potassium ferrocyanide: a reddish-brown precipitate of cupric ferrocyanide, Cu2Fe(CN)6, is obtained. (Plate III., 5.) 5. Add solution of’arsenious acid and carefully neutralize with sodium hydroxide: green cupric arsenite is precipitated. (Plate V., 2.) 6. Add sodium or potassium carbonate: green cupric carbonate with hydroxide is precipitated. (Plate III., 4.) 7. Immerse a piece of iron, or steel, showing a bright surface, in an acidified solution of copper : the latter is precipitated upon the iron, an equivalent amount of iron passing into solution: CuS04 -f- Fe = FeS()4 -)- Cu. 8. Most compounds of copper color the flame green, cupric chloride blue. 9. Cupric compounds give a blue, cuprous compounds a red borax bead. 10. Cupric salts (when not anhydrous) have mostly a blue or green color: sulphate, nitrate, chloride, and the ammonio-copper compounds are soluble, most other compounds are insoluble. 190 METALS AND THEIR COMBINATIONS. Bismuth, Bim = 208.9. Found in nature chiefly in the metallic state, disseminated, in veins, through various rocks. The extraction of the metal is a mere mechanical process, the earthy matter contain- ing it being heated in iron cylinders, and the melted bismuth collected in suitable receivers. Bismuth is grayish-white, with a pinkish tinge, very brittle, gen- erally showing a distinct crystalline structure. Occasionally it is used in alloys and in the manufacture of a few medicinal prepara- tions. Bismuth is trivalent, as shown in the chloride, BiCl3, or oxide, Bi203. A characteristic property of this metal is decomposition of the concentrated solution of any of its normal salts by the addition of much water, with the formation and precipitation of so-called oxy- salts or subsalts of bismuth, while some bismuth with a large quantity of acid remains in solution. The true constitution of these subsalts is as yet doubtful, but a com- parison of them has led to the assumption of a radical Bismuthyl, BiO, which behaves like an atom of a univalent metal. The relation between the normal or bismuth salts, and the subsalts or bismuthyl salts, will be shown by the composition of the following compounds : Bismuth chloride, BiCI3. Bismuthyl chloride, (BiO)Cl. “ bromide, BiBr3. “ bromide, (BiO)Br. “ iodide, BiT3. *• iodide, (BiO)I. “ nitrate, Bi(N03)3. “ # nitrate, (BiO)NOs. sulphate, Bi2(S04)3. “ sulphate, (Bi0)2SO4. “ carbonate, Bi2(C03)3 not known. “ carbonate, (Bi0)2C03. Bismuthyl nitrate, Subnitrate of bismuth, Bismuthi subnitras, Bi0N03.H20 ? (Oxynitrate of bismuth). By dissolving metallic bismuth in nitric acid, a solution of bismuth nitrate is obtained, nitrogen dioxide escaping: Bi + 4HN03 = Bi(N03)3 + NO + 2H20. Upon evaporation of the solution, colorless crystals of bismuth nitrate, Bi(N03)35H20, are obtained. If, however, the solution (or the dissolved crystals) be poured into a large quantity of water, the salt is decomposed with the formation of bismuthyl nitrate and nitric acid, which latter keeps in solution some bismuth : Bi(N03)3 '+ 2H20 = Bi0N03.H20 + 2HN03. LEAD — COPPER— BISMUTH. 191 Subnitrate of bismuth is a heavy, white, tasteless powder, insoluble in water, soluble in most acids. Experiment 33. Dissolve by the aid of heat about 1 gramme of metallic bis- muth in a mixture of 2 c.c. of nitric acid and 1 c.c. of water. Evaporate the clear solution to about one-half its volume, in order to remove excess of acid, and pour this solution of normal bismuth nitrate into 100 c.c. of water. Col- lect the precipitate of bismuthyl nitrate on a filter, wash and dry it. Prove the presence of bismuth in the filtrate by tests mentioned below. Bismuthyl carbonate, Subcarbonate of bismuth, Bismuthi subcar- bonas, (Bi0)2C03.H20 (?) (Oxy carbonate of bismuth, Pearl-white). Made by adding sodium carbonate to solution of bismuth nitrate, when the subcarbonate is precipitated, some carbon dioxide escaping : 2[Bi(N03)3] + 3Na2C03 + H20 = 6NaN03 + 2C02 + (Bi0)2C03.H20. A white, or pale yellowish-white powder, resembling the subnitrate. It readily loses water and carbon dioxide on heating, when the yellow oxide, Bi203, is left. Bismuthyl iodide, Subiodide of bismuth, BiOI, may be obtained by adding solution of hydriodic acid to freshly precipitated bismuth oxide: Bi203 + 2HI = 2BiOI + H20. A better method for making the compound is to pour gradually a solution, made by dissolving 95 grammes of crystallized normal bis- muth nitrate in 125 c.c. of glacial acetic acid, into a solution of 40 grammes of potassium iodide, and 55 grammes of sodinm acetate in 2500 c.c. of water. The precipitate, which has a brick-red color, is well washed and dried at 100° C. (212° F.). The decomposition is this : 2[Bi(N03)3] + 2H..0 + 2KI + 4NaC,H,0„ = 2(BiOI) + 4NaN03 + 2KN03 + 4C2H402. Analytical reactions. (Bismuth nitrate, Bi(N03)3, or bismuth chloride, BiCl3, may be used.) 1. Add to solution of bismuth, hydrogen sulphide or ammonium sulphide: a dark-brown (almost black) precipitate of bismuth sul- phide, Bi2S3, is produced (Plate III., 8): 2. Pour a concentrated solution of bismuth into water : a white pre- cipitate of a bismuthyl salt is formed. (See explanation above.) 3. Add to bismuth solution ammonium or sodium hydroxide, or 2BiCl3 + 3H2S = 6HC1 + Bi2S3. 192 METALS AND THEIR COMBINATIONS. carbonate: a white precipitate of bismuth hydroxide, Bi30H, or of bismuthyl carbonate is produced. (See explanation above.) 4. Potassium iodide precipitates brown bismuth iodide, Bil3, soluble in excess of the reagent. 5. Potassium dichromate precipitates yellow bismuthyl dichromate, (BiO)2Cr2Or 6. A small quantity of bismuth or of any bismuth compound, mixed with sulphur and potassium iodide, and heated upon charcoal before the blow-pipe, forms a scarlet-red incrustation of bismuthyl iodide, BiOI. 29. SILVER—MERCURY. Silver, Ag = 107.7 [Argentum). This metal is found sometimes in the metallic state, but generally as a sulphide, which is nearly always in combination with large quantities of lead sulphide, such ore being known as argentiferous galena. The lead manufactured from this ore contains the silver, and is separated from it by roasting the alloy in a current of air, whereby lead is oxidized and converted into litharge, while pure silver is left. Silver is the whitest of all metals, and takes the highest polish; it is the best conductor of heat and electricity, and melts at about 1000° C. (1873° F.); it is univalent, and forms but one series of salts; it is not affected by the oxygen of the air at any temperature, but is readily acted upon by traces of liydrosulphuric acid, which forms a black film of sulphide upon the surface of metallic silver. Hydrochloric acid scarcely acts on silver, nitric and sulphuric acids dissolve it. While many of the non-metallic elements have long been known to exist in allotropic forms, none of the metals had been, obtained in such a condition until quite recently, when it was shown that silver is capable of assuming a number of allotropic modifications. These are obtained chiefly by precipi- tating silver from solutions by different reducing agents. While normal silver Questions.—271. What are the properties of lead and from what ore is it obtained? 272. What is litharge, and how does it differ from red lead ? 273. Give the composition of nitrate, carbonate, and iodide of lead; how are they made? 274. State the analytical reactions for lead. 275. How is copper found in nature? 276. How many oxides of copper are known ; what is their composition, and under what conditions are they formed ? 277. What is “ blue vitriolhow is it made, and what are its properties? 278. How does ammo- nium hydroxide act on cupric solutions? 279. Mention tests for copper. 280. What is the composition of subnitrate and subcarbonate of bismuth ; how are they made from metallic bismuth, and what explanation is given in regard to their constitution ? SILVER—MERCURY. 193 is white, the allotropic forms have distinct colors—blue, bluish-green, red, pur- ple, yellow—and differ also in many other respects. Thus they are converted into silver chloride by highly diluted hydrochloric acid, which does not act on common silver; they are soluble in ammonia water, and act as reducing agents upon a number of substances, such as permanganates, ferricyanides, etc. Allo- tropic silver can be converted into the common form by different forms of energy—for instance, by heat, electricity, and the action of strong acids. Silver is too soft for use as coin or silverware, and, therefore, is alloyed with from 5 to 25 per cent, of copper, which causes it to be- come harder, and consequently gives it more resistance to the wear and tear by friction. Pure silver may be obtained by dissolving silver coin in nitric acid, when a blue solution, containing the nitrates of copper and silver, is formed. By the addition of sodium chloride to the solution a white precipitate of silver chloride forms, while cupric nitrate remains in solution. The silver chloride is washed, dried, mixed with sodium carbonate, and heated in a crucible, when sodium chloride is formed, carbon dioxide escapes, and a button of silver is found at the bottom of the crucible : 2AgCl + N82C03 = 2NaCl + C02 + 2Ag + O. Experiment 34. Dissolve a small silver coin in nitric acid, dilute with water, and precipitate the clear liquid with an excess of solution of sodium chloride. The washed precipitate of silver chloride may be treated with sodium carbon- ate, as stated above, or may be converted into metallic silver by the following method: Place the dry chloride in a small porcelain crucible and apply a gentle heat until the chloride has fused; when cold, place a piece of sheet zinc upon the chloride, cover with water, to which a few drops of sulphuric acid have been added, and set aside for a day, when the silver chloride will be found to have been decomposed with liberation of metallic silver and formation of zinc chloride: 2AgCl + Zn = ZnCl2 + 2Ag. Wash the spongy silver with dilute sulphuric acid and then with water. Use this silver for making silver nitrate by dissolving it in nitric acid, and evaporation of the solution to dryness. Use this solution for silver reactions. Silver nitrate, Argenti nitras, AgN03 = 169.7 (Nitrate of silver). Pure silver is dissolved in nitric acid : 3Ag + 4IINO3 = NO + 2H20 + 3AgN03. The solution is evaporated to dryness with the view of expelling all free acid, the dry mass dissolved in hot water and crystallized. If the silver used should contain copper, the latter may be eliminated from the mixture of silver and cupric nitrate by evaporating to dry- ness and fusing, when the latter salt is decomposed, insoluble cupric 194 METALS AND THEIR COMBINATIONS. oxide being formed. The fused mass is dissolved in water, filtered, and again evaporated for crystallization. When fused and poured into suitable moulds it yields the white cylindrical sticks which are usually known as caustic, lunar caustic, or lapis infernalis. When fused with an equal weight of potassium nitrate and formed into similar rods, it represents the diluted nitrate of silver (mitigated caustic) of the U. S. P. Silver nitrate forms colorless, transparent, tabular, rhombic crystals, or, when fused, a white, hard substance; it is soluble in less than its own weight of water, the solution having a neutral reaction. Exposed to the light, especially in the presence of organic matter, silver nitrate blackens in consequence of decomposition; when brought in contact with animal matter, it is readily decomposed into free nitric acid and metallic silver, which produces the characteristic black stain ; it is this decomposition, and the action of the free nitric acid, to which the strongly caustic properties of silver nitrate are due. Nitrate of silver is used largely in photography, and also in the manufacture of various kinds of indelible inks and hair-dyes. Silver oxide, Argenti oxidum, Ag20 = 231.4 (Oxide of silver). Made by the addition of an alkali hydroxide to silver nitrate: 2AgN03 + 2K0H = 2KN03 + H20 + Ag20. A dark-brown, almost black powder, but very sparingly soluble ' in water, imparting to the solution a weak alkaline reaction. It is a strong base, and easily decomposed into silver and oxygen. Silver iodide, Argenti iodidum, Agl = 234.3 (.Iodide of silver). Made by the addition of potassium iodide to silver nitrate : AgNOg + KI = KN03 + Agl. A heavy, amorphous, light-yellowish powder, insoluble in water, and but slightly soluble in ammonium hydroxide. Antidotes. Sodium chloride, white of egg, or milk, followed by an emetic. Analytical reactions. (Silver nitrate, AgN03, may be used.) 1. Add to solution of a silver salt, hydrogen sulphide or ammonium sulphide : a dark-brown precipitate of silver sulphide is produced : 2AgNOs + H2S = 2HN0S + Ag2S. SILVER — MERCURY. 195 2. Add hydrochloric acid, or any soluble chloride : a white, curdy precipitate of silver chloride is produced, which is insoluble in dilute acids, but soluble in ammonium hydroxide and in potassium cyanide. AgNOg -f NaCl = NaNOj + AgCl. 3. Add potassium chromate or dichromate: a red precipitate of silver chromate, Ag2Cr04, is formed (Plate II., 7). 4. Add sodium phosphate: a pale-yellow precipitate of silver phos- phate, Ag3P04, is formed, which is soluble in ammonia aud in nitric acid. 5. Alkali hydroxides precipitate dark-brown silver oxide, soluble in ammonia water. 6. Potassium iodide or bromide gives a pale-yellow precipitate. (See above.) 7. Metallic copper, zinc, or iron precipitates metallic silver. Mercury, Hydrargyrum, Hg = 199.8 (Quicksilver). Mercury is found sometimes in small globules in the metallic state, but generally as mercuric sulphide or cinnabar, a dark-red mineral. The chief supply was formerly obtained from Spain and Austria; now, how- ever, large quantities are obtained from California; it is also imported from Peru and Japan. Mercury is obtained from cinnabar either by roasting it, whereby the sulphur is converted into sulphur dioxide, or by heating it with lime, which combines with the sulphur, while the metal volatilizes, and is condensed by passing the vapors through suitable coolers. Mercury is the only metal showing the liquid state at the ordinary temperature ; it solidifies at —40° C. (—40° F.), and boils at 357° C. (675° F.); but is slightly volatile at all temperatures; it is almost silver-white, and has a bright metallic lustre; its specific gravity is 13.57 at 15° C. (59° F.). Mercury is peculiar in that its molecule contains but one atom, at least when in the state of a gas; in the liquid and solid states it may contain two atoms, like most other elements, but we have as yet no means of proving this fact. Mercury is bivalent, and forms, like copper, two series of com- pounds, distinguished as mercuric and mercurous compounds. In the former, one atom of mercury exerts its bivalence, as in HgO, HgCl2; in the mercurous compounds two atoms of mercury exert the same valence, as in Hg20, Hg2Cl2. In order to explain this behavior we have to assume that of the four points of attraction, represented METALS AND THEIR COMBINATIONS. by the two atoms of mercury, two are required to hold together or unite these two atoms, so as to leave but two for other elements. ,C1 XC1 Hg—01 I Etg—Cl Mercuric chloride. Mercurous chloride. There are known, however, some data which seem to contradict this view and make it not unlikely that the composition of mercurous chloride is HgCl, and not Hg2Cl2. Mercury is not affected by the oxygen of the air, nor by hydro- chloric acid, while chlorine, bromine, and iodine combine with it directly, and warm sulphuric and nitric acids dissolve it. Mercury is used in the metallic state for many scientific instruments (thermometer, barometer, etc.); in the silvering of looking-glasses, which is effected by means of an amalgam of tin (amalgams are alloys in which mercury is one of the constituents); for manufacturing from it all of the various mercury compounds, and those officinal prepara- tions in which mercury exists in the metallic state. These latter preparations are : Mercury with chalk, blue mass or blue pill, mercurial ointment, and mercurial plaster. Mercury exists in a metallic, but highly subdivided state in these preparations, which are made by intimately mixing (triturating) metallic mercury with the different substances used (viz., chalk, pill-mass, fat, lead-plaster) It is most probable that the action of these agents upon the animal system is chiefly due to the conversion of small quantities of mercury into mercurous oxide, which, in contact with the acids of the gastric juice or with perspiration, are converted into soluble compounds capa- ble of being absorbed. Mercurous oxide, Hg20 (Black oxide or suboxide of mercury). An almost black, insoluble powder, made by adding an alkaline hydroxide to a solution of mercurous nitrate: Hg22NOg + 2K0H = 2KN03 + H20 + Hg20. A similar decomposition takes place when alkaline hydroxides are added to insoluble mercurous chloride. A mixture of lime-water and mercurous chloride (calomel) is known as black-wash ; when the two substances are mixed, calomel is converted into mercurous oxide, while calcium chloride is formed: Hg2C)2 + Ca(OH)2 = CaCl2 + H20 + Hg20. Mercuric oxide, Oxide of mercury, HgO = 215.8. There are two SILVER—MERCURY. 197 mercuric oxides which are officinal; they do not differ in their chemi- cal composition, but in their molecular structure. The yellow oxide of mercury, Hydrargyri oxidum flavum, is made by pouring a solution of mercuric chloride into a solution of sodium hydroxide, when an orange-yellow, heavy precipitate is produced, which is washed and dried at a temperature not exceeding 30° C. (86° F.) (Plate IV., 3): HgCl2 + 2NaOH = HgO + 2NaCl + H20. The red oxide of mercury, Hydrargyri oxidum rubrum, or red pre- cipitate, is made by heating mercuric nitrate, either by itself or after it has been intimately mixed with an amount of metallic mercury equal to the mercury in the nitrate used (Plate IV., 4). In the first case, nitrous fumes and oxygen are given off, mercuric oxide remain- ing: Hg(N03)2 = HgO + 2N02 + 0. ** In the other case, no oxygen is evolved : Hg(N03)2 + Hg = 2HgO + 2N02. The red oxide of mercury differs from the yellow oxide in being more compact, and of a crystalline structure ; while the yellow oxide is in a more finely divided state, and consequently acts more energeti- cally when used in medicine Yellow oxide, when digested on a water-bath with a strong solution of oxalic acid, is converted into white mercuric oxalate within fifteen minutes, while red oxide is not acted upon by oxalic acid under the same conditions. When mercuric chloride is added to lime-water, a mixture is formed holding in suspension a yellow oxychloride of mercury; this mixture is known as yellow-wash. Experiment 35. Heat some mercuric nitrate in a porcelain dish, placed in a fume chamber, until red fumes no longer escape. The remaining red powder is mercuric oxide, which, by further heating, may be decomposed into its ele- ments. Mercurous chloride, Hydrargyrum chloridum mite, Hg2Cl2 —470.4 (Calomel, Mild chloride of mercury, Subchloride or protochloride of mercury). Mercurous chloride, like mercurous oxide, may be made by different processes, but the article used medicinally is the one obtained (except it be otherwise stated) by sublimation and the rapid condensation of the vapor in the form of powder. It is made either by subliming a mixture of mercuric chloride and mercury : HgCl2 + Hg = Hg2Cl2, 198 METALS AND THEIR COMBINATIONS. or by thoroughly mixing with mercuric sulphate a quantity of mer- cury equal to that contained in the sulphate; by this operation mer- curous sulphate is obtained, which is mixed with sodium chloride, and sublimed from a suitable apparatus into a large chamber, so that the sublimate may fall in powder to the floor : Precipitated calomel, being in a finer state of subdivision, acts more energetically when used medicinally. It is obtained by precipitation of a soluble mercurous salt by any soluble chloride : HgS04 + Hg + 2NaCl = Na2S04 + Hg2Cl2. Hg2(N03)2 + 2NaCl = 2NaN03 + Hg2Cl2. Mercurous chloride, made by either process, generally contains traces of mercuric chloride, and should, therefore, be washed with hot water until the washings are no longer acted upon by ammonium sul- phide or silver nitrate. Mercurous chloride is a white, impalpable, tasteless powder, insolu- ble in water and alcohol; it volatilizes without fusing previously; when given internally, it should not be mixed with either mineral acids, alkali bromides, iodides, hydroxides, or carbonates, except the action of the decomposition products be desired. Mercuric chloride, Hydrargyri chloridum corrosivum, HgCl2 = 270.6 (Corrosive chloride of mercury, Corrosive sublimate, Perchloride or bichloride of mercury). Made by thoroughly mixing mercuric sul- phate with sodium chloride, and subliming the mixture, when mer- curic chloride is formed, and passes off in white fumes which are condensed in the cooler part of the apparatus, while sodium sulphate is left: HgS04 + 2'NaCl = Na2S04 + HgCl2. Mercuric chloride is a heavy, white powder, or occurs in heavy, colorless, rhombic crystals or crystalline masses; it is soluble in 16 parts of cold and 2 parts of boiling water, in about 3 parts of alcohol, in 4 parts of ether, and in about 14 parts of glycerin; when heated, it fuses and is volatilized; it lias an acrid, metallic taste, an acid reaction, and strongly poisonous and antiseptic properties. Mercurous iodide, Hydrargyri iodidum viride, Hg2I2 = 652.6 (Green iodide, yellow iodide, or protiodide of mercury). Both iodides of mercury may be obtained either by rubbing together mercury and iodine in the proportions represented by the respective atomic weights, SILVER—MERCURY. 199 or by precipitation of soluble mercurous or mercuric salts by potas- sium iodide. According to the U. S. P., mercurous iodide is made by triturating 8 parts of mercury with 5 of iodine (and a little alcohol) until all globules of mercury have disappeared and a green powder has been formed, which is washed with alcohol in order to eliminate small quantities of mercuric iodide which may have been formed : 2Hg + 21 = Hg2T2. The powder finally is collected and dried between paper at a low temperature. During the whole operation light should be excluded as much as possible, as it decomposes the compound. Mercurous iodide is a yellow to greenish-yellow, tasteless powder, insoluble in water, alcohol, and ether. It is easily decomposed into mercuric iodide and mercury (Plate IV., 5). Mercuric iodide, Hydrargyri iodidum rubrum, Hgl2 = 452.8 (Red iodide or biniodide of mercury). Made by mixing solutions of potas- sium iodide and mercuric chloride, when a yellow precipitate is formed, turning red immediately (Plate IV., 6): HgCl2 + 2KI = 2KC1 + HgL,. Mercuric iodide is soluble both in solution of potassium iodide and mercuric chloride, for which reason au excess of either substance will cause a loss of the salt when prepared. It is a scarlet-red, tasteless powder, almost insoluble in water and but slightly soluble in alcohol; on heating or subliming it turns yellow in consequence of a molecular change which takes place; on cooling, and, more quickly, on press- ing or rubbing the yellow powder, it reassumes the original condition and the red color. Mercuric sulphate, HgS04. When mercury is heated with strong sulphuric acid (the presence of nitric acid facilitates the formation) chemical action takes place between the two substances, sulphur dioxide being liberated and mercuric sulphate formed, which is obtained as a heavy, white, crystalline powder : Hg + 2H2S04 = HgS04 + 2H20 + S02. Yellow subsulphate of mercury, Hydrargyri subsulphas flavus, HgS04.2Hg0 = 727.4 (Basic mercuric, sulphate, Turpeth mineral, Mer- curic oxy-sulphate). When mercuric sulphate, prepared as directed above, is thrown into boiling water, it is decomposed into an acid salt 200 METALS AND THEIR COMBINATIONS. which remains in solution, and a basic salt which is precipitated. As shown by its composition, HgS04.2HgO, it may be looked upon as mercuric sulphate in combination with mercuric oxide. It is a heavy, lemon-yellow, tasteless powder, almost insoluble in water. Mercurous sulphate, Hg2S04. When mercuric sulphate is tritu- rated with a sufficient quantity of mercury, direct combination takes place, and the mercurous salt is formed : HgS04 + Hg = Hg2S04. Nitrates of mercury. Mercurous nitrate, IIg2(V03)2, and Mercuric nitrate, Hg(N03)2, may both be obtained as white salts by dis olving mercury in nitric acid. The relative quantities of the two substances present determine whether mercurous or mercuric nitrate be formed. If mercury is present in excess the mercurous salt, if nitric acid is present in excess the mercuric salt, is formed, the latter especially on heating. Both salts are white and soluble in water. Experiment 36. Heat gently a small globule (about 1 gramme) of mercury with 2 c.c. of nitric acid until red fumes cease to escape. If some of the mer- cury remains undissolved, the solution will deposit crystals of mercurous nitrate on cooling. Use some of the solution, after being diluted with much water, for mercurous tests. Use another portion as follows: Heat the solution, or some of the crystals, with about an equal weight of nitric acid until no more red fumes escape. Add to a few drops of the diluted liquid a little hydro- chloric acid, which, if the conversion of the mercurous into mercuric salt has been complete, will give no precipitate. If, however, one should be formed, the solution has to be heated with more nitric acid until no longer a precipi- tate is formed by hydrochloric acid, when the solution is evaporated and set aside for crystallization. The respective changes may be represented by the following equations : 3[Hg2(N03)2] + 8HNO3 = 6[Hg(N03)2] + 4H20 + 2N0. 6Hg + 8HN03 = 3[Hg2(N03)2] + 4H20 + 2N0 ; Mercuric sulphide, HgS = 231.8. This compound has been men- tioned as the chief ore of mercury, occurring crystallized as cinnabar, which has a red color (Plate IV., 2). The same compound may, however, be obtained by passing hydrosulphuric acid gas through mercuric solutions, when at first a white precipitate is formed (a double compound of the sulphide of mercury in combination with the mercuric salt), which soon turns black (Place IV., 1): HgCl2 + H2S = 2HC1 + HgS. The black, amorphous, mercuric sulphide may be converted into the red, crystallized variety by sublimation, and is then the officinal preparation known as red sulphide of mercury, Hydrargyri sulphidum SILVER—MERCURY. 201 rubrum, cinnabar, or vermilion. It forms brilliant dark-red crystal- line masses, or a fine bright scarlet powder, which is insoluble in water, hydrochloric or nitric acid, but soluble in nitro-hydrochloric acid Mercuric and mercurous sulphides may be made also by triturating the elemeuts mercury and sulphur in the proper proportions, when they combine directly. Ammoniated mercury, Hydrargyrum ammoniatum, NH2HgCl = 251.2 (White precipitate, Mercur-ammonium chloride). This com- pound is made by pouring solution of mercuric chloride into water of ammonia, when a white precipitate falls, which is washed with highly diluted ammonia water and dried at a low temperature : HgCl2 + 2NH4OH = NH2HgCl + NH4C1 + 2H20. As shown by the composition of this compound, it may be re- garded as ammonium chloride, NH4C1, in which two atoms of hydro- gen have been replaced by one atom of the bivalent mercury. (There are many compounds known in which metallic atoms replace hydrogen in salts of ammonium; the ammonium copper compounds belong to this group of substances.) Ammoniated mercury is a white, tasteless, insoluble powder. Mercurous salts. (Mercurous nitrate, Hg22N03, may be used.) Mercuric salts. (Mercuric chloride, HgCl2, may be used.) 1. Hydrogen sul- phide, or ammo- nium sulphide. Black precipitate of mercuric sulphide, with mercury. Hg2(N03)2 + H2S — 2HN03 + HgS + Hg. Black precipitate of mercuric sulphide. (Precipitate may be white or gray, with an insuffi- cient quantity of the reagent.) (See above.) (Plate IV., 1.) 2. Potassium iodide. Green precipitate of mercurous iodide (Plate IV., 7): Hg2(N03)2 + 2KI = 2KN03 + Hg2T2. Red precipitate of mercuric iodide. (See above.) (Plate IV, 6.) 3. Potassium or so- dium hydroxide. Dark-brown precipitate of mer- curous oxide, Hg.,0 (Plate IV, 5). Yellow precipitate of mercuric oxide, HgO. ( See above-) (Plate IV, 3.) 4. Ammonium hy- droxide. Black precipitate of mercurous ammonium salt is formed. (The insoluble white calomel is converted into a black pow- der.) White precipitate of a mercuric ammonium salt is formed. (See explanation above.) 5. Potassium or so- dium carbonate. Yellowish precipitate of mer- curous carbonate, which is un- stable. Brownish - red precipitate of basic mercuric carbonate, HgC03 3HgO. 6. Hydrochloric acid or soluble chlorides. White precipitate of mercurous chloride is produced: Hg2(N03)2 + 2HC1 = 2HNO, + HgaCl2. No change. Analytical reactions. 202 METALS AND THEIR COMBINATIONS. 7. Stannous chloride produces, in solutions of mercury, a white precipitate, which turns dark-gray on heating with an excess of the reagent. The reaction is due to the strong reducing or deoxidizing property of the stannous chloride, which itself is converted into stannic chloride, while the mercury salt is first converted into a mercurous salt and afterward into metallic mercury : 2HgCl2 + SnCl2 = Hg2CI2 + SnCl4; Hg2Cl2 + SnCl2 = 2Hg + SnCI4. 8. Dry mercury compounds, when mixed with sodium carbonate and potassium cyanide, and heated in a narrow test-tube, are decom- posed with liberation of metallic mercury, which condenses in small globules in the cooler part of the tube. 9. A piece of bright metallic copper, when placed in a slightly acid mercury solution becomes coated with a dark film of metallic mercury, which by rubbing becomes bright and shining, and may be volatilized by heat. 10. All compounds of mercury are completely volatilized by heat, either with or without decomposition. Antidotes. Albumen (white of egg), of which, however, not too much should be given at one time, lest the precipitate formed by the mercuric salt and albumin be redissolved. The antidote should be followed by an emetic to remove the albuminous mercury compound. Questions.—281. How is silver obtained from the native ores, and how may it be prepared from silver coin? 282. State of silver nitrate: its composition, mode of preparation, properties, and names by which it is known. 283. Give analytical reactions for silver. 284. How is mercury found in nature; how is it obtained from the native ore; what are its physical and chemical properties? 285. Mention the three oxides of mercury; how are they made, what is their composition, what is their color and solubility? 286. State of the two chlorides of mercury: their names, composition, mode of preparation, solubility, color, and other properties. 287. Mention the same of the two iodides, as above, for the chlorides. 288. State the difference between mercuric sulphate, basic mer- curic sulphate, and mercurous sulphate. 289. What is formed when ammonium hydroxide, calcium hydroxide, potassium or sodium hydroxide is added to either mercurous or mercuric chloride? 290. Give tests answering for any mercury compound, and tests by which mercuric compounds may be distinguished from mercurous compounds. PLATE I-\T. MERCURY. SILVER. 1 Mercury sulphide precipitated from mercuric solutions by hydrosul- phuric acid. [Page 200,] Mercuric sulphide, Cinnabar. [Page 200.] 2 3 Yellow mercuric oxide precipi- tated from mercuric solutions by po- tassium hydroxide. [Pages 197, 201.] 4 Red mercuric oxide obtained by heating mercuric nitrate. [Page 197.] Mercurous oxide precipitated from mercurous solutions by potas- sium hydroxide.. [Page 196.] Silver Sulphide precipitated from silver solutions by hydrogen sul- phide. [Page 194.] 5 Mercuric iodide precipitated from mercuric solutions by alkali iodides. [Pages 199, 201,] 6 7 Mercruous iodide precipitated from mercurous solutions by alkali iodides. [Pages 199, 201.] Mercuric solutions with ammo- nium hydroxide. [Page 201.] Mer- curous solutions with soluble chlor- ides. \Page 198.] Silver solutions with soluble chlorides. [Page 195.j 8 SILVER— MERCURY. 203 Lead. Copper. Bismuth. Silver. Mercuric salts. Mercurous salts. Cadmium. Hydrogen sulphide Black precipi- Black precipi- Dark-brown Black precipi- Black precipi- Black precipi- Yellow pre- tate. tate. precipitate. tate. tate. tate. cipitate. Sodium hydroxide W hite precipi- Blue precipi- White precipi- Brown precipi- Brown, then Black precipi- White precipi- tate. tate, turning tate tate. yellow pre- tate. tate. brown on cipitate. boiling. Ammonia water . W hite precipi- Pale-blue pre- YV hite precipi- Brown precipi ■ White precipi- Black precipi- W hite precipi- tate. cipitate. tate tate. tate. tate. tate. In excess of reagent Insoluble. Dark blue Insoluble. Colorless solu- Insoluble. Insoluble. Colorless solu- solution. tion. tion. Sodium carbonate White precipi- Greenish-blue White precipi- Pale-yellow Reddish-brown Yellowish pre- White precipi- tate. precipitate. tate. precipitate. precipitate. cipitate. tate. Potassium iodide Yellow pre- Yellow pre- Brown pre- Pale-yellow Scarlet-red Yellowish- cipitate. cipitate. cipitate. precipitate. precipitate. green precipi- tate In excess of reagent Insoluble. Insoluble. Insoluble. Insoluble. Soluble. Partly soluble. Potassium chromate Yellow pre- Orange pre- Dark-red pre- Orange precipi- Brick-red pre- cipitate. cipitate. cipitate. tate. cipitate. Hydrochloric acid White precipi- tate, soluble . tate, soluble tate, turning in hot water. in ammonia dark with water. ammonia Sulphuric acid White precipi- tate. tate. Summary of analytical characters of metals of the lead group. 204 METALS AND THEIR COMBINATIONS. 30. ARSENIC. As = 74.9. General remarks regarding the metals of the arsenic group. The metals belonging to either of the five groups considered heretofore, show much resemblance to each other in their chemical properties, and consequently in their combinations. This is much less the case among the six metals (As, Sb, Sn, Au, Pt, Mo) which are classed together in this group. In fact, the only resemblance which unites these metals is the insolubility of their sulphides in dilute acids and the solubility of these sulphides in ammonium sulphide (or alkaline hydroxides), with which they form soluble double compounds ; the oxides have also a tendency to form acids. In all other respects no general resemblance exists between these metals. Arsenic and antimony have many.prop- erties in common, and resemble in many respects the non-metallic elements phosphorus and nitrogen, as may be shown by a comparison of their hydrides, oxides, acids, and chlorides.* nh3 n2o3 n2o5 nci8. ph3 p2o3 p2o5 h3po4 ** 3. AsI13 As203 As205 H3A.s04 AsC13. SbH3 Sb203 Sb205 ' SbCl3. Arsenic. Found in nature sometimes in the native state, but gener- ally as sulphide or arsenide. One of the most common arsenic ores is the arsenio-sulphide of iron, or mispickel, FeSAs. Realgar is the native red sulphide, As2S2, and orpiment or auripigment, the native yellow sulphide, As2S3. Arsenides of cobalt, nickel, and other metals are not infrequently met with in nature. Certain mineral waters con- tain traces of arsenic compounds. Arsenic may be obtained easily by heating arsenous oxide with charcoal, or by allowing vapors of arsenous oxide to pass over char- coal heated to redness: As203 + 3C = 3C0 + 2As. In both cases the arsenic, when liberated by the reducing action of the charcoal, exists in the form of vapor, which condenses in the cooler part of the apparatus as a steel-gray metallic mass, which when exposed to the atmospheric air, loses the metallic lustre in consequence of the formation of a film of oxide. When pure, arsenic is odorless and tasteless; it is very brittle, and volatilizes unchanged and without melting when heated to 180° C. (356° F.), without access of air. Heated in air, it burns with a bluisli- ARSENIC. 205 white light, forming arsenous oxide. Although insoluble in water, yet water digested with arsenic soon contains some arsenous acid in solution, the oxide of arsenic being formed by oxidation of the metal by the oxygen absorbed in the water. Arsenic is used in the metallic state as fly-poison, and in some alloys, chiefly in shot, an alloy of lead and arsenic. The molecule of arsenic contains four atoms, and not two, like most elements. It is trivalent in some compounds, quinquivalent in others. Arsenous oxide, Acidum arsenosum, As203 = 197.8 (.Arsenious oxide, White arsenic, Arsenic trioxide, Arsenous anhydride, improperly Arsen- ous acid). This compound frequently is obtained as a by-product in metallurgical operations during the manufacture of metals from ores containing arsenic. Such ores are roasted (heated in a current of air), when arsenic is converted into arsenous oxide, which, at that tempera- ture, is vo'atilized and afterward condensed in chambers or long flues. Arsenous oxide is a heavy, white solid, occurring either as an opaque, slightly crystalline powder, or in transparent or semi-trans- parent masses which frequently show a stratified appearance; recently sublimed, arsenous oxide exists as the amorphous semi-transparent glassy mass known as vitreous arsenous oxide, which gradually becomes opaque and ultimately resembles porcelain. This change is due to a rearrangement of the molecules into crystals which can be seen under the microscope. The two modifications of arsenous oxide differ in their solubility in water, the amorphous or glassy variety dissolving more freely than the crystallized. One part of arsenous oxide dissolves in from 30 to 100 parts of cold and in 15 parts of boiling water, the solution having at first a faint acrid and metallic, and afterward a sweetish taste. This solution contains the arsenous oxide not as such, but as arsenous acid, H3As03, which compound, however, cannot be obtained in an isolated condition, but is known in solution only : A?203 + 3H20 = 2H3 As03. The salts of arsenous acid are known as arsenites. When heated to about 220° C. (428° F.) arsenous oxide is volatil- ized without fusion ; the vapors, when ceudensed, form small, shining, eight-sided crystals; when heated on charcoal, it is deoxidized, giving off, at the same time, an odor resembling that of garlic. Arsenous oxide is frequently used in the arts and for manufacturing 206 METALS AND THEIR COMBINATIONS. purposes, as, for instance, in the manufacture of green colors, of opaque white glass, in calico-printing, as a powerful antiseptic for the preser- vation of organic objects of natural history, and, finally, as the sub- stance from which all arsenic compounds are obtained. The officinal solution of arsenous acid, Liquor acidi arseniosi, is a 1 per cent, solution of arsenous oxide in water to which 2 per cent, of hydrochloric acid has been added. The officinal solution of arsenite of potassium, Liquor potassii arse- nitis, or Fowler s solution, is made by dissolving 1 part of arsenous oxide and 1 part of potassium bicarbonate in 95 parts of water and adding 3 parts of compound tincture of lavender; the solution con- tains the arsenic as potassium arsenite. Arsenic oxide, As205 (Arsenic pentoxide, Anhydrous arsenic acid). When arsenous oxide is heated with nitric acid, it becomes oxidized and is converted into arsenic acid, H3As04, from which the water may be expelled by further heating, when arsenic oxide is left: 2H3A.s04 = As205 + 3H20. Arsenic oxide is a heavy, white, solid substance which, in contact with water, is converted into arsenic acid. This acid resembles phos- phoric acid not only in composition, but also in forming metarsenic and pyroarsenic acid under the same conditions under which the cor- responding phosphoric acids are formed. The salts of arsenic acid, the arsenates, also resemble in their constitution the corresponding phosphates. Arsenic oxide and arsenic acid are used largely as oxidizing agents in the manufacture of aniline colors. Disodium hydrogen arsenate, Sodii arsenas, Na2HAs04.7H20 = 311.9 (.Arsenate of sodium). This salt is made by fusing arsenous oxide with carbonate and nitrate of sodium. As203 H~ 2NaIs03 -|- Na2C03 = Na4As207 -j~ lv203 H- C02. Sodium pyroarsenate is formed, nitrogen trioxide and carbon dioxide escaping. By dissolving in water and crystallizing, the officinal salt is obtained in colorless, transparent crystals: Nh4A?207 + 15H20 = 2(Na2HAs04.7Hs0). Hydrogen arsenide, AsH3 (Arsine, Arsenetted or arseniuretted hydrogen). This compound is formed always when either arsenous or arsenic oxides or acids, or any of their salts, are brought in contact ARSENIC. 207 with nascent hydrogen, for instance, with zinc and diluted sulphuric acid, which evolve hydrogen : Ap203 + 12H = 2AsH3 + 3H20. As205 + 16H = 2AsH3 + 5H20. AsC13 + 6H = AsH3 + 3HC1. Hydrogen arsenide is a colorless, highly poisonous gas, having a strong garlic odor. Ignited, it burns with a bluish flame, giving off* white clouds of arsenous oxide : 2AsH3 -f~ 60 = As203 3H20. When a cold plate (porcelain answers best) is held in the flame of arsenetted hydrogen, a dark deposit of metallic arsenic (arsenic spots) is produced upon the plate (in a similar manner as a deposit of carbon is produced by a common luminous flame). The formation of this metallic deposit may be explained by the fact that the heat of the flame decomposes the gas, and that, furthermore, of the two liberated elements, arsenic and hydrogen, the latter has the greater affinity for oxygen. In the centre of the flame, to which but a limited amount of oxygen penetrates, the latter is taken up by the hydrogen, arsenic being present in the metallic state until it burns in the outer cone of the flame. It is this liberated arsenic which is deposited upon a cold substance held in the flame. Arsenetted hydrogen, when heated to redness, is decomposed into its elements; by passing the gas through a glass tube heated to red- ness, the liberated arsenic is deposited in the cooler part of the tube, forming a bright metallic ring. Sulphides of arsenic. Two sulphides of arsenic are known and have been mentioned above as the native disulphide or realgar, As2S2, and the trisulphide or orpiment, As2S3. Disulphide of arsenic is an orange-red, fusible, and volatile substance, used as a pigment; it may be made by fusing together the elements in the proper propor- tions. Trisulphide is a golden-yellow, fusible, and volatile substance, which also may be obtained by fusing the elements, or by precipitating an arsenic solution by hydrogen sulphide (Plate V., 1). Both sul- phides of arsenic are sulpho-acids, uniting with other metallic sulphides to form sulpho-salts, as, for instance, K2S.As2S3, or (NH4)2S.As2S3. These compounds are known as sulph-arsenides. Arsenous iodide, Arsenii iodidum, Asl3 == 454.7 (Iodide of arsenic), may be obtained by direct combination of the elements, and forms 208 METALS AND THEIR COMBINATIONS. orange-red crystalline masses, soluble in water and alcohol, but decom- posed by boiling with either of these liquids. It is used in the officinal preparation, Solution of iodide of arsenic and mercury, Donovan’s solu- tion, which is made by dissolving one part each of arsenous iodide and mercuric iodide in 98 parts of water. Analytical reactions. (Use arsenous oxide, As203, and sodium arsenate, Na2HAs04, respectively.) 1. Add hydrogen sulphide to an aqueous solution of arsenous oxide: a yellow coloration, but no precipitate is formed until some hydro- chloric acid is added, when yellow arsenic trisulphide, As2S3 (Plate V., 1) is precipitated : A<203 + 3H2S = 3H20 + Ap2S3 ; or 2HsAs03 + 3H2S = 6H20 + A,2S3. When hydrogen sulphide is added to a cold solution of arsenic oxide or of an arsenate, acidified with hydrochloric acid, a yellow mix- ture of arsenic trisulphide, As2S3, and sulphur is slowly precipitated : When the same substances act upon one another in hot solution, and when also an excess of hydrogen sulphide (preferably a current of the gas) is used, yellow arsenic pentasulphide is precipitated : 2H3As04 + 5H2S = 8H20 + A.«2S3 + 2S. or Ap206 4" 5H2S = 5H20 -f- A?2S5; 2H3As04 + 5H2S = 8H20 + A?2S5. 2. Add ammonium sulphide or any alkali hydroxide to the yellow precipitate of arsenous or arsenic sulphide : the precipitates are readily dissolved, but may be reprecipitated by neutralizing with an acid. 3. Ammonio-nitrate of silver (silver nitrate to which enough of water of ammonia has been added to redissolve the precipitate at first formed) produces in neutral solutions of arsenous acid a yellow precipi- tate of arsenite of silver, Ag3As03 (Plate V., 3); in arsenic acid solu- tions a reddish-brown precipitate of arsenate of silver, Ag3As04 (Plate V., 4). The two precipitates are soluble in alkalies and acids both. Silver arsenite dissolved in water of ammonia and boiled forms silver arsenate and metallic silver. 4. Ammonio-sulphate of copper (made similarly to ammonio-nitrate of silver from cupric sulphate) added to neutral arsenous solutions produces a green precipitate of cupric arsenite (CuHAsOs) known as Scheele’s green (Plate V., 2). (Arsenite of copper mixed with cupric Plate ARSENIC. ANTIMONY. TIN. 1 Arsenous sulphide, precipitat ed from arsenous solutions by hyro- gen sulphide. [Page 208] 2 Cupric arsenite, precipitated from arsenous solutions by ammonio- sulphate of copper. [Page 208] Silver arsenite, precipitated from arsenous solutions by silver nitrate. [Page 208] 3 Silver arsenate, precipitated from arsenic solutions by silver nitrate. [Page 208] 4 5 A ntimonious sulphide, precipi- tated from solutions of antimony by hydrogen sulphide. [Pages 214, 216.] Native or Crystallized anti- mon ious sulphide. [Page 211] 0 7 Stannous sulphide, precipitat- ed from stannous solutions by hydro- gen sulphide. [Page 217] Stannic sulphide, precipitated from stannic solutions by hydrogen sulphide. [Page 217] 8 ARSENIC. 209 acetate is known as Schweinfurth green). The same reagent produces in neutral solutions of an arsenate a similar green precipitate of cupric arsenate, CuHAs04. Cupric arsenite boiled with potassium hydroxide forms potassium arsenate and red cuprous oxide. Instead of using for the above tests the ammonio salts, silver nitrate or cupric sulphate may be added to the acid (or neutral) solution of arsenic, then adding water of ammonia carefully in small quantities until a neutral reaction has been obtained, when the precipitate is formed. 5. Soluble arsenates give white precipitates with soluble salts of barium, calcium, magnesium, zinc, and some other metals; soluble arsenites do not. Arsenates give, on heating with ammonium molyb- date, a yellow precipitate of ammonium arseno-molybdate, (HH4)3As04. 10MoO3. 6. Heat any dry arsenic compound, after being mixed with some charcoal and dry potassium carbonate in a very narrow test-tube (or, better, in a drawn-out glass tube having a small bulb on the end): the arsenic compound is decomposed and the metallic arsenic deposited as a metallic ring in the upper part of the contraction. (Fig. 18.) Fig. 13. 7. Heat arsenous or arsenic oxide upon a piece of charcoal by means of a blowpipe; a characteristic odor of garlic is perceptible. 8. jReinsch's test. A thin piece of copper, having a bright metallic surface, placed in a slightly acidified solution pf arsenic becomes, upon heating the solution, coated with a dark steel-gray deposit of arsenic, which can be vaporized by application of heat. 210 METALS AND THEIR COMBINATIONS. 9. Bettendorf’s test. Acid to any arsenic compound, dissolved in concentrated hydrochloric acid, an equal volume of freshly prepared solution of stannous chloride in hydrochloric acid, add a small piece of tin-foil, and apply heat: a brown color or precipitate is formed, due to the separation of arsenic. 10. Gutzeit’s test. Place a small piece (about 1 gramme) of pure zinc in a test-tube, add about 5 c.c. of dilute (5 per cent.) sulphuric acid and a few drops of any arsenic solution, which should not be alkaline. Fasten over the mouth of the test-tube a cap made of three thicknesses of pure filter paper, and moisten the upper paper with a drop of a saturated solution of silver nitrate in water, acidulated with about 1 per cent, of nitric acid. (Fig. 14.) Place the tube in a box so as to exclude all light, and examine the paper cap after awhile. Upon it will appear a bright yellow stain, rapidly if the quantity of arsenic be considerable, slowly if it be small. Upon moistening the yellow stain with water the color changes to brown or black. The action of hydrogen arsenide upon silver nitrate in the absence of water takes place with the formation of a yellow compound, thus : AsH3 -f- 6AgN03 = 3HN03 + Ag3A.s.3(AgN03). In the presence of water metallic silver is separated, showing a black or brown color: AsH3 + 6AgN03 -f- 6H20 = 6HN03 -f- EL3As03 6Ag. Compounds of antimony treated in the above manner produce a dark spot upon the paper, but cause no pre- vious yellowT color. 11. Fleitmann’s test. This is similar to the previous test, the chief difference being that hydrogen is evolved in alkaline solution, which has the advantage that the presence of antimony does not interfere, because this metal does not form autimonetted hydrogen in alkaline solutions. Place about 1 gramme of pure zinc in a test-tube, add about 5 c.c. of potassium hydroxide solution and a few drops of the arsenic solu- tion, which should not be acid. Provide paper cap as described in previous test, and set the test-tube in a box containing sand heated to about 90° C. (194° F.). A brown or black stain of metallic silver will appear upon the paper. 12. Marsh’s test. While this test is not used now for qualitative Fro. 14. ARSENIC. 211 determinations as much as formerly, it is still of great value because it may serve for collecting the total amount of arsenic present in a specimen, thus permitting quantitative estimation. The apparatus (Fig. 15) used for performing this test consists of a glass vessel (flask or Woulf’s bottle) provided with a funnel-tube and delivery-tube (bent at right angles), which is connected with a wider tube, filled with pieces of calcium chloride or plugs of asbestos ; this drying-tube is again connected with a piece of hard glass tube, about one foot long, having a diameter of inch, drawn out at intervals of about 3 inches, so as to reduce its diameter to inch. Hydrogen is generated in the flask by the action of sulphuric acid on zinc, and examined for its purity by heating the glass tube to redness at one of its wide parts Marsh’s apparatus for detection of arsenic. for at least 30 minutes; if no trace of a metallic mirror is formed at the constriction beyond the heated point, the gas and the substances used for its generation may be pronounced free from arsenic. (Both zinc and sulphuric acid often contain arsenic.) After having thus demonstrated the purity of the hydrogen, the suspected liquid, which must contain the arsenic either as oxide or chloride (uot as sulphide), is poured into the flask through the funnel- tube. If arsenic is present in not too small quantities, the gas ignited at the end of the glass tube shows a flame decidedly different from that of burning hydrogen. The flame becomes larger, assumes a bluish tint, and emits an odor of garlic, while above it a white cloud appears which is more or less dense; a cold test-tube held inverted over the 212 METALS AND THEIR COMBINATIONS. flame will be covered upon its walls with a white deposit of minute octahedral crystals of arsenous oxide; a piece of cold porcelain held in the flame becomes coated with a brown stain (arsenic spot) of metallic arsenic. (See explanation above in connection with arsenetted hydrogen.) The glass tube heated, as above mentioned, at one of its wide parts, will show a bluish-black metallic mirror at the constriction beyond. If quantitative determination is desired, the glass tube is heated in two places so as to cause all hydrogen arsenide to be decomposed. To collect, however, the arsenic from any gas that might escape, the end of the tube is inverted and placed into solution of nitrate of silver, which is decomposed by the hydrogen arsenide, silver and arsenous acid being formed. The arsenic solution should be introduced into the hydrogen generator in small por- tions, so as to produce not more hydrogen arsenide at a time than can be decomposed by the method given. The only element which, under the same condi- tions, forms spots and mirrors similar to arsenic, is antimony; there are, however, sufficiently reliable tests to distinguish arsenic spots from those of anti- mony. Arsenic spots treated with solution of hypochlorites (solution of bleaching-powder) dissolve readily ; anti- mony spots are not affected. When nitric acid is added to an arseuic spot, evaporated to dryness and moistened with a drop of silver nitrate, it turns brick-red; antimony spots treated in like manner remain white. Arsenic spots dissolved in ammonium sulphide and evaporated to dryness show a bright yellow, antimony spots an orange-red residue. Fig. 16 represents a simpler form of Marsh’s apparatus, which generally will answer for student’s tests. Fig. 16. Student’s appa- ratus for making arsenic spots. Preparatory treatment of organic matter for arsenic analysis. If organic matter is to be examined for arsenic (or for any other metallic poison) it ought to be treated as follows: The substance, if not liquid, is cut into pieces, -well mashed and mixed with water; the liquid or semi-liquid sub- stance is heated in a porcelain dish over a steam bath with hydrochloric acid and potassium chlorate until the mass has a uniform light-yellow color and has no longer an odor of chlorine. By this operation all poisonous metals (lead and silver excepted, because insoluble silver chloride and possibly insoluble lead sulphate are formed) are rendered soluble even wThen present as sulphides, ANTI MON Y. 213 and may now be separated by filtration from some remaining solid matter. The clear solution is heated and treated with hydrosulphuric acid gas for several hours, when arsenic and all metals of the arsenic and lead groups are precipi- tated as sulphides, a little organic matter also being precipitated generally. The precipitate is collected upon a small filter and treated with warm ammo- nium sulphide, which dissolves the sulphides of arsenic and antimony, leaving behind the sulphides of the lead group, which may be dissolved in nitric, or, if mercury be present, in nitro-hydrochloric acid, and the solution tested by the methods mentioned for the respective metals. The ammonium sulphide solu- tion is evaporated to dryness, this residue mixed with nitrate and carbonate of sodium, and the mixture fused in a small porcelain crucible. By the oxidizing action of the nitrate, both sulphides are converted into the higher oxides, arsenic forming sodium arsenate, antimony forming antimonic oxide. By treating the mass with warm water, sodium arsenate is dissolved and may be filtered off, while antimonic oxide remains undissolved, and may be dissolved in hydrochloric acid. Both solutions may now be used for making the respective tests for arsenic or antimony. Antidotes. Moist, recently prepared ferric hydroxide or dialyzed iron are the best antidotes, insoluble ferric arsenite or arsenate being formed. Vomit- ing should be induced by tickling the fauces or by administering zinc sulphate, but not tartar emetic. 31. ANTIMONY—TIN—GOLD—PLATINUM—MOLYBDENUM. Antimony, Sb = 119.6 (Stibium). This metal is found in nature chiefly as the trisulphide, Sb2S3, an ore which is known as black anti- mony, crude antimony, or stibnite. The metal is obtained from the sulphide by roasting, when it is converted into oxide, which is reduced by charcoal. Antimony is a brittle, bluish-white metal, having a crystalline structure; it fuses at 450° C. (842° F.), and may at a higher temperature be distilled with- Questions.—291. Which metals belong to the arsenic group? what are their characteristics? 292. Which non-metallic elements does arsenic resemble? Mention some of the compounds showing this analogy. 293. How is arsenic obtained in the metallic state ; what are its physical and chemical properties; how does heat act upon it? 294. What is white arsenic? State its compo- sition, mode of manufacture, appearance, solubility, and other properties. 295. Which three solutions, containing arsenic, are officinal, and what is their composition? 296. How is arsenic acid obtained from arsenous oxide, and which arsenate is officinal? 297. State composition and properties of arse- netted hydrogen, and explain its formation. What use is made of it in testing for arsenic? 298. State the composition of realgar, orpiment, Scheele’s green, and Schweinfurth green. 299. Give a detailed description of the process by which arsenic can be detected in organic matter. 300. Describe in detail the principal tests for arsenic. 214 METALS AND THEIR COMBINATIONS out change, provided air is excluded; heated in air it burns bril- liantly. Antimony is used in a number of important alloys, for instance, in type-metal, an'alloy of lead, tin, and antimony. Antimony trisulphide, Antimonii sulphidum, Sb2S3 = 335.2 (Anti- monious sulphide, Sulphide of antimony). The above-mentioned native sulphide, the black antimony, is purified by fusion; it forms steel- gray masses of a metallic lustre, and a striated, crystalline fracture, forming a grayish-black, lustreless powder, which is insoluble in water, but soluble in hydrochloric acid with liberation of hydrogen sulphide. When finely powdered antimonious sulphide is treated with water of ammonia to remove any traces of arsenic (which is frequently found in this ore) and the washed sulphide dried, the purified sulphide of antimony of the U. S. P. is obtained. Antimonious sulphide found in nature is crystallized and steel-gray (Plate V., 6), but it may be obtained also in au amorphous condition as an orange-red (Plate V., 5) powder, by passing hydrogen sulphide through au antimonious solution. By heating the orange-red sul- phide, it is converted into the black variety. Sulphurated antimony, Antimonium sulphuratum (Oxysulphide oj antimony, Kermes mineral), chiefly antimonious sulphide with some antimonious oxide. This preparation is made by boiling purified antimonious sulphide with solution of sodium hydroxide, and adding to the hot solution sulphuric acid as long as a precipitate is formed, which is collected and dried. The sulphides and oxides of antimony, like those of arsenic, com- bine with many metallic sulphides or oxides to form sulpho-salts or oxy-salts. Thus the sodium sulph-antimonite, Na3SbS3, and the sodium antimonite, Na3Sb03, are formed when antimonious sulphide is boiled with sodium hydroxide. Sb2S3 + 6NaOH = Na3SbS3 + Na3Sb03 + 3H20. By the addition of sulphuric acid, both salts are decomposed, sodium sulphate is formed, and antimonious sulphide and oxide are precipitated : 2Na3SbS8 + 3H2S04 = 3Na2S04 + Sb2S3 + 3H2S; 2Na3Sb03 + 3H2S04 = 3Na2S04 + Sb203 + 3H20. It is a reddish-brown, amorphous powder, insoluble in water, solu- ble in hydrochloric acid or sodium hydroxide. ANTIMONY. 215 Experiment 37. Boil about 2 grammes of finely powdered black antimony with a solution of 2 grammes of sodium hydroxide in 80 c.c. of water for about one hour, stirring frequently and occasionally adding water to preserve the same volume. Filter the warm liquid through paper or muslin and add dilute sulphuric acid so long as it produces a precipitate. Collect, wash, and dry the precipitated red powder, which is chiefly amorphous antimonious sulphide with oxide. Antimony pentasulphide, Sb2S5 (Grolden sulpliuret of antimony). A red powder, which, like antimonious sulphide, forms sulpho-salts. It may be obtained by precipitation of acid solutions of antimonic acid by hydrosulphuric acid. Antimonious chloride, SbCl3 (Terchloride of antimony, Butter of antimony). Obtained by boiling the native sulphide with hydro- chloric acid : Sb2S3 + 6HC1 = 3H2« + 2SbCl3. The clear solution is evaporated and the remaining chloride dis- tilled, when it is obtained as a white, crystalline, semi-transparent mass. By passing chlorine over antimonious chloride it is converted into antimonic chloride, SbCl5, which is a fuming liquid. Experiment 38. Boil about 2 grammes of black antimony with 10 c.c. of hydrochloric acid until most of the sulphide is dissolved. Set aside for sub- sidence, pour off the clear solution of antimonious chloride, evaporate to about half its volume and use solution for next experiment. Antimonious oxide, Antimonii oxidum, Sb203 = 287.2 (Oxide of antimony). When antimonious chloride is added to water, decompo- sition takes place, and an oxychloride of antimony, 2SbCl35Sb203, is precipitated : 12SbCl3 + 15H20 = 2SbCI3.5Sb203 + 30HC1. This white precipitate was formerly known as powder of Algaroth. It is completely converted into oxide by treating it with sodium car- bonate : 2SbCl3.5Sb203 + 3Na2C03 = 6Sb203 + 6NaCl + 3C02. The precipitate when washed and dried, is a heavy, grayish-white, tasteless powder, iusoluble in water, soluble in acids. Antimonious oxide, while yet moist, dissolves readily in potassium acid tartrate, forming the double tartrate of potassium and antimony, or tartar emetic, which salt will be more fully considered hereafter. 216 METALS AND THEIR COMBINATIONS. Experiment 39. Pour the antimonious chloride solution (obtained by Ex- periment 38) which should have been boiled sufficiently to expel all hydrogen sulphide, into 100 c.c. of water, wash by decantation the white precipitate of oxychloride thus obtained, and add to it an aqueous solution of about 1 gramme of sodium carbonate. After effervescence ceases, collect the precipi- tate on a filter, wash well and treat some of the precipitate, while yet moist, with a solution of potassium acid tartrate, which dissolves it readily, forming tartar emetic. (For the latter compound see index.) Antidotes. Poisonous doses of any preparation of antimony are generally quickly followed by vomiting: if this, however, have not occurred, the stomach-pump must be applied. Tannic acid in any form, or recently pre- cipitated ferric hydroxide, should be administered. Analytical reactions. (A solution of antimonious chloride, SbCl3> may be used.) 1. Add hydrogen sulphide to an acidified solution of antimony: an orange-red precipitate of antimonious or antimonic sulphide (Sb2S3 or Sb2S5) is produced (Plate V., 5). 2. Add ammonium sulphide to the precipitated sulphide of anti- mony : this is dissolved and may be re-precipitated by neutralizing with an acid. 3. Produce a concentrated solution of antimonious chloride by evaporation or by dissolving the sulphide in hydrochloric acid, and pour it into water : a white precipitate of oxychloride is formed. (See explanation above.) 4 Add sodium hydroxide, ammonium hydroxide, or sodium car- bonate : in either case white antimonious hydroxide, Sb30H, is pro- duced, which is soluble in sodium hydroxide. 5. Boil a piece of metallic copper in the solution of antimonious chloride : a black deposit of antimony is formed upon the copper. By heating the latter in a narrow test-tube, the antimony is volatil- ized and deposited as a white incrustation of antimonious oxide upon the glass. 6. Use Gutzeit’s or Marsh’s test as described under analytical reac- tions for arsenic. Tin, Sn = 118.8 (Stannum). This metal is found in nature chiefly as stannic oxide or tin-stone, Sn02, from which the metal is easily obtained by heating with coal : Sn02 + 2C = Sn + 2C0. TIN—GOLD. 217 Tin is an almost silver-white, very malleable metal, fusing at the comparatively low temperature of 228° C. (440° F.). It is used in many alloys, in the silvering of looking-glasses by tin-amalgam, and chiefly in the manufacture of tin-plate, which is sheet-iron covered with a thin layer of tin. Tin is bivalent in some compounds, quadrivalent in others. These combinations are distinguished as stannous and stannic compounds. Stannous chloride, SnCl2 (Protochloride of tin). Obtained by dis- solving tin in hydrochloric acid by the aid of heat: Sn + 2HC1 = SnCI2 + 2H. Sufficiently evaporated, the solution yields crystals of the composi- tion SnCl2.2H20. Stannous chloride is a strong deoxidizing agent, frequently used as a reagent for arsenic, mercury, and gold, which metals are precipitated from their solutions in the metallic state. It is used also in calico-printing. Stannic chloride, SnCl4 (Perchloride of tin). Stannous chloride may be converted into stannic chloride either by passing chlorine through its solution or by heatiug with hydrochloric and nitric acids. Analytical reactions. (Stannous chloride, SnCl2, and stannic chloride, SnCl4, may be used.) 1. Add hydrogen sulphide to solution of a stannous salt: brown stannous sulphide is precipitated (Plate V., 7): SnCl2 + H2S = 2HC1 + SnS. The precipitate is soluble in ammonium sulphide. 2. Add hydrosulphuric acid to a solution of a stannic salt: yellow stannic sulphide is precipitated (Plate V., 8): SnCl2 + 2H2S = 4HC1 + SnS2, The precipitate is soluble in ammonium sulphide. 3. Sodium or potassium hydroxide added to a stannous salt, pro- duces a white precipitate of stannous hydroxide, Sn20H. • The same reagents added to a stannic salt produce white stannic acid, H2Sn03. Both precipitates are soluble in excess of the alkali. Gold, Au = 196.7 (Aurum). Gold occurs in nature chiefly in the free state, often associated with silver, copper, and possibly with other metals. This impure gold is separated from most of the adhering 218 METALS AND THEIR COMBINATIONS. sand and rock by a mechanical process of washing, in which advan- tage is taken of the high specific gravity of the metallic masses. The remaining mixture of heavy material is treated with mercury, which dissolves gold and silver, leaving behind most other impurities. The gold amalgam is placed in a retort and heated, when the mercury distils over, while the gold is left behind. If this should contain silver, the metals may be separated by treating the alloy with hot sulphuric acid, which dissolves silver, leaving gold behind. Pure gold is too soft for general use, and therefore is alloyed with various proportions of silver and copper. American coin is an alloy of 90 parts of gold and 10 parts of copper; jeweller’s gold contains generally 75 per cent. (18 carat) of gold, the other 25 per cent, being copper and silver; the varying proportions are well indicated by the color. Gold is not affected by either hydrochloric, nitric, or sulphuric acid, but is dissolved by nitro-hydrochloric acid, by free chlorine or bromine, and by mercury, with which it forms an amalgam. Gold is trivalent generally, as in auric chloride, AuCls, but most likely also univalent in some compounds, as in aurous chloride, AuCl. Auric chloride, AuC13. Obtained by dissolving pure gold in nitro- hydrochloric acid and evaporating the solution to dryness. A mixture of equal parts of auric chloride and sodium chloride is officinal under the name of chloride of gold and sodium. It is an orange-yellow, very soluble powder. Analytical reactions. (Auric chloride, AuC13, may be used.) 1. Add hydrogen sulphide to solution of gold: brown auric sul- phide, Au2S3, is precipitated, which is soluble in yellow ammonium sulphide. 2. Add ferrous sulphate to solution of gold and set aside for a few hours; metallic gold is precipitated as a dark powder, which, by fusion, is converted into a metallic mass. 3. Many other reagents cause the separation of metallic gold from its solution, as, for instance, oxalic acid, sulphurous and arsenous acids, potassium nitrite, etc. Platinum, Pt = 194.3. Platinum, like gold, is found in nature in the free state, the chief supply being derived from the Ural moun- PLATINUM—MOLYBDENUM. 219 tains, where it is found associated with a number of metals (iridium, ruthenium, osmium, palladium, rhodium) resembling platinum in their properties. Platinum is of great importance and value on account of its high fusing-point and its resistance to the action of most chemical agents, for which reason it is used in the manufacture of vessels serving in chemical operations. Platinum, when dissolved in nitrohydrochloric acid, forms platinic chloride, PtCl4, a salt frequently used as a reagent for potassium or ammonium salts, with which it forms insoluble double chlorides of the composition PtCl4.2KCl and PtCl4.2NH4Cl. By heating the latter salt sufficiently it is decomposed and metallic platinum is left as a gray spongy mass. Molybdennm, Mo = 95.9. This metal is found in nature chiefly as sulphide, MoS2, from which, by roasting, molybdic oxide, Mo03, is obtained. The oxide, when dissolved in water, forms an acid which combines with metals, forming a series of salts termed molybdates. Of interest is ammonium molybdate, a solution of which in nitric acid is an excellent reagent for phosphoric acid, with which it forms a yellow precipitate, insoluble in acids, soluble in ammonium hydroxide. Questions.—301. How is antimony found in nature, and what are the prop- erties of this metal ? 302. State the composition of antimonious sulphide, and its color when crystallized and amorphous. 303. How do hydrochloric acid and alkali hydroxides act upon antimonious sulphide? 304. What is the sul- phurated antimony of the U. S. P. ? 305. Mention the two chlorides of anti- mony and state their properties. 306. How is antimonious oxide made, and what is it used for? 307. Give tests for antimony. 308. State the use made of tin in the metallic state; mention the two chlorides of tin, and what stannous chloride is used for. 309. How are gold and platinum found in nature; by what acid may they be dissolved, and what is the composition of the com- pounds formed? 310. Which is the most important compound of molybdenum, and what is it used for? 220 METALS AND THEIR COMBINATIONS. Summary of analytical characters of metals of the arsenic group. Arsenic. Antimony. Tin. Gold. Platinum. Hydrogen sulphide . Yellow pre- cipitate. Orange precipitate. Yellow or brown precipitate. Black precipitate. Dark- brown precipitate. Precipitate heated 1 in strong hydro- j- Insoluble. Soluble. Soluble. Insoluble. Insoluble. chloric acid. j Potassium hydroxide White precipitate. White precipitate. Brownish precipitate, soluble in excess. With ex- cess of hydro- chloric acid a Ammonia water . . Gutzeit’s test . . . Fleitmann’s test . . Yellow stain, turning dark with water. Dark stain. White precipitate. Dark stain. White precipitate. Brownish- yellow precipitate. - yellow precipi- tate. V. ANALYTICAL CHEMISTRY. 32. INTRODUCTORY REMARKS AND PRELIMINARY EXAMINATION. General remarks. Analytical chemistry is that part of chemistry which treats of the different analytical methods by which substances are recognized and their chemical composition determined. This determination may be either qualitative or quantitative, and, accord- ingly, a distinction is made between a qualitative analysis, by which simply the nature of the elements (or groups of elements) present in the substance under examination is determined, and a quantitative analysis, by which also the exact amount of these elements is ascer- tained. In this book qualitative analysis will be considered chiefly, as the methods for quantitative determinations of the different elements are so numerous and so varied that a detailed description of them would occupy more space than can be devoted to analytical chemistry in this work. Some brief directions concerning quantitative determinations, especially by volumetric methods, are given in Chapter 37. Every- one studying analytical chemistry should do it practically, that is, should perform for himself in a laboratory all those reactions which have been mentioned heretofore as characteristic of the different ele- ments and their compounds, and, furthermore, should make himself acquainted with the methods by which substances are recognized when mixed with others, by analyzing various complex substances. Such a course of practical work in a practical laboratory is of the greatest advantage to all studying chemistry, and students cannot be too strongly advised to avail themselves of any facilities offered in performing chemical experiments, analytically or otherwise. Apparatus needed for qualitative analysis. 1. Iron stand. (Fig. 17.) 2. Bunsen lamp with flexible tube (Fig. 17), or (where without gas-supply) spirit- lamp and alcohol. 222 ANALYTICAL CHEMISTRY. 3. Test-tube stand and one dozen assorted test-tubes. (Fig. 18.) 4. Three small beakers from 100 to 150 c.c. capacity. (Fig. 19, A.) 5. Two flasks of 100 to 150 c.c. capacity. (Fig. 19, B.) Fig. 17. Fig. 18. Fig. 19. 6. Wash-bottle of about 400 c.c. capacity. (Pig. 20, A.) 7. Three small’glass funnels, about one and a half to two inches in diameter. (Fig. 20, B.) INTRODUCTORY REMARKS 223 8. A few pieces of glass tubing about ten inches long, and some India-rubber tubing to fit the glass tubing. 9. Three glass rods. Fig. 20. 10. Three small porcelain evaporating dishes, about two inches in diameter. (Fig. 21, A.) 11. Blowpipe. (Fig. 21, B.) 12. Crucible tongs. (Fig. 21, C.) Fig. 21. 13. Round and triangular file. 14. Wire gauze, about six inches square, or sand tray. 15. One square inch of platinum foil (not too light), and six inches of platinum wire. 16. Filter-paper. 17. Pair of scissors. 18. One or two dozen assorted corks. 19. Sponge and towel. 20. Two watch-glasses. 21. Small pestle and mortar. (Fig. 21, D.) 22. Small porcelain crucible. 23. Small platinum crucible. (Fig. 21, E.) 24. Wire triangle to support the crucible. (Fig. 21, F.) 224 ANALYTICAL CHEMISTRY Reagents needed in qualitative analysis. a. Liquids. 1. Sulphuric acid, sp. gr. 1.84, H2S04. 2. Sulphuric acid diluted, sp. gr. 1.068 (1 part sulphuric acid, 9 parts water). 3. Hydrochloric acid, sp. gr. 1.16, HC1. 4. Hydrochloric acid diluted, sp. gr. 1.049 (6 parts hydrochloric acid, 13 parts water). 5. Nitric acid, sp. gr. 1.42, HN03. 6. Acetic acid, sp. gr. 1.048, C2H402. 7. Hydrogen sulphide, either the gas or its solution in water, H2S. 8. Ammonium sulphide, (NH4)2S. 9. Ammonium hydroxide (water of ammonia), NH4OH. 10. Ammonium carbonate (NH4)2C03. A solution of one part of the commercial salt in a mixture of four parts of water and one part of water of ammonia. 11. Ammonium chloride, NH4C1; ten per cent, solution. 12. Ammonium oxalate, (NH4)2C204; five percent, solution. 13. Ammonium molybdate, (NH4)2Mo04. A five per cent, solution of the salt in a mixture of equal parts of water and nitric acid. 14. Sodium hydroxide, ISTaOH. 15. Sodium carbonate, Na2C03. 16. Sodium phosphate, Na2HP04. 17. Sodium acetate, NaC2H302. 18. Potassium chromate, K2Cr04. 19. Potassium dichromate, K2Cr2Q7. Ten per cent, solutions. 20. Potassium iodide, KI. 21. Potassium ferrocyanide, K4Fe(CN)6. 22. Potassium ferricyanide, K6Fe2(CN)12. 23. Potassium sulphocyanate, KCNS. Five per cent, solutions. 24. Magnesium sulphate, MgS04. 25. Barium chloride, BaCl.2 26. Calcium chloride, CaCl2. Ten per cent, solutions. 27. Calcium hydroxide, Ca(OH)2 (lime-water). no „ , . i u j. n 28. Calcium sulphate, CaS04. Saturated solutions. 29. Ferric chloride, Fe2016. 30. Lead acetate, Pb.(C2H302)2. 31. Silver nitrate, AgN03. 32. Mercuric chloride, HgCl2. 33. Platinic chloride, PtCl4. Five per cent, solutions. 34. Stannous chloride, Sn012.2H20 ; ten per cent, solution. 35. Solution of indigo. 36. Alcohol, C2H5OH. 37. Sodium cobaltic nitrite solution, Co2(N02)6.6NaN02 -f- H20. Four grammes of cobaltous nitrate, Co(N03)2.6H20, and 10 grammes of sodium nitrite, NaN02, are dissolved in about 50 c.c. of water, 2 c.c. of acetic acid are added, and then water to make 100 c.c. 38. Alkaline mercuric-potassium iodide solution (Nessler’s solution). Five grammes of potassium iodide are dissolved in hot water, and to this is added a hot solution, made by dissolving 2.5 grammes of mercuric chloride in 10 c.c. of water. To the turbid red mixture is added a solution made by dis- INTRODUCTORY REMARKS. 225 solving 16 grammes of potassium hydroxide in 40 c.c. of water, and the whole diluted to 100 c.c. Some mercuric iodide deposits on cooling, and may be left in the bottle, the clear solution being decanted as needed. b. Solids. 1. Litmus or blue and red paper. 2. Turmeric paper. 3. Sodium carbonate, dried, Na2C03. 4. Sodium biborate, borax, Na2Bo4O7.10H2O. 5. Sodium-ammonium-hydrogen phosphate (microcosmic salt), Na(NH4)HP04.4H20. 6. Potassium carbonate, K2C03. 7. Potassium nitrate, KN03. 8. Potassium chlorate, KC103. 9 Potassium permanganate, KMn04. 10. Potassium cyanide, KCN. 11. Calcium hydroxide, Ca(OH)2. 12. Ferrous sulphide, FeS. 13. Ferrous sulphate, FeS04.7II20. 14. Manganese dioxide, Mn02. 15. Zinc, granulated, Zn. 16. Copper, Cu. 17. Cupric oxide, CuO. 18. Cupric sulphate, CuS04.5H20. 19. Tartaric acid, H2C4H406. 20. Tannic acid, H.C14H909. 21. Pyrogallic acid, C6H3(OH)3. 22. Diphenylamine, (CfiH,),NH. 23. Starch, C6H10O5. While the apparatus and reagents here enumerated are the more important ones, the analyst will occasionally require others not men- tioned in the above list. General mode of proceeding in qualitative analysis. Every step taken in analysis should be properly written down in a note-book, and these remarks should be made directly after a reaction has been performed, and not after the nature of the substance has been revealed by perhaps numerous reactions. Not only the reactions by which positive results have been obtained should be noted, but also those tests and reagents mentioned which have been applied with negative results—that is, which have been applied without revealing the presence of any substance, or any group of substances. Such negative results are, however, positive in so far as they prove the absence of a certain substance, or certain substances, for which reason they are of direct value, and should be noted. In comparing, finally, the result obtained by the analysis with the 226 ANALYTICAL CHEMISTRY notes taken during the examination, none of them should be contra- dictory to the conclusions drawn. If, for instance, the preliminary examination showed the substance to have been volatilized by heating upon platinum foil with the exception of a very slight residue, and if, afterward, other tests show the presence of ammonia and hydrochloric acid and the absence of everything else, and if, then, the conclusion be drawn that the substance is pure ammonium chloride, this conclu- sion must be incorrect, because pure ammonium chloride is wholly volatile, and does not leave a residue. It will then be the task of the operator to find where the mistake occurred, and to correct it. Use of reagents. A mistake made by most beginners in analyzing is the use of too large quantities both of the substance applied for testing and of the reagents added. This excessive use of material is not only a waste of money, but, what is of greater importance, a waste of time. Some experience in analyzing will soon convince the student of the truth contained in this remark, and will also enable him to select the correct quantities of materials to be used, which rarely exceed 0.2-1.0 gramme. A smaller amount frequently may answer, and a much larger quantity may occasionally be needed, as, for instance, in cases where highly diluted reagents, such as calcium sulphate solution, lime-water, hydrogen sulphide water, etc., are applied. Preliminary examination. This examination includes the follow- ing points : 1. Physical properties. Solid or liquid; crystallized or amor- phous ; color, odor, hardness, gravity, etc. (On account of possible poisouous properties, the greatest care should be exercised in tasting a substance.) 2. Action on litmus. Examined by holding litmus-paper in the liquid, or by placing the powdered solid upon red and blue litmus- paper, moistened with water. (It should be remembered that many normal salts, as, for instance, aluminium sulphate, ferrous sulphate, etc., have au acid reaction to litmus-paper, and that such a reaction consequently is not conclusive of the presence of a free acid, nor even of an acid salt.) 3. Heating- on platinum foil or in a dry glass tube, open at both ends. (If the substance to be examined be a liquid, it should be evaporated in a small porcelain dish to see whether a solid residue be INTRODUCTORY REMARKS. 227 left or not. If a residue be left, it should be treated like a solid.) The heating of a small quantity of a solid substance upon platinum foil held over the flame of a Bunsen burner or of an alcohol lamp, is a test which should never be omitted, as it discloses in most cases the fact whether the substance is of an organic or inorganic nature. Most organic (nori-volatile) substances, when thus heated, will burn with a luminous flame, leaving in many cases a black residue of carbon, which, upon further heating, disappears. In cases where the organic nature of a compound is not clearly demonstrated by heating on plati- num foil, the substance is heated with an excess of cupric oxide in a test-tube or other glass tube, provided with a delivery-tube, which passes into lime-water. Upon heating the mixture, the carbon of the organic matter is converted into carbon dioxide, which renders lime- water turbid. The analytical processes by which the nature of an organic sub- stance is determined, are not considered in this part of the book, but will be mentioned when considering the carbon compounds. An inorganic substance, heated on platinum foil, may either be volatilized, fused, change color, become oxidized, suffer decomposition, or remain unchanged. (See Table I., page 231.) Fig. 22. Fig. 23. Heating of solids in bent glass tube Heating on charcoal by means of blowpipe. Some substances, containing small quantities of water enclosed between the crystals (common salt, for instance), decrepitate when heated, the small fragments being thrown from the foil; such sub- stances should be heated in a dry test-tube to expel the water and then be examined on platinum foil. 228 ANALYTICAL CHEMISTRY. In many cases it is preferable to heat the substance in a bent glass tube, as shown in Fig. 22, instead of on platinum foil, because vola- tile products evolved during the process of heating may become re- condensed in the cooler part of the tube, and thus saved for further examination. The presence of water, sulphur, mercury, arsenic, etc., may often be readily demonstrated by this mode of operating. 4. Heating on charcoal by means of the blowpipe. This test reveals the presence of chlorates and nitrates by the vivid combustion of the charcoal (known as deflagration), which takes place in conse- quence of the oxidizing action of these substances. Arsenic is indicated by a characteristic odor of garlic. 5. Heating on charcoal with sodium carbonate and potassium cyanide. A small quantity of the finely powdered substance is mixed with twice its weight of potassium cyanide and dry sodium carbonate. This mixture is placed in a small hole made in a piece of charcoal, and heat applied by means of the blowpipe (see Fig. 23). Many metallic compounds may be recognized by this test, the metals being liberated and found as metallic globules or shining particles in the fused mass after this has been removed from the charcoal and washed with water in a small mortar. (See Fig. 24.) Fig. 24. A characteristic incrustation is formed by some metals, due to the precipitation of a metallic oxide around the heated spot on the char- coal. If sulphur as such, or in any form of combination, be present in the substance examined by this test, the fused mass contains a sulphide of the alkali (hepar), which may be recognized by placing it on a piece of bright silver (coin) moistened with a drop of water, when the silver will be stained black in consequence of the formation of silver sul- INTRODUCTORY REMARKS. 229 phide. The presence of the alkali sulphide may also be demonstrated by the addition of a few drops of hydrochloric acid to the fused mass, when hydrogen sulphide is evolved and may be recognized by its odor. 6. Flame tests. Many substances impart a characteristic color to a non-luminous flame. The best mode of performing this test is as follows : A platinum wire is cleaned by washing in hydrochloric acid and water, and heating it in the flame until the latter is no longer colored. One end of the wire is fused in a short piece of glass tubing (see Fig. 25), the other end is bent so as to form a small loop, which Fig. 25 is heated, dipped into the substance to be examined, and again held in the lower part of the flame, which then becomes colored. Some substances show the color-test after being moistened with hydrochloric or sulphuric acid. A second method of showing flame reactions is to mix the substance with alcohol in a small dish; the alcohol, upon being ignited, shows a colored flame, especially in the dark. 7. Colored borax beads. The compounds of some metals when fused with glass, impart to it characteristic colors. For analytical purposes not the silica-glass, but borax-glass is generally used. This latter is made by dipping the loop of a platinum wire in powdered borax and heating it in the flame (directly, or by means of the blow- pipe) until all water has been expelled and a colorless, transparent bead has been formed. To this colorless bead a little of the finely powdered substance is added and the bead strongly heated. The metallic compound is chemically acted upon by the boric acid, a borate being formed which colors the bead more or less intensely, according to the quantity of the metallic compound used. Some metals (copper, for instance) forming two series of compounds, give different colors to the bead when present in either the higher or lower state of oxidation. By modifying the blowpipe flame so as either to oxidize (by supply- ing an excess of atmospheric oxygen) or deoxidize (by allowing some unburnt carbon to remain in the flame), the metallic compound in the bead may be made to assume the higher or lower state of oxidation. 230 ANALYTICAL CHEMISTRY. A copper bead may thus be changed from blue to red or red to blue, the blue bead containing the copper in the cupric, the red bead in the cuprous form. In some cases microcosmic salt, NaNH4HP04, is used for making the bead. 8. Liquefaction of solid substances. Most solid substances have to be dissolved for analysis. The solution obtained may be either a simple or chemical solution. In a simple solution the dissolved body retains all of its original properties, with the exception of its shape, and may be re-obtained by evaporation. Sodium chloride and sugar dissolved in water form simple solutions. A chemical solution is one in which the chemical composition of the substance has been changed during the process of dissolving, as, for instance, when calcium carbo- nate is dissolved in hydrochloric acid ; this solution now contains and leaves on evaporation calcium chloride. The solvents used are water, or the mineral acids for substances insoluble in water, especially dilute, or, if necessary, strong hydrochloric acid. The dissolving action of the acid should be facilitated by the aid of heat. Nitric or even nitro- hydrochloric acid may have to be used in some cases. Three mistakes are frequently made by beginners in dissolving substances in acids, viz.: The substance is not powdered as finely as it should be; sufficient time is not given for the acid to act; too large an excess of the acid is used. If a substance is partly dissolved by water and partly by one or more other solvents, it may be well to examine the different solutions separately. Substances insoluble in water and in acids have to be rendered soluble by fusion with a mixture of potassium and sodium carbonate, or with potassium acid sulphate, or by the action of hydrofluoric acid. The insoluble sulphates of the alkaline earths, when fused with the alkaline carbonates, are converted into carbonates, while the sulphates of the alkalies are formed. The latter compounds may be eliminated by washing the fused mass with water and filtering : the solid residue upon the filter contains the carbonates of the alkaline earths, which may be dissolved in hydrochloric acid. Insoluble silicates may be decomposed by the methods mentioned on page 98. Questions.—311. What is analytical chemistry, and what is the object of qualitative and of quantitative analysis? 212. What properties of a substance should be noticed first in making a qualitative analysis ? 313. By what tests may organic compounds be distinguished from inorganic compounds? 314. Ex- INTRODUCTORY REMARKS. 231 Heat the solid sub- stance upon plati- num foil, or in a dry, narrow glass tube open at both ends. Combustible are: All organic compounds, carbon, sulphur, phosphorus, etc. Easily volatilized are: All compounds of ammonium and mercury, most of arsenic, some of antimony, etc. (Heat in a glass tube as directed on page 226.) Fusible are: Most of the salts of the alkalies, and some of those of the alkaline earths, many metals, etc. Infusible are: Salts of the earths, and most salts of the alkaline earths and heavy metals, most silicates, etc. Assume a darker color: Many oxides of the heavy metals and their salts (oxides of zinc, antimony, lead, etc.). Evolve water: Many salts containing water of crystallization some hydroxides, etc. Decrepitate: Some salts, sodium chloride, for instance. Heat the solid sub- stance on charcoal. Deflagrate: Nitrates, chlorates, iodates, bromates, etc. Give garlic odor: Most compounds of arsenic. Heat the substance, mixed with sodium carbonate and po- tassium cyanide, on charcoal. Give hepar: Sulphur and all its compounds Give bright metallic grains without incrustation: Compounds of gold, silver, copper, tin. Give bright metallic grains with incrustation : Compounds of lead, bismuth, antimony. Give gray infusible powder: Compounds of iron, cobalt, nickel, platinum. Heat the substance on platinum wire in a non luminous flame Yellow flame, compounds of sodium. Violet flame, compounds of potassium. Crimson flame, compounds of lithium or strontium. Orange flame, compounds of calcium. Yellowish-green flame, compounds of barium or molybdenum. Green flame, compounds of copper, phosphoric or boric acids. Blue flame, compounds of arsenic, antimony, lead, or cupric chloride. Heat a colorless borax bead with very little of the substance. Green bead, compounds of chromium. Blue bead, compounds of cobalt or copper in the oxidizing flame. Red bead, compounds of copper in the reducing flame. Violet bead, compounds of manganese. Yellow to brown bead, compounds of iron. Colorless bead, compounds of the light metals and those of the arsenic group; also silver, bismuth, lead, etc. Table I—Preliminary examination. 232 ANALYTICAL CHEMISTRY. 33. SEPARATION OF METALS INTO DIFFERENT GROUPS. General remarks. The preliminary examination will, in most cases, decide whether or not a metal or metals are present in the substance to be examined. If there be metals, the solution should be treated according to Table II., page 235, in order to find the group or groups to which these metals belong, aud also to separate them into these groups, the individual nature of the metals themselves being afterward demonstrated by special methods. The simplest method of separating from each other the 55 metals known, when all in one solution, would be to add successively 55 different reagents, each of which should form an insoluble compound with but one of the metals. By separating this insoluble compound from the metals remaining in solution (by filtration), and by thus pre- cipitating one metal after the other, they all could be easily separated. We have, however, no such 55 reagents, and are, consequently, com- pelled to precipitate a number of metals together, and the reagents used for this purpose are known as group-reagents. They are: 1. Hydrogen sulphide, added to the solution previously acidified by hydrochloric acid. Precipitated are: the metals of the arsenic and lead groups as sulphides. 2. Ammonium, sulphide, added after supersaturating with ammonium hydroxide. Precipitated are : the metals of the iron group and of the earths as sulphides or hydroxides. 3. Ammonium carbonate. Precipitated are: the metals of the al- kaline earths as carbonates. 4. In solution are left: the metals of the alkalies and magnesium. The order in which these group-reagents are added cannot be reversed or changed, because ammonium sulphide added first would precipitate not only the metals of the iron group and the earths, but plain the terms decrepitation and deflagration. 315. Mention some substances which are completely volatilized by heat, some which are fusible, and some which are not changed by heating them. 316. What is meant by “ hepar,” and which element is indicated by the formation of hepar? 317. Mention some metals which may be liberated from their compounds by heating on charcoal with potassium cyanide and carbonate. 318. Which metallic compounds and which acids are capable of coloring a non-luminous flame? Name the colors imparted. 319. State the metals which impart characteristic colors to a borax bead. 320. Which solvents are used for liquefying solids, and what precau- tions should be observed in this operation ? SEPARATION OF METALS INTO DIFFERENT GROUPS. 233 also the metals of the lead group; ammonium carbonate would pre- cipitate also most of the heavy metals. For the same reasons, in separating metals of the diiferent groups, the group- reagents must be added in excess, that is, enough of them must be added to precipitate the total quantity of the metals of one group, before it is possible to test for metals of the next group. Suppose, for instance, a solution to con- tain a salt of bismuth only. Upon the addition of hydrogen sulphide to the acidified solution, a dark-brown precipitate (of bismuth sulphide) is produced, indicating the presence of a metal of the lead group. Suppose, further, that hydrogen sulphide has not been added in sufficient quantity to precipitate the whole of the bismuth, then ammonium sulphide, as the next group reagent, would produce a further precipitation in the filtrate, which fact would lead to the assumption that a metal of the iron group was present, which, however, would not be the case. If the solution contain but one metal, the group-reagents are added successively in small quantities to the same solution, until the reagent is found which causes a precipitation, which reagent is then added in somewhat larger quantity in order to produce a sufficient amount of the precipitate for further examination. Acidifying the solution. Hydrosulphuric acid has to be added to the acidified solution for two reasons, viz.: In a neutral or alkaline solution some metals of the arsenic group (which are to be precipitated) would not be precipitated by hydrogen sulphide; some of the metals of the iron group (which are not to be precipitated) would be thrown down. The best acid to be used in acidifying is dilute hydrochloric acid; but this acid forms insoluble compounds with a few of the metals of the lead group, causing them to be precipitated. Completely precipi- tated by hydrochloric acid are mercurous and silver compounds; par- tially precipitated are compounds of lead, chloride of lead being some- what soluble in water. The precipitate formed by hydrochloric acid may be examined by Table III., page 237. Hydrochloric acid added to a solution may, in a few cases (other than those just mentioned), cause a precipitate, as, for instance, when added to solutions containing certain compounds of antimony or bis- muth (the precipitated oxychlorides of these metals are soluble in excess of the acid), to metallic oxides or hydroxides which have been dissolved by alkali hydroxides (for instance, hydroxide of zinc dis- solved in potassium or ammonium hydroxide), to solutions of alkali silicates, when silica separates, etc. Addition of hydrosulphuric acid. This reagent is employed either in the gaseous state (by passing it through the heated solution) or as 234 ANALYTICAL CHEMISTRY. hydrogen sulphide water. The latter reagent answers in those cases where but one metal is present; if, however, metals of the arsenic and lead groups are to be separated from metals of other groups, the gas must be used. Fig. 26. Fig. 27. Apparatus for generating hydro- sulphuric acid. Apparatus for generating hydro- sulphuric acid. For generating hydrosulphuric acid the directions given on page 107 may be followed. In place of the apparatus there mentioned for generating the gas, others may be used which have the advantage to the analyst that the supply of gas may be better regulated. Fig. 26 shows such an apparatus for the continuous preparation of the gas. It consists of three glass bulbs; the upper bulb, prolonged by a tube reaching to the bottom of the lowest one, is ground air-tight into the neck of the second. Ferrous sulphide is introduced into the middle bulb through the tubulure, which is then closed by a per- forated cork through which connection is made with the wash bottle. Acid poured in through the safety tube, runs into the bottom globe and rises to the ferrous sulphide in the second bulb. Upon closing the delivery tube, the pressure of the generated gas forces the liquid from the second bulb through the lower to the upper, thus preventing contact of acid and ferrous sulphide until the gas is used again. A convenient and cheaper apparatus is shown in Fig. 27. A glass tube, drawn at its lower end to a small point and partly filled with pieces of ferrous sulphide, is suspended through a cork (not air-tight) in a cylinder containing the acid. The gas supply is regulated by closing or opening the stop-cock, and also by raising or lowering the tube in the acid. SEPARATION OF METALS INTO DIFFERENT GROUPS. 235 Add the following reagents successively to the same solution. Every time a precipitate is formed an excess of the precipitant is to be used; the precipitate is collected upon a filter, well washed, and treated by the tables mentioned. To the clear filtrate the next group-reagent is added. If the solution contains but one metal, generally it is sufficient to find by means of this table the group to which it belongs, and then to use the original solution for testing according to Tables III.-VIII. Dilute hydrochloric acid precipitates: Hydrogen sulphide precipitates: Ammonium hydrox- ide and sulphide precipitate: Ammonium car- bonate precipitates: In solution are left: Metals of the lead group. Arsenic group. Iron group and earths. Alkaline earths. Alkalies and mag- nesium. Silver chloride, ] 1 Lead chloride, A precipitate may be caused by other sub- stances than those men tioned. See page 233. See Table III. Insoluble in ammo- nium sulphide. Soluble in ammo- nium sulphide. Ferrous sulphide, black. Cobaltous sulphide, black. Niekelous sulphide, black. Manganous sulphide, flesh-colored. Zinc sulphide, white. Chromic hydroxide, green. Aluminium hydroxide, white A precipitate may be caused by other substances than those mentioned. See page 236. See Table VI. Calcium 'j carbonate, Barium carbonate, >- Strontium carbonate, J See Table VII. Magnesium. Potassium. Sodium. Lithium. Ammonium. See Table VIII. Lead sulphide, Mercuric sul- phide, Bismuth sul- phide, Cupric sul- phide, Cadmium sulpt yellow. See Table IV. M LJ w tide, Arsenous sulphide, yellow. Antimonious sul- phide, orange. Stannous sulphide, brown. Stannic sulphide, yellow. Auric sulphide, brown. Platinic sulphide, brown. See Table V. Table II.—Separation of metals into different groups. 236 ANALYTICAL CHEMISTRY. In some cases sulphur is precipitated on the addition of hydrogen sulphide, while a change in color may take place. This change is due to the deoxidizing action of hydrogen sulphide, the hydrogen of this reagent becoming oxidized and converted into water, while sul- phur is liberated. Thus, brown ferric compounds are converted into pale-green ferrous compounds; red solutions of acid chromates become green; and red permanganates or green manganates are decol- orized. The same deoxidizing action of hydrogen sulphide is the reason why this reagent cannot be employed in a solution containing free nitric acid, which latter compound oxidizes the hydrogen sulphide. Separation of the metals of the arsenic from those of the lead group. The precipitate produced by hydrogen sulphide in acid solu- tion contains the metals of the arsenic and lead groups. They are separated by means of ammonium sulphide, which dissolves the sul- phides of the arsenic group, but does not act on those of the lead group. Addition of ammonium sulphide. This reagent should never be added to the acid solution, but the solution should be previously supersaturated by ammonium hydroxide, as, otherwise, a precipitate of sulphur may be formed. The yellow ammonium sulphide is almost invariably a polysulphide of ammonium, that is, ammonium sulphide which has combined with one or more atoms of sulphur. If an acid be added to this compound, an ammonium salt is formed, hydrogen sulphide is liberated, and sulphur precipitated : (NH4)2S2 + 2HC1 = 2NH4C1 + H2S + S. Ammonium sulphide precipitates the metals of the iron group as sulphides, with the exception of chromium, which is precipitated as hydroxide; aluminium is precipitated in the same form of combina- tion. Ammonium sulphide (or ammonium hydroxide) causes also the precipitation of metallic salts which have been dissolved in acids, as, for instance, of the phosphates, borates, silicates, or oxalates of the alkaline earths, magnesium, and others. The processes by which the nature of some of these precipitates is to be recognized are found in Table VI., page 239. Addition of ammonium carbonate. The reagent used is the com- mercial salt, dissolved in water, to which some ammonia water has SEPARATION OF THE METALS OF EACH GROUP. 237 been added. Heating facilitates complete precipitation of the carbon- ates of the alkaline earths. 34. SEPARATION OF THE METALS OF EACH GROUP. Table III.—Treatment of the precipitate formed by hydrochloric acid. The precipitate may contain silver, mercurous, and lead chlorides. Boil the washed precipitate with much water, and filter while hot. Filtrate may contain lead chloride. Add dilute sulphuric acid; a white precipitate of lead sul- phate is produced. Eesidue may consist of mercurous and silver chlor- ides Digest residue with ammonia water. Solution may contain sil- ver. Neutralize with nitric acid, when silver chloride is reprecipi- tated. A dark gray residue indi- cates mercury, the white mercurous chloride hav- ing been converted into dimercurous ammonium chloride. Treatment of the precipitate formed by hydrogen sulphide in warm acid solution. The precipitate is collected upon a small filter, well washed with water, and then examined for its solubility in ammonium sulphide. This is done by placing a portion of the washed precipi- tate in a test-tube, adding ammonium sulphide, and warming gently. It is either wholly insoluble (metals of the lead group), and treated according to Table IV., or fully soluble (metals of the arsenic group), and treated according to Table V., or it is partly soluble and partly insoluble (metals of both groups). In the latter case, the total quan- tity of the washed precipitate is to be treated with warm ammonium sulphide; upon filtering, an insoluble residue is left, which is treated according to Table IV.; to the filtrate, diluted sulphuric acid is added Questions.—321. State the three groups of heavy, and the three groups of light metals. 322. By which two reagents may all heavy metals be precipi- tated ? 323. Why is a solution acidified before the addition of hydrogen sul- phide, when testing for metals? 324. Which metals are precipitated by hydrochloric acid? 325. Which two groups of metals are precipitated by hydrogen sulphide in acid solution? 326. How are the sulphides of the arsenic group separated from those of the lead group ? 327. Why is an acid solution neutralized or supersaturated by ammonium hydroxide, before adding ammonium sulphide? 328. Which two groups of metals are precipitated by ammonium sulphide, and in what forms of combination? 229. Name the group-reagent for the alkaline earths. 330. Which metals maybe left in solu- tion after hydrogen sulphide, ammonium sulphide, and ammonium carbonate have been added ? 238 ANALYTICAL CHEMISTRY. as long as a precipitate is formed, which precipitate contains the metals of the arsenic group as sulphides, generally with some sulphur from the ammonium sulphide. Table IV.—Treatment of that portion of the hydrogen sulphide precipitate which is insoluble in ammonium sulphide. The precipitate may contain the sulphides of lead, copper, mercury, bismuth, and cadmium. Heat the well-washed precipitate with nitric acid in a test-tube, and filter. Residue may con- sist of: Mercuric sulph- ide, which is black and easily dissolves in nitro- hydrochloric acid, which solution, after sufficient evaporation, is tested by potas- sium iodide, etc. Lead sulphate is white, pulveru- lent, and soluble in ammonium tar- trate. Sulphur is yellow and combustible. Filtrate may contain the nitrates of lead, copper, bis- muth, and cadmium. Add to the solution a few drops of dilute sulphuric acid. Precipitated is lead, as white lead sulphate, which is solu- ble in ammo- nium tartrate with excess of ammonium hydroxide. Solution may contain copper, bismuth, and cadmium Supersaturate with am- monium hydroxide. Precipitated is white bis- muth hy droxide. Dissolve in hydrochloric acid and ap- ply tests for bismuth. Solution may contain copper and cadmium. Divide solution in two parts, and test for copper by potas- sium ferrocyanide in the acidified solution; a red pre- cipitate indicates copper. To second part add potas- sium cyanide and hydro- sulphuric acid. A yellow precipitate indicates cad- mium. Table V.—Treatment of the hydrogen sulphide precipitate which is soluble in ammonium sulphide. The precipitate may contain the sulphides of arsenic, antimony, tin, and a few of those metals which are but rarely met with in qualitative analysis, such as gold, platinum, molybdenum, and others, which latter metals, if suspected, may be detected by special tests. Boil the washed precipitate with strong hydrochloric acid. An insoluble yellow residue consists of arsenous sulphide. The residue is dissolved by boiling with hydrochloric acid and a little potassium chlorate, and the solu- tion examined by Fleitmann’s test. A dark-colored residue may indi- cate gold or platinum for which use special tests. The solution may contain the chlorides of antimony and tin. The solution is introduced into Marsh’s appara- tus when all antimony is gradually evolved as antimoniuretted hydrogen, while tin re- mains with the undissolved zinc as a black metallic powder, which may be collected, washed, dissolved in hydrochloric acid, and the solution tested by the special tests for tin. SEPARATION OF THE METALS OF EACH GROUP. 239 The precipitate may contain the sulphides of iron, manganese, and zinc (cobalt and nickel),1 the hydroxides of chro- mium and aluminium, and possibly the phosphates of barium, calcium, strontium, and magnesium.2 Dissolve the washed precipitate in the smallest possible quantity of warm, dilute hydrochloric acid, and heat the solution with a few drops of nitric acid. To the clear solution add ammonium chloride, and supersaturate it with ammonium hydroxide. The precipitate may contain the hydroxides of iron, aluminium, and chromium, and the phosphates of the alkaline earths or of mag- nesium. Dissolve the precipitate in a little hydrochloric acid, and super- saturate with potassium hydroxide. The solution may contain zinc, manganese, cob;.It, and nickel. Acidulate the ammonia solution with acetic acid, and add hydrogen sulphide. Precipitated is ferric hydroxide, reddish-brown. Dissolve in dilute hydrochloric acid, and add potassium ferrocyanide. A blue precipitate in- dicates iron. Precipitate may also contain the phosphates of the alkaline earths or magnesium. To the solution of the precipitate in hydrochloric acid add ammonium hydroxide until a precipitate is formed which does not redissolve on stirring. Add a few drops of acetic acid to dissolve the precipitate, and then ammonium oxalate. A white precipitate indicates calcium. Barium and strontium are indicated by the addition of calcium sulphate, and distinguished by flame reaction. Test for phosphoric acid by ammonium molybdate. Solution may contain aluminium, and, if green, chromium Supersaturate the alkaline solution slightly with hydrochloric acid, and add ammonium carbonate. A white gelati- nous precipitate indicates aluminium. If the precipitate is green it may contain chro- mium, or aluminium and chromium. Separate as follows: Dissolve the washed precipitate in the smallest possible quantity of nitric acid, evaporate solution in porcelain dish nearly to dryness, add about two volumes of strong nitric acid and frag- ments of potassium chlorate ; heat until solution has assumed a bright orange color. Dilute with water and supersaturate with ammonia; a white precipitate indicates aluminium. To filtrate add barium chloride; a yellow precipitate indicates chromium Precipitate may contain the sulphides of zinc, cobalt, and nickel. If the precipitate be white, it is zinc sulphide only; if black, it may contain cobalt and nickel. If the latter be present, the precipi- tate is dissolved in hydrochloric acid with little nitric acid, and the solution supersaturated with potassium hydroxide. The fil- trate then contains the zinc, the black precipitate cobalt and nickel. Solution may contain manganese, which is verified by adding ammo- nium hydroxide and sul- phide, which produce a flesh-colored precipitate. 1 The sulphides of cobalt and nickel are but sparingly soluble in hydrochloric acid, but dissolve readily in nitro-hydrochlorie acid. 2 In the absence of a sufficient quantity of ammonium chloride some magnesium hydroxide may also be precipitated. Table VI —Treatment of the precipitate formed by ammonium hydroxide and ammonium sulphide. 240 ANALYTICAL CHEMISTRY. The precipitation of sulphur, in the absence of metals of the arsenic group, frequently leads beginners to the assumption that metals of this group are present. The precipitate consisting only of sulphur is white and milky, but flocculent, and more or less colored in the presence of the metals of the arsenic group. Table VII.—Treatment of the precipitate formed by ammonium carbonate. The precipitate may contain the carbonates of barium, calcium, and strontium.1 Dissolve the precipitate in acetic acid, and add potassium chromate. Precipitated is barium, as pale yellow barium chromate. Solution may contain calcium and strontium. Add very dilute sulphuric acid, and let stand for fifteen minutes. Precipitated is stron- tium, as white stron- tium sulphate. Solution may contain calcium. Super- saturate with ammonium hydroxide and add ammonium oxalate. A white precipitate indicates calcium. Table VIII.—Detection of the alkalies and of magnesium. The fluid which has been treated with hydrochloric acid, hydrogen sulphide, am- monium hydroxide, sulphide, and carbonate, may contain magnesium and the alkalies. Divide solution into two portions. To the first portion add sodium phosphate. A white crystalline precipitate indi- cates magnesium.2 The second portion is evaporated to dryness, further heated (or ignited) until all ammonium compounds are expelled, and white fumes are no longer given off. The residue is dissolved in water, and sodium cobaltic nitrite added. A yellow precipitate indicates potassium. The residue is also examined by flame test: a yellow color indicating sodium, a red color lithium. Ammonium compounds have to be tested for in the original fluid by treating it with calcium hydroxide, when ammonia gas is liberated. 1 If an insufficient quantity of ammonium chloride should have been present, some magnesia may also be contained in this precipitate, and may be redissolved by treating it with ammonium chloride solution. 2 If an insufficient quantity of ammonium chloride has been produced in the original solution by the addition of hydrochloric acid and ammonium hydroxide, a portion of the magnesia may have been precipitated by the ammonium hydroxide or carbonate. Questions.—331. By what tests can mercurous chloride be distinguished from the chloride of silver or lead ? 332. How can it be proved that a pre- cipitate produced by hydrogen sulphide in an acid solution contains a metal DETECTION OF ACIDS. 241 35. DETECTION OF ACIDS. General remarks. There are no general methods (similar to those for the separation of metals) by which all acids can be separated, first into different groups, and afterward into the individual acids. It is, moreover, impossible to render all acids soluble (when in combination with certain metals) without decomposition, as, for instance, in the case of carbonic acid when in combination with calcium ; calcium carbonate is insoluble in water, and when the solution is attempted by means of acids, decomposition takes place with liberation of carbon dioxide. Many other acids suffer decomposition in a similar manner, when attempts are made to render soluble the substances in which they occur. It is due to these facts that a complete separation of all acids is not so easily accomplished as the separation of metals. There is, how- ever, for each acid a sufficient number of characteristic tests by which it may be recognized; moreover, the preliminary examination, as well as the solubility of the substance, and the nature of the metal or metals present, will aid in pointing out the acid or acids which are present. If, for instance, a solid substance be completely soluble in water, and if the only metal found were iron, it would be unnecessary to test for carbonic, phosphoric, and hydrosulphnric acids, because the combinations of these acids with iron are insoluble in water; there might, however, be present sulphuric, hydrochloric, nitric, and many other acids, which form soluble salts with iron. Detection of acids by means of the action of strong- sulphuric acid upon the dry substance. The action of sulphuric acid upon a dry powdered substance often furnishes such characteristic indications of or metals of either the arsenic or lead group ? 333. How can mercuric sul- phide be separated from the sulphides of copper and bismuth? 334. Plow does ammonium hydroxide act on a solution containing bismuth and copper? 335. State the action of strong, hot hydrochloric acid on the sulphides of arsenic and antimony. 336. Suppose a solution to contain salts of iron, aluminium, zinc, and manganese ; by what process could these four metals be separated and recognized? 337. How can barium, calcium, and strontium be recognized when dissolved together? 338. By what tests is magnesium recog- nized? 339. State a method of separating potassium when mixed with other metallic compounds. 340. How are ammonium compounds recognized when in solution with other metals? 242 ANALYTICAL CHEMISTRY. the presence or absence of certain acids, that this treatment should never be omitted when a search for acids is made. When the substance under examination is liquid, a portion should be evaporated to dryness, and, if a solid residue remains, it should be treated in the same manner as a solid. Most non-volatile, organic substances (including most organic acids) color sulphuric acid dark when heated with it. Dry inorganic salts when heated with sulphuric acid either are decomposed, with liberation of the acid (which may escape in the gaseous state), or with liberation of volatile products (produced by the decomposition of the acid itself), or no apparent action takes place. See Table IX. Detection of acids by means of reagents added to their neutral or acid solution. Whenever a substance is soluble in water, there is little difficulty of finding the acid by means of Table X.; but if the substance is insoluble in water, and has to be rendered soluble by the action of acids, this table may, in some cases, be of no use, because the acid originally present in the substance may have been liberated, and escaped in a gaseous state (as, for instance, when dissolving in- soluble carbonates in acids), or the tests mentioned in the table may refer to neutral solutions, while it is impossible to render the solution neutral without reprecipitatiug the dissolved acid. If calcium phos- phate, for instance, be dissolved by hydrochloric acid, the maguesium test for phosphoric acid cannot be used, because this test can be applied to a neutral or an alkaline solution only; in attempting, however, to neutralize the hydrochloric acid solution, calcium phosphate itself is reprecipitated. Table XI., showing the solubility or insolubility (in water) of over 300 of the most important inorganic salts, oxides, and hydroxides, will greatly aid the student in studying this important feature. It will also guide him in the analysis of inorganic substances, as it gives directions for over 300 (positive or negative) tests for metals, and an equal number for acids. To understand this, it must be remembered that any salt (or oxide or hydroxide) which is insoluble in water may be produced and pre- cipitated by mixing two solutions, one containing the metal, the other containing the acid of the insoluble salt to be formed. For instance : Table XI. states that the carbonates of most metals are insoluble in water. To produce, therefore, the carbonate of any of these metals (zinc, for instance) it becomes necessary to add to any solution of zinc DETECTION OF ACIDS. 243 (sulphate, chloride, or nitrate of zinc) any soluble carbonate (sodium or potassium carbonate), when the insoluble zinc carbonate is pro- duced. Soluble carbonates consequently are reagents for soluble zinc salts, while at the same time soluble zinc salts are reagents for soluble car- bonates. For similar reasons soluble zinc salts are, according to Table XI., reagents for soluble phosphates, arsenates, arsenites, hydroxides, and sulphides, but not for iodides, chlorides, sulphates, nitrates, or chlorates. The insolubility of a compound in water is not an absolute guide for preparing this compound according to the general rule given above for the precipitation of insoluble compounds, there being some excep- tions. For instance: Cupric hydroxide is insoluble in water; therefore, by adding solution of cupric sulphate to any soluble hydroxide, the insoluble cupric hydroxide should be precipitated, and is precipitated by the soluble hydroxides of potassium and sodium, but not by the soluble hydroxide of ammonium, on account of the formation of the soluble ammonium cupric sulphate. There are not many such exceptions, and to mention them in the table would have greatly interfered with its simplicity, for which reason they have been omitted. For the same reason some compounds, which are not known at all, have not been specially mentioned. For instance, according to Table XI., aluminium carbonate and chromium carbonate are insoluble salts: actually, however, these compounds can scarcely be formed, the affinity between the weak carbonic acid and the feeble bases not being suffi- cient to unite them. Finally, it may be stated that no well-defined line can be drawn between soluble and insoluble substances. There is scarcely any sub- stance which is not slightly soluble in water, and many of the so-called soluble substances are but very sparingly soluble, as, for instance, the hydroxide and sulphate of calcium. Table XII. shows the solubility of a large number of compounds more accurately than Table XI.; it may be used for reference. 244 ANALYTICAL CHEMISTRY. A small quantity of the finely powdered substance is treated with about four times its weight of concentrated sulphuric acid, in a test-tube, care being taken not to heat to the boiling-point of sulphuric acid. No apparent change takes place. No gas is evolved. A colorless gas is evolved. A colored gas is evolved. Sulphuric acid (hepar and barium test). Phosphoric acid (molybdate of ammonium test). Boric acid (green flame after moistening with sulphuric acid). Arsenic acid, ") Silicic acid, Molybdic acid, - Special tests. Phosphorous acid, Arsenous acid, Hydrochloric acid (silver test). Carbonic acid (the gas is generated also by diluted acids in the cold, and renders lime- water turbid). Nitric acid (the vapors turn red on the addition of ferrous sulphate). Sulphurous acid (odor). Hydrosulphuric acid (odor). Hydrofluoric acid (corrodes glass). Acetic acid (odor of acetic ether on the addition of alcohol). Many organic acids are decomposed with liberation of colorless gases. Hydriodic acid (violet vapors of iodine). Hydrobromic acid (brown vapors of bro- mine). Bromic acid (brown vapors; deflagration on charcoal). Chloric acid (the greenish-yellow gas ex- plodes readily). Nitric acid (vapors more red on adding fer- rous sulphate). Nitrous acid (red vapors). Whenever one or more acids are suspected or are indicated by the above tests, their presence is to be verified by the tests in Tables X. and XI., or by the reactions given in connection with the consideration of the acids themselves. For latter tests see Index. Table IX.—Preliminary examination for inorganic acids. Barium chloride pre- cipitates : Calcium chloride pre- Ferric chloride precipi- tates from neutral solution : Magnesium sulphate precipitates in the pres- Silver nitrate precipitates cipitates from neutral or alkaline solution : ence of ammonium hydroxide and chloride. from neutral solution: from neutral or acid solution : Sulphuric acid, white. Sulphurous acid, white. Phosphoric acid, white. Phosphorous acid, white. Carbonic acid, white. Boric acid, white. Sulphuric acid, white. Sulphurous acid, white. Phosphoric acid, white. Hydrochloric acid, white. Hydrobromic acid, Sulphurous acid, white. Phosphoric acid, pale yellow. Phosphorous acid, white, then black. Carbonic acid, white. Boric acid, white. Phosphoric acid, yel- lowish-white. Phosphoric acid, white. white. Hydriodic acid, white. Iodic acid, white. Carbonic acid, white. Boric acid, white. (Ferric hydroxide is precipi- tated and carbon dioxide escapes.) Boric acid, yellowish. Hydrocyanic acid, white. Ferrocyanides, white. Arsenic acid, white. Arsenous acid, white. Chromic acid, pale yel- low. Arsenic acid, white. Arsenic acid, yellowish- white. Ferrocyanidos, blue. Sulphocyanides, red coloration. Hydrosulphuric acid, black. Oxalic acid, yellow. Arsenic acid, white. Arsenic acid, brownish- red. Arsenous acid, yellow. Chromic acid, red. Ferricyanides, reddish- brown. Sulphocyanides, white. Hydrosulphuric acid, black. Oxalic acid, white. Oxalic acid, white. Oxalic acid, white. Not precipitated are: Tartaric acid, white. Citric acid, white. All the above precipitates are soluble in hydrochloric and most other acids, with the exception of barium sul- phate. Tartaric acid, white. Citric acid, white. Tannic acid, black. Acetic acid, a reddish- brown coloration is pro- duced, and, on boiling, a reddish-brown precipitate. Tartaric acid, white; precipitate forms slowly. Tartaric acid, white Citric acid, white. Nitric acid. Nitrous acid. Chloric acid. Hypochlorous acid. Acetic acid. DETECTION OF ACIDS. 245 Table X.—Detection of the more important acids by means of reagents added to the solution. 246 ANALYTICAL CHEMISTRY. Table XI. Systematically arranged table showing the solubility and insolubility of inorganic salts and oxides in water. The dark squares represent insoluble, the white soluble compounds. Carbonate. Phosphate, Arsenate. Arsenite. Oxide. Hydroxide. Sulphide. Iodide. Chloride. Sulphate. Nitrate. Chlorate. Potassium Sodium • Ammonium Alkalies. Calcium Barium Strontium Alkaline Earths. Magnesium Aluminium Ferric Ferrous Zinc. Chromium Nickel, Cobalt Manganese Iron Group. Stannic Stannous Arsenic Arsenous Antimony Gold Platinum Arsenic Group. Copper Bismuth Cadmium Mercuric Mercurous Silver. Lead Lead Group. DETECTION OF ACIDS. 247 w = soluble in water, a = insoluble in water, soluble in acids (HC1, HN03). t = insoluble in water and acids, w a = sparingly soluble in water, but soluble in acids, w t = sparingly soluble in water and acids. at = insoluble in water, sparingly soluble in acids. Aluminium. s p fc O s s ◄3 Antimony. Barium. Bismuth. Cadmium. Calcium. Chromium. Cobalt. H M P* CU C O Ferrous. Ferric. Q ◄ M A & P W e < s Manganese. Mercurous. Mercuric. Nickel. Potassium. Silver. Sodium. Strontium Zinc. Acetate W w W w w w W w w vv w w w w w a w w W w w w VV Arsenate . a w a a a a a a a a a a a a a a a a w a vv a Arsenite . vv a a a a a a a a a a a a a w a w a Borate a vv a a w a a a a a a a a w a a a w a vv a a Bromide . vv w vv a w vv a w w w t w vv vv w w t w w a t w w w a vv vv vv Carbonate . a w a a a a a a a a a a a a a a w a w a a Chlorate . w vv w w w w w vv w w w w w w w w w w w w vv vv Chloride . w vv vv a w vv a vv vv vv vv w vv w w t w w a t vv w w t w w vv Chromate . w a a a a w a a a w w a t w w a w a a w a vv vv a w Citrate w w a a w a w w vv w w a w a a w a w w a w a vv a Cyanide w \v a a w a a t a a t a w a w a t w t w w a Ferricyanide vv w t t w w a w t t w t w a Ferrocyanide vv w a \v t t t t a w a t w t w vv a t Fluoride . w vv w a t w \v a a w w a a w a w a a t a w a vv a w w vv a t vv a Hydroxide a w a w a a a w a a a a a a a a a a w w vv a a Iodide w w w a w a w w w vv w w w vv a w w a a w w t w w vv Nitrate w w w w w w w w w vv w vv w w w w w w w vv w vv Oxalate a w a a a a a vv a a a a a a a w a a a a w a w a a Oxide a or t a vv a a a w a a t a a a a a a a a a a w a w vv a a Phosphate a w vv a vv or a a a vv a a a a a a a a a a a a w a w a a Silicate a t a a a a a a a a a a a a w w a a Sulphate . w w a a w vv t w a w w vv w a t \v w w a w w w w a w t vv Sulphide . a w a vv a a vv a a t a a a a a a a a a a w a w vv a Tartrate . vv w a a a w a a vv w w w a w a w a w a w a a a w a w a a Table XII—Table of solubility. 248 ANALYTICAL CHEMISTRY. 36. DETECTION OF IMPURITIES IN OFFICINAL INORGANIC CHEMICAL PREPARATIONS. General remarks. Very little has been said, heretofore, about impurities which may be present in the various chemical preparations, and this omission has been intentional, because it would have increased the bulk of this book beyond the limits considered necessary for the beginner. Impurities present in chemical preparations are either derived from the materials used in their manufacture, or they have been intention- ally added as adulterations. In regard to the last, no general rule for detecting them can be given, the nature of the adulterating article varying with the nature of the substance adulterated; the general properties of the substance to be examined for purity will, in most cases, suggest the nature of those substances which possibly may have been added, and for them a search has to be made, or, if necessary, a complete analysis, by which is proved the absence of everything else but the constituents of the pure substauce. Impurities derived from the materials used in the manufacture of a substance (generally through an imperfect or incorrect process of manufacture), or from the vessels used in the manufacture, are usually but few in number (in any one substance), and their nature can, in most cases, be anticipated by one familiar with the process of manu- facture itself. For one not acquainted with the mode of preparation, it would be a rather difficult task to study the nature of the impurities which might possibly be present. Questions.—341. Why is sulphuric acid added to a solid substance when it is to be examined for acids ? 342. Mention some acids which cause the libera- tion of colorless, and some which cause the liberation of colored gases when the salts of these acids are heated with sulphuric acid. 343. Mention an acid which is precipitated by barium chloride in acid solution, and some acids which are precipitated by the same reagent in neutral solution. 344. Which acids may be precipitated by silver nitrate from neutral solutions, and which from either neutral or acid solutions? 345. Mention some acids which form soluble salts only. 346. Mention three soluble, and three insoluble carbonates, phos- phates, arsenates, sulphates, and sulphides respectively. 347. Which oxides or hydroxides are soluble, and which are insoluble in water? 348. Mention some metals, the solutions of which are precipitated by soluble chlorides, iodides, and sulphides. 349. State a general rule according to which most insoluble salts may be formed from two other compounds. 350. Why is it sometimes impossible to render a substance soluble in order to test for the acid in the solution obtained? DETECTION OF IMPURITIES. 249 The same remarks apply to the methods by which the impurities can be detected. One familiar with analytical chemistry can easily find, in most cases, a good method by which the presence or absence of an impurity can be demonstrated; but to one unacquainted with chemistry, it might be an impossibility to detect impurities, even if a method were given. For these reasons little stress has been laid upon the occurrence of impurities in the various chemical preparations heretofore considered; moreover, the U. S. P. itself gives, in most cases, directions for the detection of impurities, and, finally, the analyst may avail himself of books specially treating of the examination of chemicals. In order to give the student, and especially the beginner, a guide to examination for impurities, the following pages furnish a few direc- tions in regard to the impurities which may be present in the more important chemical preparations, and also some of the tests used for recognizing them. Examination of sulphuric, sulphurous, nitric, phosphoric, and hydrochloric acids. The pure acids should be colorless, and upon evaporation leave no residue whatever (phosphoric acid requires a red heat for complete evaporation). After being neutralized with an alkali, and then slightly acidulated with hydrochloric acid, they should give no precipitate with hydrogen sulphide, ammonium hydroxide, or sulphide, or carbonate. Sulphuric, phosphoric, and hydrochloric acids sometimes (nitric acid very rarely) contain arsenic, and should, therefore, be examined by Bettendorlf’s or Gutzeit’s test. Sulphuric acid may contain sulphurous or nitric acid; the former may be recognized by potassium permanganate, which is decolorized by sulphurous acid; the latter is detected by carefully pouring solu- tion of ferrous sulphate upon the acid, when, if nitric acid be present, a brownish zone appears at the line of contact. Sulphurous acid frequently contains sulphuric acid ; it should not give more than a very slight turbidity with barium chloride (limit of sulphuric acid). Nitric acid diluted with 5 parts of water should afford no precipi- tate with either barium chloride or silver nitrate (absence of sulphuric and hydrochloric acids). Phosphoric acid should be tested for phosphorous, nitric, sulphuric, hydrochloric, pyrophosphoric, and metaphosphoric acids. Phosphorous acid is indicated by the formation of a dark color when silver nitrate 250 ANALYTICAL CHEMISTRY. is added to the diluted acid ; sulphuric and hydrochloric acids by the formation* of white precipitates on the addition of barium chloride or silver nitrate, respectively, to the diluted acid; pyrophosphoric and metaphosphoric acids, by the formation of a precipitate upon adding to the acid an equal volume of tincture of chloride of iron. Hydrochloric acid, after being diluted, should not give a precipitate with barium chloride (absence of sulphuric acid), and should not liberate iodine from potassium iodide (absence of free chlorine). The officinal acids should have the strength required by the U. S. P. The amount of actual acid can be determined either by the specific gravity of the liquid acid, or by the quantity of an alkali required to neutralize a certain quantity of the acid. (See next chapter, on volu- metric analysis.) Examination of potassium, sodium, and ammonium compounds. The acidulated solution of the hydroxides, carbonates, bicarbonates, sulphates, nitrates, and phosphates of potassium, sodium, and ammo- nium should afford no precipitate with hydrogen sulphide; or with hydroxide, sulphide, carbonate, or phosphate of ammonium (absence of heavy metals, alkaline earths, and magnesium). The nitrates, hydroxides, carbonates, and bicarbonates (the latter three, after being supersaturated with nitric acid), should be tested with solution of barium chloride and silver nitrate for sulphates and chlorides. Bicarbonates may contain carbonates ; the latter are shown by dis- solving the salt in cold water and adding a few drops of solution of phenol-phthalein. In the presence of carbonates the solution turns red. Iodides may contain an iodate, which is indicated by a blue color on the addition of gelatinized starch and some diluted sulphuric acid; chlorides and bromides may be detected bv precipitating a solution of the iodide with an excess of silver nitrate, collecting and washing the precipitated iodide (bromide and chloride) of silver, and digesting it with ammonium hydroxide, which dissolves the chloride and bromide, but not (or only traces of) the iodide; the filtered ammonium solution, upon being supersaturated with nitric acid, will give a white precipi- tate if chlorides or bromides are present; a slight turbidity may be due to traces of dissolved iodide. Potassium salts should impart a violet color to a non-luminous flame, a yellow color indicating sodium. All compounds of ammomium should be completely volatilized by heat, leaving no residue. DETECTION OF IMPURITIES. 251 Examination of compounds of calcium. Calcium chloride and cal- cium carbonate, the latter after having been converted into chloride by neutralizing with hydrochloric acid, should not be precipitated by hydrogen sulphide or by ammonium hydroxide (absence of heavy metals and of aluminium); the solution, after adding ammonium hydroxide, ammonium chloride, and an excess of ammonium carbonate, and fil- tering off the precipitated calcium carbonate, should not be precipi- tated by sodium phosphate (absence of magnesium). Distilled water digested with calcium carbonate or with tricalcium phosphate, should leave no residue upon evaporation. Calcium carbonate and phosphate should be tested for sulphates and chlorides by dissolving the salts in dilute nitric acid and adding barium chloride aud silver nitrate. Iron will be indicated in calcium phosphate by saturating its solu- tion in hydrochloric acid with hydrosulphuric acid, and then adding an excess of ammonium hydroxide; the precipitate should be white, a dark color indicating iron (or possibly other heavy metals). Calcium hypochlorite (bleach ing-powder) has to be examined by quantitative (volumetric) methods to ascertain the amount of hypo- chlorous acid present. Examination of magnesium compounds. Magnesium sulphate, oxide, and carbonate, the latter two after being dissolved in dilate hydro- chloric acid, should not be precipitated by hydrogen sulphide, nor by ammonium hydroxide, sulphide, or carbonate after a sufficient amount of ammonium chloride has been added (absence of heavy metals and alkaline earths). Magnesium sulphate should be tested for chlorides by the addition of silver nitrate. Water digested with oxide or carbonate should leave no residue upon evaporation. The same compounds dissolved in nitric acid should not be precipitated by either barium chloride or silver nitrate. The oxide should not effervesce with acids (absence of carbonic acid). Examination of aluminium compounds. Aluminium sulphate and aluminium-potassium sulphate (aium) when dissolved in water, or aluminium hydroxide when dissolved in sulphuric acid, should not be precipitated by hydrogen sulphide and should show no blue color on the addition of potassium ferrocyanide (absence of iron). Potassium or sodium hydroxide, added to the above solutions, should cause a white gelatinous precipitate, which is completely dissolved by an ex- 252 ANALYTICAL CHEMISTRY. cess of the reagent. Water digested with aluminium hydroxide should leave no residue upon evaporation. Examination of compounds of iron. The solution of ferrous salts, acidulated with hydrochloric acid, should not be precipitated by hydrogen sulphide ; solutions of ferric salts give with the same re- agent an almost white precipitate of sulphur ; a more or less colored precipitate would indicate the presence of metals of the arsenic or lead group. Solutions of ferric salts (or of ferrous salts after they have been converted into ferric salts by heating with hydrochloric and nitric acids), after having been precipitated by an excess of ammonium hydroxide and filtered, should not impart a blue color to the filtrate (absence of copper), and this filtrate should not be precipitated by ammonium carbonate or phosphate (absence of alkaline earths and magnesium) and should not leave a residue on evaporation and gentle ignition (absence of zinc, manganese, alkalies, etc.). When solution of potassium hydroxide is added in excess to ferric solutions, the filtrate should not be precipitated by ammonium car- bonate after the alkaline solution has been neutralized by hydro- chloric acid (absence of aluminium). Compounds of iron may contain arsenic, and should therefore be examined by Gutzeit’s test. Ferrous salts (sulphate) should give no immediate blue precipitate with potassium ferrocyanide (absence of ferric compounds). Ferric salts should give no blue color with potassium ferricyanide (absence of ferrous compounds). Ferric chloride should also be tested for nitric acid by placing a crystal of ferrous sulphate into the solution to which sulphuric acid has been added (a brown color indicating nitric acid), and for sul- phuric acid by barium chloride. Examination of compounds of zinc. Chloride or sulphate of zinc, when dissolved in water, oxide or carbonate of zinc when dissolved in dilute hydrochloric acid, should give no precipitate, in the acidulated solution, with hydrogen sulphide (absence of the metals of the arsenic and lead groups). Ammonium hydroxide added to zinc solutions should cause a white precipitate, which is completely dissolved by an excess of the reagents (absence of iron, aluminium, etc.). The filtrate of zinc solutions from which the metal has been pre- DETECTION OF IMPURITIES. 253 cipitated by an excess of ammonium sulphide, should not be precipi- tated by ammonium carbonate or phosphate (absence of alkaline earths and magnesium), nor should it leave a residue on evaporation and gentle ignition (absence of alkalies). Water digested with zinc oxide or carbonate should leave no residue on evaporation. Zinc sulphate should be tested for chlorides by silver nitrate, and zinc chloride for sulphates by barium chloride. Examination of compounds of manganese. The native manganese dioxide invariably contains other mineral matter; according to the U. S. P., it should contain not less than 66 per cent, of the dioxide, and should, consequently, be examined by quantitative analysis. A solution of manganese sulphate, acidulated with hydrochloric acid, should give no precipitate with hydrogen sulphide (absence of copper, etc.); the aqueous solution of the salt from which the manga- nese has been precipitated by ammonium sulphide should leave no residue on evaporation and gentle ignition (absence of magnesium, alkalies, etc.). Potassium permanganate is completely decolorized (deoxidized) by oxalic acid or sulphurous acid. The colorless solution should be tested with ferrous sulphate for nitric acid, and with silver nitrate for hydrochloric acid. It should also be tested quantitatively. Examination of compounds of chromium. Solution of potassium dichromate and chromium trioxide, acidulated with nitric acid, should not be precipitated by barium chloride (absence of sulphuric acid) ; both substances, when neutralized with potassium hydroxide, should give an immediate red precipitate with silver nitrate, a white precipi- tate indicating chlorides. Examination compounds of lead. Solutions of compounds of lead, when completely precipitated with either sulphuric acid or with hydrogen sulphide, should yield a filtrate which leaves no residue on evaporation and gentle ignition (absence of most other metals). Lead oxide (litharge) and lead carbonate should dissolve completely in nitric acid. Lead iodide should be tested for lead chromate by triturating it with 2 parts of ammonium chloride and water, which dissolve the iodide, but not the chromate. 254 ANALYTICAL CHEMISTRY. Examination of compounds of copper. Solutions of cupric salts (sulphate), when completely precipitated by hydrogen sulphide, should yield a filtrate which gives no precipitate with ammonium hydroxide, sulphide, or carbonate. Excess of ammonium hydroxide added to a cupric solution should form a dark-blue solution without leaving a residue. Examination of compounds of bismuth. Subnitrate and subcar- bonate of bismuth should be completely soluble in about 8 parts of a mixture of equal parts of nitric acid and water. When this solution is poured into 50 parts of water a white precipitate falls; the filtrate from this precipitate may be used to test for lead by sulphuric acid, for silver by hydrochloric acid, for sulphates by barium chloride, for chlorides by silver nitrate (all of which reagents give white precipi- tates if the impurities named are present), and for copper by an excess of ammonium hydroxide, which precipitates the bismuth yet in solution, while copper would be indicated by the blue color of the filtrate. Another portion of the bismuth salt is dissolved in hydrochloric acid and all bismuth precipitated by hydrogen sulphide; the filtrate should leave no residue (absence of metals of the iron group and of the light metals). Examination of compounds of silver. A solution of silver nitrate in water, or of silver oxide iu nitric acid, should give a white precipi- tate with hydrochloric acid, which precipitate should be completely dissolved by ammonium hydroxide. The filtrate from a solution from which all silver has been precipitated by hydrochloric acid, should leave no residue on evaporation. Silver oxide is readily dissolved by ammonia water. Examination of compounds of mercury. All compounds of mercury are completely volatilized on heating. The oxides of mercury are dissolved by heating with about 10 parts of dilute nitric acid; mercuric oxide, when heated in a tube, should evolve no red fumes (absence of nitric acid). Mercuric chloride should be soluble in water and in alcohol; it should be tested for arsenic. Mercurous chloride, when digested with water, should yield a filtrate which, on evaporation, leaves no residue, and which is not changed by either hydrogen sulphide, silver nitrate, or potassium iodide (absence DETECTION OF IMPURITIES. 255 of mercuric chloride). Mercurous chloride is soluble in warm hydro- chloric acid to which a little nitric acid is gradually added. Mercuric iodide should be dissolved by 25 parts of boiling alcohol, or by digesting it with potassium iodide and water. Mercurous iodide should be tested for mercuric iodide by digesting it with alcohol, which, upon evaporation, should leave no residue. Ammoniated mercury should be wholly soluble, without efferves- cence, in warm hydrochloric, nitric, or acetic acid; when digested with diluted alcohol, the filtrate should not be acted upon by hydrogen sulphide or potassium iodide (absence of mercuric chloride). The basic mercuric sulphate is almost insoluble in cold water, but soluble in diluted hydrochloric or nitric acid. Examination of compounds of arsenic. Arsenous and arsenic oxides, arsenous bromide and iodide are completely volatilized by heating (a residue indicating non-volatile impurities). The four compounds are also soluble in water, the bromide and iodide with decomposition. Examination of compounds of antimony. Antimonious oxide should be completely soluble in tartaric acid; the solution should be tested for chlorides and sulphates by means of silver nitrate and barium chloride; also for iron, by potassium ferrocyanide. Antimonious and antimonic suljphides are soluble in concentrated hydrochloric acid, with liberation* of hydrogen sulphide and formation of the trichloride. Antimonic oxide, as well as antimonious oxychloride, is dissolved by boiling with sodium or potassium hydroxide. Questions.—351. Give some general methods by which the mineral acids may be examined for metallic impurities and their strength determined. 352. How is sulphuric acid to be tested for sulphurous and nitric acids, and how is nitric acid tested for sulphuric and hydrochloric acids? 353. What impurities are sometimes present in phosphoric acid ? How is their presence demon- strated ? 354. By what tests is sodium carbonate detected in sodium bicarbo- nate, and potassium iodate in potassium iodide ? 355. How can the presence of ferrous sulphate he demonstrated in either ammonium chloride, calcium phosphate, magnesium carbonate, aluminium sulphate, or cupric sulphate? 356. State the various methods by which preparations of iron are examined for metallic impurities, and how ferric chloride can also be tested for ferrous chloride, nitric and sulphuric acids. 357. By what methods can zinc oxide be tested for any other metal, and for carbonic, sulphuric, or hydrochloric acid? 358. Give methods for detecting metallic impurities in compounds of lead, copper, bismuth, and silver. 359. What is the action of heat upon mercury 256 ANALYTICAL CHEMISTRY. 37. METHODS FOR QUANTITATIVE DETERMINATIONS. General remarks. Quantitative determination of the different elements or groups of elements may be accomplished by various methods, which differ generally with the nature of the substance to be examined. But even one and the same substance may often be analyzed quantitatively by entirely different methods, of which the two principal ones are the gravimetric and volumetric methods. In the gravimetric method, the quantities of the constituents of a substance are determined by separating and weighing them either as such, or in the form of some compound, the exact composition of which is known. For instance: From cupric sulphate, the copper may be precipitated as such by electrolysis and weighed as metallic copper, or it may be precipitated by sodium hydroxide as cupric oxide, CuO, and weighed as such. Knowing that every 79.2 parts by weight of cupric oxide contain of oxygen 16 parts and of copper 63.2 parts, the weight of copper contained in the cupric oxide found may be readily calculated. In the volumetric method, the determination is accomplished by add- ing to a weighed quantity of the substance to be examined, a solution of a reagent of a known strength until the reaction is just completed, no excess being allowed. For instance: We know that every 80 parts by weight of sodium hydroxide precipitate 79.2 parts by weight of cupric oxide, containing 63.2 parts by weight of copper. There- fore, if we add a solution of sodium hydroxide of known strength to a weighed portion of cupric sulphate until all the copper is precipitated, we may calculate from the volume of soda solution used the weight of sodium hydroxide, and from this the weight of copper which has been precipitated. The operation of volumetric analysis is termed titration. Gravimetric methods. While the quantitative determinations by these methods differ widely in some cases, there are a number of oper- ations so often and so generally employed that a few remarks may be of advantage to the beginner. A small quantity (generally from 0.5 to 1 gramme) of the substance to be analyzed is very exactly weighed on a delicate balance, transferred to a beaker, and dissolved in a suit- able ageut (water or acid). From this solution the constituent to be determined is precipitated completely, which is ascertained by allowing compounds, and by what solvents may the various officinal mercury prepara- tions be dissolved? 360. How is mercuric chloride detected in mercurous chloride, and how mercuric iodide in mercurous iodide ? METHODS FOR QUANTITATIVE DETERMINATIONS. 257 the precipitate to subside and adding to the clear liquid a few drops more of the agent used for precipitation. The precipitate is next collected upon a small filter of good filter paper containing as little of inorganic con- stituents (ash) as possible; the particles of precipitate which may adhere to the beaker are carefully washed off by means of a earners-hair brush. The precipitate is well washed (generally with pure water) until free from adhering solution, and dried by placing funnel and contents in a drying-oven, Fig. 28, in which a constant temperature of about 100° C. (212° F.) is maintained. The dried filter is then taken from the funnel and its contents are transferred to a platinum (or porcelain) Drying-oven. crucible, which has been previously weighed and stands on a piece of glazed, colored paper in order to collect any particle of the dried pre- cipitate which may happen to fall beside the crucible. The filter, from which the precipitate has been removed as completely as possible, by slightly rubbing it, is now folded, placed upon the lid of the crucible, which rests ou a triangle over a gas burner, and completely incinerated. The remaining filter-ash, with particles of the precipitate mixed with it, is transferred to the crucible, which is now placed over the burner and heated until all water (or possibly other substances) is completely expelled. After cooling, the crucible is weighed, the weight of the empty crucible and that of the filter-ash (the latter having been pre- 258 ANALYTICAL CHEMISTRY. viously determined by burning a few filters of the same kind) deducted, and thus the quantity of the precipitate determined. As platinum crucibles and many precipitates, after ignition, absorb moisture from the air, it is Avell to allow the heated crucible to cool in a desiccator. This is a closed vessel in which the contained air is kept dry by means of concentrated sulphuric acid. Fig. 29 shows a convenient form of desiccator. The empty crucibles should be Aveighed under the same conditions —i. e.) after having been heated and cooled in a desiccator. Some precipitates (as, for instance, potassium platinic chloride), cannot be ignited without suffering partial or complete decomposition. Fig. 29. Fig. 30 Desiccator. Watch-glasses for weighing filters. It is for this reason that some precipitates are collected upon filters which have been previously dried at 100° C. (212° F.) and weighed ca*refully. The precipitate is then collected upon the weighed filter, well washed, dried at 100° C. (212° F.) and Aveighed. The weighing of dried filters is best accomplished by placing them between two Avatch-glasses held together by means of a brass or nickel clamp, as shown in Fig. 30. The above-described methods may be employed for the determina- tion of those substances which can be precipitated from their solu- tions in the form of some stable compound. Aluminium, zinc, iron, bismuth, copper, etc., may, for instance, be precipitated as hydroxides and Aveighed as oxides, into which the precipitated compound is con- verted by ignition. Sulphuric acid may be precipitated and Aveighed as barium sulphate, phosphoric acid may be precipitated by magnesia mixture aud Aveighed as magnesium pyrophosphate, etc. Some sub- METHODS FOR QUANTITATIVE DETERMINATIONS. 259 stances, like nitric acid, chloric acid, etc., cannot be precipitated from their solutions, for which reason other methods have to be employed for their determination. Volumetric methods. The great advantage of volumetric over gravimetric analysis consists chiefly in the rapidity with which these Fig. 32. Fig. 31. Litre flask. Pipettes. determinations are performed. Unfortunately, volumetric methods cannot be employed to advantage for the estimation of all substances. The special apparatus required for volumetric analysis consists of a few flasks, some pipettes, burettes, aud a burette-holder. The flasks 260 ANALYTICAL CHEMISTRY. should have a mark on the neck, indicating a capacity of 100,'250, 500, and 1000 c.c. respectively. (See Fig. 31.) Of pipettes (Fig. 32) are mostly used those having a capacity of 5, 10, 25, and 50 cubic centimetres. Mohr’s burette and clamp. Mohr's burette and holder. Of burettes many different forms are used; in most cases Mohr’s burette (Figs. 33 aud 34) answers all requirements, but its application is excluded whenever the test solution is chemically affected by rubber, as in the case of solutions of silver, permanganate, and a few odier substances. For such solutions Mohr’s burette with glass stop-cock, or Gay Lussac’s burette (Fig. 35), is generally used. METHODS FOR QUANTITATIVE DETERMINATIONS. 261 Standard solutions. The test solutions used in volumetric analysis are adjusted according to a uniform system, so that each solution contains in a litre (1000 c.c.) the weight of one atom or one molecule of the active reagent expressed in grammes. This rule refers to all cases of univalent elements (Ag, Cl, I), or monobasic acids (HC1, HN03), or monacid bases (KOH, NH4OH). In case a bivalent element (O, S), or dibasic acids (H2S04, H2C204), or di-acid bases [Ca(OH)2], are used in volumetric solutions, only one-half of the atomic or molecular weight in grammes is used per litre, in order to have the saturating or neutralizing power the same for an equal number of cubic centimetres of univalent and bivalent substances. To illustrate why this is done, if we were to take the molecular weight of hydro- chloric acid, 36.4, and of sulphuric acid, 98, in grammes, diluted to 1000 c.c., the saturating power of 1 c.c. of the diluted sulphuric acid would be equal to that of 2 c.c. of hydrochloric acid solution, because 36.4 parts by weight of hydrochloric acid saturate 40 parts by weight of sodium hydroxide, and 98 parts by weight of sul- phuric acid saturate 80 parts by weight of sodium hydroxide. The solutions thus obtained are known as normal solutions. For some operations these normal solutions are too concen- trated, and are diluted to one-tenth of their strength, and are then called deci-normal solutions. Normal solutions are generally designated by deci-normal solu- tions by centi-normal solutions by ; solutions containing twice the amount are designated as double normal, - ; half the amount semi-normal, In some instances volumetric solutions are prepared which do not Fig. 35. Gay Lussae’s burette. 262 ANALYTICAL CHEMISTRY. belong to the above system of normal solutions, but are adjusted to correspond to a certain unit of the special substance they are to act upon. Such solutions are called empirical solutions; as an instance, may be mentioned Fehling’s solution, used for the determination of sugar. This solution is so adjusted that 1 c.c. decomposes or indicates 0.005 gramme of grape-sugar. Different methods of volumetric determination. Of these we have at least three, which may be called the direct, the indirect, and the method of rest or residue. The direct methods are used in all cases in which the quantities of volumetric solutions can be added until the reaction is complete : for instance, until an alkaline substance has been neutralized by an acid, or a ferrous salt has been converted into a ferric salt by potassium permanganate, etc. In the indirect methods one substance, which cannot well be deter- mined volumetrically, is made to act upon a second substance, with the result that, by this action, an equivalent amount of a substance is generated or liberated, which may be titrated. For instance : Per- oxides, chromic and chloric acids when boiled with strong hydro- chloric acid, liberate chlorine, which is not determined directly, but is caused to act upon potassium iodide, from which it liberates the iodine, which may be titrated with sodium thiosulphate. The methods of residue are based upon the fact that while it is im- possible or extremely difficult to obtain complete decomposition between certain substances and reagents, when equivalent quantities are added to one another, such a complete decomposition is accom- plished by adding an excess of the reagent, which excess is afterward determined bv a second volumetric solution. For instance: Car- bonate of calcium, magnesium, zinc, etc., cannot well be determined directly, for which reason an excess of normal acid is used for their decomposition, this excess being titrated afterward by means of an alkali. Indicators. In all cases of volumetric determination it is of the greatest importance to observe accurately the completion of the reac- tion. In some cases the final point is indicated by a change in color, as, for instance, in the case of potassium permanganate, which changes from a red to a colorless solution, or chromic acid, which changes from orange to green under the influence of deoxidizing agents. In other cases the determination is indicated by the formation or cessa- METHODS FOR QUANTITATIVE DETERMINATIONS. 263 tion of a precipitate, and in yet others the final point could not be noticed with precision unless rendered visible by a third substance added for that purpose. Such substances are termed indicators. Litmus, phenol-phthale'in, methyl-orange, etc., are used as indicators in acidimetry and alka- limetry. Starch paste is an indicator for iodine, potassium chromate for silver, etc. Litmus solution. This is made by exhausting coarsely powdered litmus with boiling alcohol, which removes a red coloring matter, erythrolitmin. The residue is treated with about an equal weight of cold water so as to dissolve the excess of alkali present in litmus. The remaining mass is extracted with about five times its weight of boiling water, and filtered. The solution should be kept in wide-mouthed bottles, stoppered with loose plugs of cotton to ex- clude dust but to admit air. Blue and red litmus paper is made by impregnat- ing strips of unsized white paper with the blue solution obtained by the above process, or with this solution after just enough hydrochloric acid has been added to impart to it a distinct red tint. Phenol-phthalein solution. 1 gramme of phenol-phthaleln is dissolved in 100 c.c. of diluted alcohol. The colorless solution is colored deep purplish-red by alkali hydrates or carbonates, but not by bicarbonates; acids render the red solution colorless. The solution is not suitable as an indicator for ammonia or bicarbonates. Methyl-orange solution. 1 gramme of methyl-orange (also known as helian- thin, tropseolin D, or Poirier’s orange III), dimethylamido-azobenzol-sul- phonic acid, (CH3)2C6H4.N.NC6H4.S03H, is dissolved in 1000 c.c. of water. The solution is yellow when in contact with alkaline hydrates, carbonates, or bicarbonates. Carbonic acid does not affect it, but mineral acids change its color to crimson. Rosalie acid solution. 1 gramme of commercial rosolic acid (chiefly C20H16O3) is dissolved in 10 c.c. of alcohol, and water added to make 100 c.c. The solu- tion turns violet-red with alkalies, yellow with acids. Titration. This term is used for the process of adding the volu- metric solution from the burette to the solution of the weighed sub- stance until the reaction is completed. We also speak of the standard or litre of a volumetric test-solution, when we refer to its strength per volume (per litre or per cubic centimetre). Of the principal processes of titration, or of volumetric methods used, may be mentioned those based upon neutralization (acidimetry and alkalimetry), oxidation and reduction (permanganates and chro- mates as oxidizing, oxalic acid and ferrous salts as reducing agents) and precipitation (silver nitrate by sodium chloride). Acidimetry and alkalimetry. Preparing the volumetric test-solu- 264 ANALYTICAL CHEMISTRY. tions is often more difficult than to make a volumetric determination. Whenever the reagents employed can be obtained in a chemically pure condition it is an easy task to prepare the solution, because a definite weight of the reagent is dissolved in a definite volume of water. In many instances, however, the reagent cannot be obtained absolutely pure, and in such cases a solution is made and its standard adjusted afterward by methods which will be spoken of later. Neither the common mineral acids, such as sulphuric, hydrochloric, and nitric acids, nor the alkaline substances, such as sodium hydrox- ide or ammonium hydroxide, are sufficiently pure to permit of being used directly for volumetric solutions, because these substances con- tain water, and an absolutely correct determination of the amount of this water is an operation which involves a knowledge of gravimetric methods. It is for this reason that the basis in preparing a volumetric normal acid solution is oxalic acid, a substance which can be readily obtained in a pure crystallized condition. Normal acid solution. Crystallized oxalic acid has the composition H2C204.2H2O and a molecular weight of 125.7. Being dibasic, only half of its weight is taken for the normal solution, which is made by placing 62.85 grammes of pure crystallized oxalic acid in a litre flask, dissolving it in pure water, filling up to the mark at the tem- perature of 15° C. (59° F.) and mixing thoroughly. Normal solutions of sulphuric or hydrochloric acid are, for various reasons, often preferred to oxalic acid. These solutions are best made by diluting approximately the acids named, titrating the solution with normal sodium hydroxide, using phenol-phthalein as an indicator, and adding water until equal volumes saturate one another. For instance, if it should be found that 10 c.c. normal alkali solution neutralize 7.6 c.c. of the acid, then 24 c.c. of water have to be added to every 76 c.c. of the acid in order to obtain a normal solution. Normal sul- phuric acid contains 48.91 grammes of H2S04, and normal hydro- chloric acid 36.37 grammes of HC1 per litre. These normal solutions can be made conveniently by diluting either 29 c.c. of pure, concentrated sulphuric acid of sp. gr. 1.84, or 130 c.c. of hydrochloric acid of sp. gr. 1.16 to 1000 c.c. The solutions thus obtained are yet too con- centrated and are adjusted as described above. Other methods of determining the exact standard of normal acids depend upon the precipitation of 10 c.c. of the sulphuric acid solution by barium chloride, or of 10 c.c. of the hydrochloric acid solution by silver nitrate, and METHODS FOR QUANTITATIVE DETERMINATIONS. 265 weighing the precipitated barium sulphate or silver chloride. Ten c.c. of normal sulphuric acid give 1.1736 gramme of barium sulphate, and 10 c.c. of normal hydrochloric acid 1.43 gramme of silver chloride. A third method depends on the formation of, and the weighing as, an ammonium salt. Ten c.c. of either acid are neutralized (or slightly super- saturated) with ammonia water. The solution is evaporated in a previously weighed platinum dish over a water-bath, the dry salt is repeatedly moistened with alcohol,* and finally dried in an air-bath at a temperature of 105° C. (221° F.) for about half an hour. Ten c.c. of normal sulphuric acid give of ammonium sulphate 0.6592 gramme, and 10 c.c. of normal hydrochloric acid of ammonium chloride 0.5338 gramme. Normal alkali solution. A normal solution of sodium carbonate may be made by dissolving 52.92 grammes (one-half the molecular Fig. 36. Titration. weight) of pure sodium carbonate (obtainable by heating pure sodium bicarbonate to a low red-heat) in water, and diluting to one litre. This solution, however, is not often used, but may serve for standard- izing acid solutions, as it has the advantage of being prepared from a substance that can be easily obtained in a pure condition, which is not the case in preparing the otherwise more useful normal solutions of 'potassium or sodium hydroxide, both of which substances contain and absorb water. The solutions are made by dissolving about 60 grammes of potas- sium hydroxide or 50 grammes of sodium hydroxide in about 1000 266 ANALYTICAL CHEMISTRY. c.c. of water, titrating this solution with normal acid, and diluting it with water, until equal volumes of both solutions neutralize one another exactly. The indicators used in alkalimetry are chiefly solution of litmus or phenol-phthale'in, only a few drops of either solution being needed for a determination. Whenever carbonates are titrated with acids, or vice versa, the solu- tion has to be boiled toward the end of the reaction in order to drive otf the carbon dioxide, as neither of the two indicators mentioned gives reliable results in the presence of carbonic acid or an acid car- bonate. The proper mode of performing the operation of titration is shown in Fig. 36. Neutralization equivalents. The normal solutions of acid and alkali may be used for the determination of a large number of sub- stances, either directly (as in the case of free acids, caustic and alka- line carbonates and bicarbonates) or indirectly (as in the case of salts of most of the organic acids, with alkalies, which are first converted into carbonates by ignition). One c.c. of normal acid is the equivalent of: Ammonia, NH3 Gramme. . 0.01701 Ammonium carbonate, NH4HC03.NH4lSrH2C02 . . 0.05226 Lead acetate, crystallized, Pb(C2H302)2.3H20 . 0.18900 Lead subacetate, Pb20(C2H302)2 .... . 0.13662 Lithium carbonate, Li2COg ..... . 0.03693 Lithium citrate, Li3C6H507 1 .... . 0.06986 Potassium acetate, KC2H302 1 . 0.09789 Potassium bicarbonate, KHC03 .... . 0.09988 Potassium bitartrate, KHC4H406 1 ... . 0.18767 Potassium carbonate, K2C03 . 0.06895 Potassium citrate, crystallized, K3C6H507.H20 1 . . 0.10789 Potassium hydroxide, KOH ..... . 0.05599 Potassium permanganate, KMn04 2 . 0.03153 Potassium sodium tartrate, KNaC4H406.4H20 1 . 0.14075 Potassium tartrate, 2K2C4H406.H20 1 . 0.11734 Sodium acetate, NaC2H302.2H20 1 ... . 0.11778 Sodium bicarbonate, NaHC03 .... . 0.08385 Sodium borate, crystallized, Na2B4O7.10H2O . . 0.19096 Sodium carbonate, crystallized, Na2CO3.10H2O . 0.14272 Sodium carbonate, Na2C03 . 0.05295 Sodium hydroxide, NaOH . 0 03996 1 After ignition. 2 With oxalic acid only. METHODS FOR QUANTITATIVE DETERMINATIONS. 267 One c.c. of normal sodium carbonate, or sodium hydroxide, is the equivalent of: Acetic acid, HC2H302 .... Gramme. . 0.05980 Citric acid, crystallized, H3C6H507.H20 . . 0.06983 Hydrobromic acid, HBr . 0.08076 Hydrochloric acid, HC1 .... . 0.03637 Hydriodic acid, HI .... . 0.12753 Hypophosphorous acid, HPH202 . . 0.06588 Lactic acid, HC3H503 .... . 0.08989 Nitric acid, HNOs . 0.06289 Oxalic acid, crystallized, H2C204.2H30 . . 0.06285 Potassium dichromate, K2Cr207 . 0.14689 Sulphuric acid, H2S04 .... . 0.04891 Tartaric acid, H2C4H406 . 0.07482 Oxidimetry. Potassium permanganate. The substances generally used as oxidizing agents are potassium permanganate and potassium dichromate, both of which salts can be obtained in a pure crystallized condition. Potassium permanganate, KMn04 = 157.67, acts generally in the presence of free acids, upou deoxidizing substances, by losing 5 atoms of oxygen of the 8 atoms contained in two molecules, as is shown in the following equations : 2KMn04 + 5H2C204 + 3H2S04 = K2S04 + 2MnS04 + 10CO2 + 8H20. 2KMn04 + 10FeSO4 + 8H2S04 = K2S04 + 2MnS04 + 5Fe23S04 + 8H20. It follows that two-fifths of the molecular weight of potassium per- manganate, or 63.068 grammes, are the equivalent of 1 oxygen atom. But as oxygen is diatomic and the volumetric normal is calculated for monatomic values, this number must be divided by 2, and 31.534 grammes of pure crystallized potassium permanganate is therefore the amount to furnish 1 litre of normal solution, but as this is too con- centrated for most determinations, a deci-normal solution containing 8.1534 grammes to the litre is generally employed. Permanganate solution, when made with pure water, does not alter its standard, if contained in a tight (glass-stoppered) bottle, kept in a dark closet. If the water used contain organic matter, the standard will alter until all that matter is oxidized. The standardizing of permanganate solution may be done by dissolv- ing 0.2 gramme of pure, thin iron wire in about 20 c.c. of dilute sul- phuric acid (1 acid, 5 water) by the aid of heat, and in a flask arranged as in Fig. 37. The flask is provided, by means of a perforated cork, with a piece of glass tubing, to which is attached a piece of rubber 268 ANALYTICAL CHEMISTRY. tubing in which is cut a vertical slit about one inch long and which is closed at the upper end by a piece of glass rod ; gas or steam gener- ated in the flask may escape, while atmospheric air cannot enter, the ferrous solution being thus protected from oxidation. The iron solution, obtained from the 0.2 gramme of iron, is cooled and diluted with about 300 c.c. of water, and then deci-normal potas- sium permanganate solution is added with constant stirring until the solution is tinged pinkish. Fig. 37. Flask for dissolving iron for volumetric determination. As 1 c.c. of deci-normal permanganate solution corresponds to 0.005588 gramme of metallic iron, the 0.2 gramme iron wire used will consume 35.7 c.c. of the solution. 10 c.c. of the normal oxalic acid solution, or a solution made by dissolving 0.6285 gramme of oxalic acid in water, may also be used for titrating the permanganate solution, of which 100 c.c. should be decolorized. The titration is accomplished by diluting the 10 c.c. of oxalic acid solution with about 50 c.c. of water, to which a few c.c. of sulphuric acid are added, heating moderately, and adding the permanganate solution as in the above instance. Permanganate is generally used in determinations of iron and iron compounds. Many of the latter contain iron in the ferric state, which must be converted into ferrous compounds before titration. This con- version is accomplished by heating the solution of a weighed quantity of the ferric compound with nascent hydrogen—i. e., with metallic zinc and dilute sulphuric acid—in a flask arranged as the one spoken of above, and shown in Fig. 37. A very much quicker reduction of the ferric into a ferrous compound METHODS FOR QUANTITATIVE DETERMINATIONS. 269 may be accomplished by adding very slowly with constant stirring a saturated solution of sodium sulphite to the boiling, acidified iron solu- tion contained in the flask until the liquid becomes colorless. All excess of sulphur dioxide is expelled before titrating, by boiling the solution (which should contain a sufficient quantity of hydrochloric acid to decompose all sodium sulphite) for about ten minutes in a flask, arranged as the one mentioned above. One c.c. of deci-normal potassium permanganate is the equiva- lent of: Gramme. Iron, in ferrous compounds, Fe . 0.005588 Ferrous carbonate, FeC03 ..... . 0.011573 Ferrous oxide, FeO . 0.007195 Ferrous sulphate, FeS04 . 0.015170 Ferrous sulphate, crystallized, FeS04.7H20 . . 0.027753 Hydrogen dioxide, H202 . 0.001696 Oxalic acid, crystallized, H2C204.2H20 . . 0.006285 Potassium dichromate, K2Cr207 = 293.78. Whenever this salt acts in the presence of free acid, as an oxidizing agent, it tranfers 3 atoms of oxygen upon the deoxidizing agent, thus : K2Cr207 + 6FeS04 + 7H2S04 = K2S04 + Cr23S04 + 7H20 + 3(Fe23S04). A normal solution should, therefore, contain one-sixth of the molec- ular weight, or 48.963 grammes per litre. The U. S. P., however, considers potassium dichromate as a dibasic salt, using, therefore, one- half the molecular weight, 146.89, and prepares a deci-normal solution, which is obtained by dissolving 14.689 grammes of pure potassium dichromate in water, and diluting to 1000 c.c. The disadvantage of this solution is, that the final point of titration cannot be well seen, for which reason, in the determination of iron, for which it is chiefly used, the end of the reaction is determined by taking out a drop of the solution and testing it on a wThite porcelain plate with a drop of freshly prepared potassium ferricyanide solution; when this no longer gives a blue color, the reaction is at an end. In all determinations by this solution dilute sulphuric acid has to be added, because both the potassium and chromium require an acid to combine with, as shown in the above equation. One c.c. potassium dichromate solution, containing of this salt 0.014689 gramme, is the equivalent of: Gramme. Iron, in ferrous combinations, Fe . . 0.016764 Ferrous carbonate, FeC03 . 0.034719 Ferrous sulphate, crystallized, FeS04.7H20 . . 0.083226 Potassium hydroxide, KOH 1 .... . 0.005599 1 With phenol-phthale'in as indicator. 270 ANALYTICAL CHEMISTRY. Iodimetry. Solutions of iodine and of sodium thiosulphate (hypo- sulphite) act upon one another with the formation of sodium iodide and sodium tetrathionate: 21 + 2Na2S203 = 2NaI + Na2S406. A normal solution of one can be standardized by a normal solution of the other. As indicator is used starch solution, which is colored blue by minute portions of free iodine. Starch solution is made by mixing 1 gramme of starch with 10 c.c. of cold water, and then adding enough boiling water, under constant stirring, to make about 200 c.c. of a transparent jelly. If the solution is to be preserved for any length of time, 10 grammes of zinc chloride should be added. Many other substances, such as sulphurous acid, hydrogen sulphide, arsenous oxide, act upon iodine with the formation of colorless com- pounds, and may, therefore, be estimated by normal solution of iodine, while the iodine may be standardized by the thiosulphate solution. In many cases the latter solution is also used for the determination of chlorine, which is caused to act upon potassium iodide, the liberated iodine being titrated. Deci-normal iodine solution is the one generally used, and is made by dissolving 12.653 grammes of pure iodine in a solution of 18 grammes of potassium iodide in about 900 c.c. of water, diluting the solution to 1000 c c. To the article to be estimated by this solution is added a little starch solution, and then the iodine solution until, on stirring, the blue color ceases to be discharged. One c.c. of deci-normal iodine solution, containing of iodine 0.012653 gramme, is the equivalent of: Arsenous oxide, As203 Gramme. 0.004942 Potassium sulphite, crystallized, K2S03.2H20 0.009692 Sodium bisulphite, NaHS03 0.005193 Sodium hyposulphite (thiosulphate), Na2S203.5H20 0.024764 Sodium sulphite, crystallized, Na2S03.7H20 . 0.012579 Sulphur dioxide, S02 0.003195 Tartrate of antimony and potassium,2KSb0C4H406.H20 0.016560 Sodium thiosulphate (Hyposulphite). The crystallized salt, Na2S203. 5H20 = 247.64, is used for making the deci-normal solution by dis- solving 24.764 grammes of the pure crystallized salt in water to make METHODS FOR QUANTITATIVE DETERMINATIONS. 271 1000 c.c. If the salt should not be absolutely pure, a somewhat larger quantity (30-32 grammes) should be dissolved in 1000 c.c. of water, and this solution titrated with deci-normal solution of iodine and diluted with a sufficient quantity of water to obtain the deci- normal solution. The article to be tested, containing free iodine, either in itself or after the addition of potassium iodide, is treated with this solution until the color of iodine is nearly discharged, when a little starch liquor is added, and the addition! of the solution continued until the blue color has just disappeared. One c.c. of deei-normal solution of sodium thiosulphate, containing of the crystallized salt 0.024764 gramme, is the equivalent of: Gramme. Bromine, Br . 0.007976 Chlorine, Cl ... . . 0.008537 Iodine, I . 0.012653 Iron, Fe, in ferric salts . . 0.005588 Deci-normal solution of silver. The pure, dry crystallized silver nitrate, AgNOs = 169.55, is used for this solution, which is made by dissolving 16.955 grammes of the salt in water to make 1000 c.c. The standard of this solution may be found by means of a deci-normal solution of sodium chloride, containing of this salt 5.837 grammes in one litre. Volumetric silver solution is used directly for the estimation of most chlorides, iodides, bromides, and cyanides, including the free acids of these salts. Insoluble chlorides must first be converted into a soluble form by fusing them with sodium hydroxide, dissolving the fused mass (containing sodium chloride) in water, filtering and neu- tralizing with nitric acid. The hydroxides and carbonates of alkali metals and of alkaline earths may be converted into chlorides by evaporation to dryness with pure hydrochloric acid, and heating to about 120° C. (248° F.). The chlorides thus obtained may be titrated with silver solution. In the case of chlorides, iodides, and bromides, normal potassium chromate is used as an indicator. This salt forms with silver nitrate a red precipitate of silver chromate, but not before the silver chloride (bromide or iodide) has been precipitated entirely. In case free acids are determined by silver, these are neutralized with sodium hydroxide before titration. The operation is conducted as follows : The weighed quantity of the chloride is dissolved in 50-100 c.c. of water, neutralized if neces- ANALYTICAL CHEMISTRY. sary, mixed with a little potassium chromate, and silver solution added from the burette until a red coloration is just produced, which does not disappear on shaking. In estimating cyanides, the operation can be conducted as above described, or it can be modified, use being made of the formation of soluble double cyanides of silver and an alkali metal. The reaction takes place thus: 2KCN + AgN03 = AgK(CN’)2 + KN03. If to this soluble double compound more silver nitrate be added, it is decomposed with the formation of a precipitate of silver cyanide: AgK(CN)2 + AgN03 = 2AgCN + KN03. The estimation of hydrocyanic acid or of simple cyanides, accord- ing to this method, is accomplished by first rendering slightly alkaline the solution of the substance to be examined by the addition of sodium hydroxide, and then adding the silver solution until a perma- nent cloudiness is produced in the liquid, which shows that all cyan- ogen present has been converted into the soluble double salt. As but one-half of the silver solution has been added which is needed for the complete conversion of the cyanogen present into silver cya- nide, the number of c.c. of the standard silver solution employed will indicate exactly one-half of the equivalent amount of cyanide present in the solution. One c.c. of deci-normal silver nitrate solution, containing 0.016955 gramme of AgNOa, is the equivalent of: Ammonium bromide, NH4Br Gramme. 0.009777 Ammonium chloride, NH4CI 0.005338 Ammonium iodide, NH4I 0.014454 Ferrous bromide, FeBr2 . • 0.010775 Ferrous iodide, Fel2 ....... 0.015425 Hydriodic acid, HI . 0.012753 Hydrobromic acid, HBr 0.008076 Hydrochloric acid, HOI 0.003637 Hydrocyanic acid, HON,to first formation of precipitate 0.005396 Hydrocyanic acid, HCN, with indicator 0.002698 Potassium bromide, KBr 0.011879 Potassium chloride, KC1 0.007440 Potassium cyanide, KCN, to first formation of precipi- tate .......... 0.013002 Potassium cyanide, KCN, with indicator 0.006501 Potassium iodide, KI 0.016556 Sodium bromide, NaBr ....... 0.010276 Sodium chloride, NaCl 0.005837 METHODS FOB QUANTITATIVE DETERMINATIONS. 273 Gramme. Sodium iodide, Nal 0.014953 Zinc bromide, ZnBr2 0.011231 Zinc chloride, ZnCi2 0.006792 Zinc iodide, Znl2 0.015908 Deci-normal solution of sodium chloride is made by dissolving 5.837 grammes of pure sodium chloride in enough water to make 1000 c.c. The titration is made in neutral solution, normal potassium chromate being used as an indicator. (See explanation in previous paragraph on silver solution.) One c.c. of deci-normal sodium chloride solution, containing 0.005837 of NaCl, is the equivalent of: Gramme. Silver, Ag 0.010766 Silver nitrate 0.016955 Silver oxide 0.011564 Gas-analysis. The analysis of gases is generally accomplished by measur- ing gas volumes in graduated glass tubes (eudiometers) over mercury (in some cases over water), noting carefully the pressure and temperature at which the volume is determined. From gas mixtures, the various constituents present may often be eliminated by causing them to be absorbed one after another by suitable agents. For instance: From a measured volume of a mixture of nitrogen, oxygen, and carbon dioxide, the latter compound may be removed by allowing the gas to remain in contact for a few hours with potassium hydroxide, which will absorb all carbon dioxide, the diminution in volume indicating the quantity of carbon dioxide originally present. The volume of oxygen may next be determined by introducing a piece of phosphorus, which will gradually absorb the oxygen, the remaining volume being pure nitrogen. In some cases gaseous constituents of liquids or solids are eliminated and measured as gases. Thus, the carbon dioxide of carbonates, the nitrogen dioxide evolved from nitrates, the nitrogen of urea and other nitrogenous bodies, are instances of substances which are eliminated from solids in the gaseous state and determined by direct measurement. The gas volume thus found is, in most cases, converted into parts by weight The basis of this calculation is the weight of 1 c.c. of hydrogen, which, at the temperature of 0°C. (32° F.) and a pressure of 760 mm., is 0.0000896 gramme. 1 c.c. of any other gas weighs as many more times as the molecule of this substance is heavier than that of hydrogen. Thus the molecular weight of carbon dioxide is 22 times greater than that of hydrogen, consequently 1 c.c. of carbon dioxide weighs 22 times heavier than 1 c.c. of hydrogen, or 0.0019712 gramme. It has been shown on pages 21 and 25 that heat and pressure cause a regular increase or decrease in volume. The data there given are used in calculating the volume of the measured gas for the temperature of 0° C (32° F.) and a pressure of 760 mm. 274 ANALYTICAL CHEMISTRY. Questions.—361. Explain the principles which are made use of in gravi- metric and volumetric determinations. 362. Give an outline of the operations to be performed in the gravimetric determination of copper in cupric sulphate. 363. What are normal and deci-normal solutions, and how are they made? 364. What is the use of indicators in volumetric analysis? Mention some indicators and explain their action. 365. Why is oxalic acid preferred in preparing normal acid solution ? What quantity of oxalic acid is contained in a litre, and why is this quantity used? 366. Suppose 2 grammes of crys- tallized sodium carbonate require 14 c.c. of normal acid for neutralization : What are the percentages of crystallized sodium carbonate and of pure sodium carbonate contained in the specimen examined? 367. Ten grammes of dilute hydrochloric acid require 35.5 c.c. of normal sodium hydroxide solution for neutralization. What is the strength of this acid? 368. Explain the action of potassium permanganate and of potassium dichromate when used for volu- metric purposes. 369. Which substances may be determined volumetrically by solutions of iodine and sodium thiosulphate? Explain the mode in which the determinations by these agents are accomplished. 370. Suppose 1 gramme of potassium iodide requires for titration 60 c.c. of deci-normal solution of silver nitrate : What quantity of pure potassium iodide is indicated by this determination ? VI. CONSIDERATION OF CARBON COMPOUNDS, OR ORGANIC CHEMISTRY. 38. INTRODUCTORY REMARKS. ELEMENTARY ANALYSIS. Definition of organic chemistry. The term organic chemistry was originally applied to the consideration of compounds formed in plants and in the bodies of animals, and these compounds were believed to be created by a mysterious power, called “ vital force/’ supposed to reside in the living organism. This assumption was partly justified by the failure of the earlier attempts to produce these compounds by artificial means, and also by the fact that the peculiar character of the compounds, and the numer- ous changes which they constantly undergo in nature, could not be sufficiently explained by the experimental methods then known, and the laws then established. It was in accordance with these views that a strict distinction was made between inorganic and organic compounds, and accordingly between inorganic and organic chemistry, the latter branch of the science considering the substances formed in the living organism, and those compounds which were produced by their decomposition. Since that time it has been shown that many substances which formerly were believed to be exclusively produced in the living organ- ism, under the influence of the so-called vital force, can be formed artificially from inorganic matter, or by direct combination of the elements. It was in consequence of this fact that the theory of the supposed “ vital force,” by which organic substances could be formed exclusively, had to be abandoned. An organic compound, according to modern views, is simply a compound of carbon generally containing hydrogen, frequently also oxygen and nitrogen, and sometimes other elements. Organic chemistry may consequently he defined as the chemistry of 276 CONSIDERATION OF CARBON COMPOUNDS. carbon compounds. The old familiar terms, organic compounds and organic chemistry, are, however, still in general use. In a strictly systematically arranged text-book of chemistry organic compounds should be considered in connection with the element carbon itself, but as these carbon compounds are so numerous, their composi- tion often so complicated, and the decompositions which they suffer under the influence of heat or other agents so varied, it has been found best for purposes of instruction to defer the consideration of these compounds until the other elements and their combinations have been studied. Elements entering into organic compounds. Organic compounds contain generally but a small number of elements. These are, besides carbon, chiefly hydrogen, oxygen, and nitrogen, and sometimes sulphur and phosphorus. Other elements, however, enter occasionally into organic compounds, and by artificial means all metallic and non- metallic elements may be made to enter into organic combinations. Here the question presents itself: Why is it that the four elements carbon, hydrogen, oxyen, and nitrogen are capable of producing such an immense number (in fact, millions) of different combinations? To this question but one answer can be given, which is that these four elements differ more widely from each other, in their chemical and physical properties, than perhaps any other four elements. Carbon is a black, solid substance, which has never yet been fused or volatilized, while hydrogen, oxygeu, and nitrogen are colorless gases which can only be converted into liquids with difficulty. More- over, hydrogen is very combustible, oxygen is a supporter of combus- tion, whilst nitrogen is perfectly indifferent. Finally, hydrogen is univalent, oxygeu bivalent, nitrogen trivalent, and carbon quadri- valent. These elements are, therefore, capable of forming a greater number and a greater variety of compounds than would be the case if they were elements of equal valence and of similar properties. It will be shown,later that carbon atoms have, to a higher degree than the atoms of any other element, the power of combining with one another by means of a portion of the affinities possessed by each atom, thus increasing the possibilities of the formation of complex compounds. General properties of organic compounds. The substances formed by the union of the four elements just mentioned have properties in some respects intermediate to those of their components. Thus, no IN TRODUCTORY REMARKS. 277 organic substance is either permanently solid 1 like carbon, nor an almost permanent gas like hydrogen, oxygen, and nitrogen. Some organic substances are solids, others liquids, others gases; they are generally solids when the carbon atoms predominate; they are liquids or gases when the gaseous elements, and especially hydro- gen, predominate; likewise, it may also be said that compounds con- taining a small number of atoms in the molecule are gases or liquids which are easily volatilized; they are liquids of high boiling-points, or solids, when the number of atoms forming the molecules is large. Th e combustible property of carbon and hydrogen is transferred to all organic substances, every one of which will burn when sufficiently heated in atmospheric air. (If carbon dioxide, carbonic acid and its salts be considered organic compounds, we have an exception to the rule, as they are not combustible.) The properties possessed by organic compounds are many and widely different. There are organic acids, organic bases, and organic neutral substances ; there are some organic compounds which are per- fectly colorless, tasteless, and odorless, whilst others show every possi- ble variety of color, taste, and odor; many serve as food, whilst others are most poisonous; in short, organic substances show a greater variety of properties than the combinations formed by any other four elements. And yet, the cause of all the boundless variety of organic matter is that peculiar attraction called chemical affinity, acting between the atoms of a comparatively small number of elements and uniting them in many thousand different proportions. It would, of course, be entirely inconsistent with the object of this book, if all the thousands of organic substances already known (the number of which is continually being increased by new discoveries) were to be considered or even mentioned. It must be sufficient to state the general properties of the various groups of organic substances, to show by what processes they are produced artificially or how they are found in nature, how they may be recognized and separated, and, finally, to point out those members of each group which claim a special attention for one reason or another. Difference in the analysis of organic and inorganic substances. The analysis of organic substances differs from that of inorganic sub- 1 Non-volatile organic substances are decomposed by heat with generation of volatile products. 278 CONSIDERATION OF CARBON COMPOUNDS. stances, in so far as the qualitative examination of an organic substance furnishes in many cases but little proof of the true nature of the substauce (except that it is organic), whilst the qualitative analysis of an inorganic substance discloses in most cases the true nature of the substance at once. For instance: If a white, solid substance, upon examination, be found to contain potassium and iodine, and nothing else, the conclu- sion may at once be drawn that the compound is potassium iodide, containing 39 parts by weight of potassium, and 126.5 parts by weight of iodine. Or, if another substance be examined, and found to be composed of mercury and chlorine, the conclusion may be drawn that the compound is either mercurous or mercuric chloride, as no other compounds containing these two elements are known, and >vhether the examined substance be the lower or higher chloride of mercury, or a mixture of both, can easily be determined by a few simple tests. Whilst thus the qualitative examination discloses the nature of the substance, it is different with organic compounds. Many thousand times the analysis might show carbon, hydrogen, and oxygen to be present, and yet every one of the compounds examined might be entirely different; it is consequently not only the quality of the ele- ments, but chiefly the quantity present which determines the nature of an organic substance, and in order to identify an organic substance with certainty, it frequently becomes necessary to make a quantitative determination of the various elements present, and this quantitative analysis by which the elements in organic substances are determined is generally called ultimate or elementary analysis. There are, however, for many organic substances such character- istic tests that these substances may be recognized by them; these reactions will be mentioned in the proper places. An analysis by which different organic substances, when mixed together, are separated from each other is frequently termed proximate analysis. Such an analysis includes the separation and determination of essential oils, fats, alcohols, sugars, resins, organic acids, albuminous substances, etc., and is one of the most difficult branches of analytical chemistry. Qualitative analysis of organic substances. The presence of carbon in a combustible form is decisive in regard to the organic nature of a compound. If, consequently, a substance burns with generation of ELEMENTARY ANALYSIS. 279 carbon dioxide (which may be identified by passing the gas through lime-water), the organic nature of this substance is established. The presence of hydrogen can be proven by allowing the gaseous products of the combustion to pass through a cool glass tube, when drops of water will be deposited. It is difficult to show by qualitative analysis the presence or absence of oxygen in an organic compound, and its determination is therefore generally omitted. The presence of nitrogen is determined by heating the substance with dry soda-lime (a mixture of two parts of calcium hydroxide and one part of sodium hydroxide), when the nitrogen is converted into ammonia gas, which may be recognized by its odor or by its action on paper moistened with solution of cupric sulphate, a dark-blue color indicating ammonia. Ultimate or elementary analysis. While the student must be referred to books on analytical chemistry for a detailed description of the apparatus required and the methods employed for elementary analysis, it may here be stated that the quantitative determination of carbon and hydrogen is generally accomplished by the following pro- cess : A weighed quantity of the pure and dry substance is mixed Fig. 38. Gas-furnace for organic analysis. with a large excess of dry cupric oxide, and this mixture is introduced into a glass tube, the open end of which is connected by means of a perforated cork and tubing with two glass vessels, the first one of which (generally a U-shaped tube) is filled with pieces of calcium chloride, the other (usually a tube provided with several bulbs) with solution of potassium hydroxide. The two glass vessels, containing 280 CONSIDERATION OF CARBON COMPOUNDS. the substances named, are weighed separately after having been filled. Upon heating the combustion-tube in a suitable furnace, the organic matter is burned by the oxygen of the cupric oxide, the hydrogen is converted into water (steam), which is absorbed by the calcium chloride, and the carbon is converted into carbon dioxide, which is absorbed by the potassium hydroxide. The apparatus repre- sented in Fig. 38 shows the gas-furnace in which rests the combustion- tube with calcium chloride tube and potash bulb attached. Upon re-weighing the two absorbing vessels at the end of the operation, the increase in weight will indicate the quantity of water and carbon dioxide formed during the combustion, and from these figures the amount of carbon and hydrogen present in the organic matter may easily be calculated. For instance: 0.81 gramme of a substance having been analyzed, furnishes, of carbon dioxide 1.32 gramme, and of water 0.45 gramme. As every 44 parts by weight of carbon dioxide contain 12 parts by weight of carbon, the above 1.32 gramme contains of carbon 0.3fi gramme, or 44.444 per cent. As every 18 parts of water contain 2 parts of hydrogen, the above 0.45 gramme consequently contains 0.05 gramme, or 6.172 per cent. Oxygen is scarcely ever determined directly, but generally indi- rectly, by determining the quantity of all other elements and deduct- ing their weight, calculated to percentages from 100. The difference is oxygen. If, in the above instance, 44.444 per cent, of carbon aud 6.172 per cent, of hydrogen were found to be present, and all other elements, except oxygen, to be absent, the quantity of oxygen is, then, equal to 49.384 per cent, and the composition of the substance is as follows: is oxygen. Carbon .... Hydrogen .... . 6.172 “ Oxygen .... . 49.384 “ 100.000 Determination of nitrogen. Nitrogen is generally determined by heating the substance with soda-lime and passing the generated ammonia gas through hydrochloric acid, contained in a suitable glass vessel. Upon evaporation of the acid solution in a weighed platinum dish over a water-bath, ammonium chloride is left, from the weight of which compound the quantity of nitrogen may be calculated. Or the ammonia gas may be passed through a measured volume of normal hydrochloric acid and the unsaturated portion of the acid determined volumetrically. ELEMENTARY ANALYSES. 281 Determination of sulphur and phosphorus. These elements are determined by mixing the organic substance with sodium carbonate and nitrate, and heating the mixture in a crucible. The oxidizing action of the nitrate converts all carbon into carbon dioxide, hydrogen into water, sulphur into sulphuric acid, phosphorus into phosphoric acid. The latter two acids combiue with the sodium of the sodium carbonate, forming sulphate and phosphate of sodium. The fused mass is dissolved in water, and sulphuric acid precipitated by barium chloride, phosphoric acid by magnesium sulphate and ammonium hydroxide and chloride. From the weight of barium sulphate and magnesium pyrophosphate the weight of sulphur and phosphorus is calculated. Determination of atomic composition from results obtained by elementary analysis. The elementary analysis gives the quantity of the various elements present in percentages, and from these figures the relative number of atoms may be found by dividing the figures by the respective atomic weights. For instance: The analysis above men- tioned gave the composition of a compound, as carbon 44.444 per cent., hydrogen 6.172 per cent., and oxygen 49.384 per cent. By dividing each quantity by the atomic weight of the respective element, the following results are obtained : 44 444 . = 3.703 12 61172 = 6.172 1 49 384 = 3.087 16 The figures 3.703, 6.172, and 3.087, represent the relative number of atoms present in a molecule of the compound examined. In order to obtain the most simple proportion expressing this relation, the greatest divisor common to the whole has to be found, a task which is sometimes rather difficult on account of slight errors made in the quantitative determination itself. In the above case, 0.6172 is the greatest divisor, which gives the following results : 3.7°3 _ 6. 06172 ~~ ’ 6.172 _ 10. 0.6172 “ ’ 3.G87 _ g 0.6172 — The simplest numbers of atoms are, accordingly, carbon 6, hydrogen 10, oxygen 5, or the composition is C6H10O5. Empirical and molecular formulas. A chemical formula is termed empirical when it merely gives the simplest possible expression of the 282 CONSIDERATION OF CARBON COMPOUNDS. composition of a substance. In the above case, the formula C6II10O5 would be the empirical formula. It might, however, be possible that this formula did not represent the actual number of atoms in the molecule, which might contain, for instance, twice or three times the number of atoms given, in which case the true composition would be expressed by the formula C12H20O10 or C18H30O15. If it could be proven that one of the latter formulas is the correct one, it would be termed the molecular formula, because it expresses not only the numerical relations existing between the atoms, but also the absolute number of atoms of each element contained in the molecule. The best method for determining the actual number of atoms con- tained in the molecule is the determination of the specific weight of the gaseous compound, taking hydrogen as the unit. For instance : Assume the analysis of a liquid substance gave the following result: Carbon 92.308 per cent. Hydrogen 7.692 “ 100.000 From this result the empirical formula, CH, is deducted by apply- ing the method stated above. If this formula were the molecular formula, the density of the vapors of the substance would, when com- pared with hydrogen (according to the law of Avogadro), be equal to 6.5, because a molecule of hydrogen weighs 2 and a molecule of the compound CH weighs 13. Suppose, however, the density of the gaseous substance is found to be 39, then the molecular formula would be expressed by C6H6, because its molecular weight (6 X 12 + 6 X 1) is equal to 78, which weight, when compared with the molecular weight of hydrogen = 2, gives the proportions 78 : 2, or 39 : 1. Not all organic compounds can be converted into gases or vapors without undergoing decomposition, and the determination of the molecular formulas of such compounds has to be accomplished by other methods. If the substance, for instance, is an acid or a base, the molecular formula may be determined by the analysis of a salt formed by these substances. For instance : The empirical formula of acetic acid is CII20 ; the analysis of the potassium acetate, however, shows the composition KC2H302, from which the molecular formula HC2H302 is deducted for acetic acid. In many cases, however, it is as yet absolutely impossible to give with certainty the molecular formula of some compounds. ELEMENTARY ANALYSIS. 283 Rational, constitutional, structural, or graphic formulas. These formulas are intended to represent the theories which have been formed in regard to the arrangement of the atoms within the molecule, or to represent the modes of the formation and decomposition of a compound, or the relation which allied compounds bear to one another. The molecular formula of acetic acid, for instance, is C2H402, but different constitutional formulas have been used to represent the struc- ture of the acetic acid molecule. Thus, H.C2H302 is a formula analogous to H.N03, indicating that acetic acid (analogous to nitric acid), is a monobasic acid, containing one atom of hydrogen, which can be replaced by metallic atoms. C2H301.0Hi is a formula indicating that acetic acid is composed of two univalent radicals which may be taken out of the molecule and replaced by other atoms or groups of atoms. This formula indicates also that acetic acid is analogous to hydroxides, the radical C2HaO having replaced one atom of hydrogen in H20. CHi3.C02H1 is a formula indicating that acetic acid is composed of the two compound radicals, methyl and carboxyl. It may be said finally, that quite a number of other rational formulas have been applied, or, at least, have been proposed by differ- ent chemists and at different times, to represent the structure of acetic acid, but it should be remembered that these formulas are not intended to represent the actual arrangement of the atoms in space, but only, as it were, their relative mode of combination, showing which atoms are combined directly and which only indirectly, that is, through the medium of others. Questions.—371. What is organic chemistry, according to modern views? 372. Mention the chief four elements entering into organic compounds, and name the elements which may be made to enter into organic compounds by artificial processes. 373. State the reason why the four elements, carbon, hydrogen, oxygen, and nitrogen, are more apt to form a larger number of compounds than most other elements. 374. State the general properties of organic compounds. 375. Why does a qualitative analysis of an organic com- pound, in most cases, not disclose its true nature? 376. By what test may the organic nature of a compound be established ? 377. By what tests may the presence of carbon, hydrogen, and nitrogen be demonstrated in organic com- pounds? 378. State the methods by which the elements carbon, hydrogen, oxygen, sulphur, and phosphorus are determined quantitatively. 379. By what general method may a formula be deducted from the results of a quanti- tative analysis? 380. What is meant by an empirical, molecular, and consti- tutional formula; how are they determined, and what is the difference between them ? 284 CONSIDERATION OF CARBON COMPOUNDS. 39. CONSTITUTION, DECOMPOSITION, AND CLASSIFICATION OF ORGANIC COMPOUNDS. Radicals or residues. The nature of a radical or residue has been stated already in Chapter 8, but the important part played by rad- icals in organic compounds renders it necessary to consider them more fully. A radical is an unsaturated group of atoms obtained by removal of one or more atoms from a saturated compound. It is not neces- sary that this removal of atoms should be practically accomplished in order to call a group of atoms a radical, but it is sufficient to prove that the unsaturated group of atoms exists as such in a number of compounds, and that it can be transferred from one compound into another without suffering decomposition. Radicals exist in organic and inorganic compounds; an inorganic radical spoken of heretofore is the water residue or hydroxyl, OH, obtained by removal of one atom of hydrogen from one molecule of water. Hydroxyl does not exist in the separate state, but it exists in hydrogen dioxide, H202, or HO—OH, and is also a constituent of the various hydroxides, as, for instance, of KOH, Ca(OH)2, Fe2(OH)6, etc. If one atom of hydrogen be removed from the saturated hydro- carbon methane, CH4, the univalent residue methyl, CH3, is left, which is capable of combining with univalent elements, as in methyl chloride, CH3C1, or, with univalent residues, as in metyl hydroxide, ch3oh. If two atoms of hydrogen be removed from CH4, the bivalent resi- due methylene, CH2, is left, capable of forming the compounds CH2C12, CH2(OH)2, etc. If three atoms of hydrogen be removed from CH4, the trivalent residue CH is left, capable of combining with three atoms of univa- lent elements, as in CHC13, or with another trivalent radical, etc. Chains. The expression, chain, designates a series of multivalent atoms (generally, but not necessarily, of the, same element), held together in such a manner that affinities are left unsaturated. For instance: —0—0—, —0—0—0—, _0—0—0—0—, are oxygen chains, each one of which has two free affinities which may be saturated, for instance, with the following results : II—0—0—II, H—0—0—0—Cl, H—0—0—0—0—Cl, Hydrogen peroxide. Chloric acid. Perchloric acid. CONSTITUTION OF ORGANIC COMPOUNDS. 285 In a similar manner, carbon atoms unite, forming chains, as, for instance : I I —C—c—, I I -UJ-, I I I _0 + 2C02. c. Two molecules, either of the same kind, or of different sub- stances, may unite together directly : Grape-sugar. Alcohol. Carbon dioxide. C2H4 + 2Br = C2H4Br2. d. Atoms may be removed from a compound without replacing them by other atoms : Ethylene. Bromine. Ethylene bromide. c2hbo + o = c2h4o + h2o. e. Atoms may be removed and replaced by others at the same time (substitution): Alcohol. Oxygen. Aldehyde. Water. C2H402 + 2C1 = C2H3C102 + HC1 Acetic acid. Chlorine. Monochloracetic acid. Hydrochloric acid. Action of heat upon organic substances. As a general rule, organic bodies are distinguished by the facility with which they decompose under the influence of heat or chemical agents; the more complex the body is, the more easily does it undergo decomposition or transforma- tion. Heat acts differently upon organic substances, some of which may be volatilized without decomposition, whilst others are decomposed by heat with generation of volatile products. This process of heating non-volatile organic substances in such a manner that the oxygen of the atmospheric air has no access, and to such an extent that decom- position takes place, is called dry or destructive distillation. The nature of the products formed during this process varies not DECOMPOSITION OF ORGANIC COMPOUNDS. 289 only with the nature of the substance heated, but also with the tem- perature applied during the operation. The products formed by destructive distillatiou are invariably less complex in composition, that is, have a smaller number of atoms in the molecule, than the substance which suffered decomposition ; in other words, a complex molecule is split up into two or more molecules less complex in com- position. Otherwise, the products formed show a great variety of properties; some are gases, others volatile liquids or solids, some are neutral, others basic or acid substances. In most cases of destructive distillation a non-volatile residue is left, which is nearly pure carbon. Action of oxygen upon organic substances. Combustion. Decay. All organic substances are capable of oxidation, which takes place either rapidly with the evolution of heat and light and is called com- bustion, or it takes place slowly without the emission of light, and is called slow combustion or decay. The heat generated during the decay of a substance is the same as that generated by burning the substance; but as this heat is liberated in the first instance during weeks, months, or perhaps years, its generation is so slow that it can scarcely be noticed. No organic substauce found or formed in nature contains a suffi- © cient quantity of oxygen to cause the complete combustion of the combustible elements (carbon and hydrogen) present; by artificial processes such substances may, however, be produced, and are then either highly combustible or even explosive. During common combustion, provided an excess of atmospheric oxygen be present, the total quantity of carbon is converted into carbon dioxide, hydrogen into water, sulphur and phosphorus into sulphuric and phosphoric acids, while nitrogen is generally liberated in the elementary state. During the process of decay the compounds mentioned above are produced finally, although many intermediate products are generated- For instance : If a piece of wood be burnt, complete oxidation takes place; intermediate products also are formed chiefly in consequeuce of the destructive distillation of a portion of the wood, but they are consumed almost as fast as they are produced, as was mentioned in connection with the consideration of flame. Again, when a piece of wood is exposed to the action of the atmosphere, it slowly burns or decays. The intermediate products formed in this case are entirely different from those produced during common combustion. 290 CONSIDERATION OF CARBON COMPOUNDS. Common alcohol has the composition C2H60; in burning, it re- quires six atoms of oxygen, when it is converted into carbon dioxide and water : C2H60 + 60 = 2C02 + 3H20. But alcohol may also undergo slow oxidation, in which case oxygen first removes hydrogen, with which it combines to form water, whilst at the same time a compound known as acetic aldehyde, C2H40, is formed: CaH60 + O = C2H40 + H20. This aldehyde, when further acted upon by oxygen, takes up an atom of this element, thereby forming acetic acid : C2H40 + 0 = C2H402. The three instances given above illustrate the action of oxygen upon organic substances, which action may consist in a mere removal of hydrogen, in a replacement of hydrogen by oxygen, or in an oxida- tion of both the carbon and hydrogen, and also of sulphur and phos- phorus, if they be present. An organic substance, when perfectly dry and exposed to dry air only, may not suffer decay for a long time (not even for centuries), but in the presence of moisture and air this oxidizing action takes place almost invariably. Besides the slow oxidation or decay which all dead organic matter undergoes in the presence of moisture, there is another kind of slow oxidation, called respiration, which takes place in the living animal; this process will be more fully considered in the physiological part of this book. Fermentation and putrefaction. These terms are applied to pecu- liar kiuds of decomposition, by which the molecules of certain organic substances are split up into two or more molecules of a less compli- cated composition. These decompositions take place when three factors are simultaneously acting upon the organic substance. These factors are: presence of moisture, favorable temperature, and presence of a substance generally termed ferment. The most favorable temperature for these decompositions lies between 25° and 40° C. (77° and 104° F.), but they may take place at lower or higher temperatures. No substance, however, will either ferment or putrefy at or below the freezing-point, or at or above the boiling-point. The nature of the various ferments differs widely, and their true DECOMPOSITION OF ORGANIC COMPOUNDS. 291 action cannot, in many cases, be explained; what we do know is, that the presence of comparatively small (often minute) quantities of one substance (the ferment) is sufficient to cause the decomposition of large quantities of certain organic substances, the ferment itself suf- fering often no apparent change during this decomposition. Ferments may be divided into two classes : 1. Soluble ferments, which are in most cases nitrogenous substances, closely related to the proteids; 2. Living microbrganisms of either vegetable or animal origin. The nature of the ferment generally determines the nature of the decomposition which a substance suffers, or, in other words, one and the same substance will under the influence of one ferment decompose with liberation of certain products, while a second ferment causes other products to be evolved. Sugar, for instance, under the influence of yeast, is converted into alcohol and carbon dioxide, while under the influence of certain other ferments it is converted into lactic acid. The difference between fermentation and putrefaction is, that the first term is used in those cases where the decomposing substance con- tains carbon, hydrogen, and oxygen only, while substances containing, in addition to these three elements, either nitrogen or sulphur (or both) undergo putrefaction. The two last-named elements are generally evolved as ammonia or derivatives*of ammonia and hydrogen sulphide, which gases give rise to an offensive odor. Sugar, having the composition 06H12O6, undergoes fermentation, whilst albuminous substances which contain also nitrogen and sulphur putrefy. The oxygen of the air takes no part in either fermentation or putre- faction, but the presence or absence of atmospheric air may cause or prevent decomposition, inasmuch as the atmosphere is filled with millions of minute germs of organic nature, which germs may act as ferments when in contact with organic matter under favorable con- ditions. Whenever organic bodies (a dead animal, for instance) undergo decomposition in nature, the processes of fermentation and putrefac- tion are generally accompanied by oxidation or decay. The conditions under which a substance will ferment or putrefy have been stated above, and the non-fulfilment of these conditions enables us to prevent decomposition artificially. Thus, we freeze substances (meat); or expel all water from or dry them (fruit, etc.), in order to prevent decomposition. The action of the ferments is counteracted either by the so-called antiseptic agents (salt, carbolic or salicylic acid, etc.) Avhich are incompatible with 292 CONSIDERATION OF CARBON COMPOUNDS. organic life, or by excluding the air, and with it the ferments, by enclosing the substances in air-tight vessels (glass jars, tin cans, etc.), which, when filled, are heated sufficiently to destroy any germs which may have been present, and are then sealed. Antiseptics and disinfectants. While the term antiseptics is applied to those substances which retard or prevent fermentation and putre- faction, the term disinfectants refers to those agents actually destroying the organisms which are the causes of these decompositions. If we assume that all infectious diseases are due to microorganisms, or germs of various kinds, disinfectants may be considered as equivalent to germicides. Disinfectants are generally antiseptics also, but the latter are not in all cases disinfectants. The solution of a substance of certain strength may act as a disinfectant and antiseptic, while the same solution diluted further may act as an antiseptic only, but not as a disinfectant. Deodorizers are those substances which convert the strongly smell- ing products of decomposition into inodorous compounds. Strong oxidizing agents are generally good deodorizers, as, for instance, chlorine, potassium permanganate,* hydrogen dioxide, etc. Among the best antiseptics and disinfectants are chlorine (generally used in the form of a 4 per cent, solution of hypochlorite of calcium), mercuric chloride (a solution of 1 : 500 or 1 : 1000), carbolic acid (a 5 per cent, solution), and some other substances. Action of chlorine and bromine. These two elements act upon organic substances (similarly to oxygen) in three different ways, viz., they either (rarely, however) combine directly with the organic sub- stance, or remove hydrogen, or replace hydrogen. The following equations illustrate this action: Ethylene. C2H4 + 2l3r = C2H4Br2. Bromine. Ethylene bromide. Ethyl alcohol. C2H60 + 2C1 = C2H40 + 2HCI. Chlorine. Aldehyde. Hydrochloric acid. C2H402 + 2C1 = C2H3C102 + HC1 Acetic acid. Chlorine. Monochloracetic acid. Hydrochloric acid. Iu the presence of water, chlorine and bromine often act as oxidiz- ing agents by combining with the hydrogen of the water and liberating oxygen ; iodine may act in a similar manner as an oxidizing agent, but it rarely acts directly by substitution. DECOMPOSITION OF ORGANIC COMPOUNDS. 293 Action of nitric acid. This substance acts either by direct combi- nation with organic bases forming salts, or as an oxidizing agent, or by substitution of nitryl, N02, for oxygen. As instances of the latter action, may be mentioned the formation of nitro-benzene and nitro- cellulose : c6h6 + hno3 = c6h5no2 + h2o. Benzene. Nitric acid. Nitro-benzene. Water. C6H10O5 + 8HNO3 = C6H7(N02)305 + 3H20 Cellulose. Nitric acid. Trinitro-cellulose. Water. The additional quantity of oxygen thus introduced into the mole- cules, renders them highly combustible, or even explosive. Action of dehydrating agents. Substances having a great affinity for water, such as strong sulphuric acid, phosphoric oxide, and others, act upon many organic substances by removing from them the elements of hydrogen and oxygen, and combining with the water formed, while, at the same time, frequently dark or even black compounds are formed, which consist largely of carbon. The black color imparted to sul- phuric acid by organic matter depends on this action. Action of alkalies. The hydroxides of potassium and sodium act in various ways on organic substances. In some cases direct combination takes place : CO + KOH = KCH02. Carbonic oxide. Potassium hydroxide. Potassium formate. Salts are formed: C2H402 + NaOH = NaC2H302 + H20. Acetic acid. Sodium hydroxide. Sodium acetate. Water, Fats are decomposed with the formation of soap C3H5(C18H3302)3 + 3NaOH = C3H.(HO)3 + 3NaC18H3302. Oxidation takes place, while hydrogen is liberated : Oleate of glyceril. Sodium hydroxide Glycerin. Sodium oleate. C2H60 + KOH = KC2H302 + 4H. Ethyl alcohol. Potassium hydroxide. Potassium acetate. Hydrogen, From compounds containing nitrogen, ammonia is evolved : NH2C2H30 + KOH = KC2H302 + NH3. Acetamide. Potassium hydroxide. Potassium acetate. Ammonia, Action of reducing agents. Deoxidizing or reducing agents, espe- cially hydrogen in the nascent state, act upon organic substances either by direct combination : 294 CONSIDERATION OF CARBON COMPOUNDS. C2H40 + 2H = C2H60 or byTremoving oxygen (and also chlorine or bromine): Ethene oxide. Ethyl alcohol. C7H602 + 2H = C7H60 + H,,0. In some cases hydrogen replaces oxygen : Benzoic acid. Benzoic aldehyde. C6H5N02 + 6H = C6H3NH2 + 2H20. Nitrobenzene. Aniline. Classification of organic compounds. There are great difficulties in arranging the immense number of organic substances properly, and in such a manner that natural groups are formed the members of which are similar in composition and possess like properties. Various modes of classification have been proposed, some of which, however, are so complicated that the beginuer will find it difficult to make use of them. The grouping of organic substances here adopted, while far from being perfect, has the advantages of being simple, easily understood, and remembered. 1. Hydrocarbons. All compounds containing the two elements carbon and hydrogen only. For instance, CH4, C6H6, C10H16, etc. 2. Alcohols. These are unsaturated hydrocarbons or hydrocarbon residues in combination with hydroxyl, OH. For instance, ethyl alcohol, C2Hi5OH, glycerin, C3HlU5 (OH)3, etc. 3. Aldehydes. Unsaturated hydrocarbons in combination with the radical COH; they are compounds intermediate between alcohols and acids, or alcohols from which hydrogen has been removed. For instance: C2H60, C2H40, c2h4o2. 4. Organic acids. Unsaturated hydrocarbons in combination with carboxyl, a radical having the composition C02H, or compounds formed by replacement of hydrogen in hydrocarbons by carboxyl. Otherwise, organic acids have the general properties of inorganic acids. 5. Ethers. Compounds formed from alcohols by replacement of the hydrogen of the hydroxyl by other unsaturated hydrocarbons, or, what is the same, by other alcohol radicals. For instance: Ethyl alcohol. Aldehyde. Acetic acid. c2ha0 H/u’ Ethyl alcohol. c2h5\ C2H5x ’ Ethyl ether. c.2h5\0 c Hs/U- Ethyl-methyl ether. 6. Compound ethers or esters. Formed from alcohols by replace- ment of the hydrogen of the hydroxyl by acid radicals, or from acids CLASSIFICATION OF ORGANIC COMPOUNDS. 295 by replacement of the hydrogen of carboxyl by alcohol radicals. For instance: 4- CH3CO\q C2H5\0 . H\0 H/u + Ily — CH3CO/ + H/U' Ethyl alcohol. Acetic acid. Acetic ether. Water. The various fats belong to this group of compound ethers. 7. Carbohydrates. (Sugars, starch, gum, etc.) These compounds contain 6 atoms of carbon (or a multiple of 6) in the molecule, and hydrogen and oxygen in the proportion of 2 atoms of hydrogen to 1 atom of oxygen, or in the proportion to form water. Most carbo- hydrates are capable of fermentation, or of being easily converted into fermentable bodies. Instances : C6H1206, C6H10O5, etc. Glucosides are substances the molecules of which may be split up in such a manner that several new bodies are formed, one of which is sugar. 8. Amines and amides. Substances formed by replacement of hydrogen in ammonia by alcohol or acid radicals. For instance: ethyl amine, NII2.C2H5, urea, H2II4.CO, etc. The alkaloids belong to this group. 9. Cyanogen and its compounds. Substances containing the radical cyanogen, CN. For instance : potassium cyanide, KCN. 10. Proteids or albuminous substances. These contain, besides carbon, hydrogen, aud oxygen, always nitrogen and sulphur, some- times also other elements. Instances : albumin, casein, fibrin, etc. In connection with each of these groups have to be considered the derivatives obtained from them directly or indirectly. As all those organic compounds the constitution of which has been explained, may be looked upon as derivatives of either methane, CII4, or benzene, C6II6, a separation of organic compounds is made into two large classes, each one embodying all the derivatives of one of the two hydrocarbons named. The derivatives of methane are termed fatty compounds, those of benzene aromatic compounds. Fatty compounds have representatives in each one of the above ten groups : aromatic compounds are missing in a few. As far as practicable, the two classes will be considered separately, because the properties of fatty and aromatic compounds differ so widely, in some respects, that this method of studying the nature of carbon compounds is to be preferred. Questions.—381. Explain the term residue or radical. 382. What is under- stood by the expression chain, when used in chemistry? 383. What are the characteristics of an homologous series ? 384. Give an explanation of the terms isomerism, metamerism, and polymerism. 385. How does heat act upon organic 296 CONSIDERATION OF CARBON COMPOUNDS. 40. HYDROCARBONS. Occurrence in nature. Hydrocarbons are seldom derived from animal sources, being generally products of vegetable life; thus, the various essential oils (oil of turpentine and others) of the composition Ci0H16 or C20H32 are frequently found in plants, where they are formed from carbon dioxide and water : 10CO2 4- 8H20 = C10H16 + 28 0. This equation, whilst showing the possibility of the formation of an essential oil in the plant, must not be taken to mean that 10 molecules of carbon dioxide and 8 molecules of water are simultaneously decom- posed, with the production of an essential oil; on the contrary, we know that many intermediate substances are formed, and the formula simply gives the final result, not the intermediate stages of the process. Other hydrocarbons are found in nature as products of the decom- position of organic matter. Thus methane, CH4, is generally formed during the decay of organic matter in the presence of moisture; the higher members of the methane series are found in crude coal-oil. Formation of hydrocarbons. It is difficult to combine the two elements carbon and hydrogen directly ; as an instance of such direct combination, may be mentioned acetylene, C2H2, which is formed when electric sparks pass between electrodes of carbon in an atmosphere of hydrogen. Many hydrocarbons are obtained by destructive distillation of organic matter, and their nature depends on the composition of the material used and upon the degree of heat applied for the decompo- sition. Hydrocarbons may also be obtained by the decomposition (other than destructive distillation) of numerous organic bodies, such as alcohols, acids, amines, etc., and from derivatives of these sub- stances. The hydrocarbons found in nature are generally separated from other matter, as well as from each other, by the process known as fractional distillation. As the boiling-points of the various compounds differ more or less, they may be separated by carefully distilling off compounds? 386. What is destructive distillation? 387. State the difference between combustion, decay, fermentation, and putrefaction ; what is the nature of these processes, and under what conditions do they take place? 388. How do chlorine, nitric acid, and alkalies act upon organic substances? 389. What is the action of hydrogen, and of dehydrating agents, upon organic substances ? 390. Mention the chief groups of organic compounds. HYDROCARBONS. 297 the compounds of lower boiling-points, while noting the temperature of the vapors above the boiling liquid by means of an inserted ther- mometer, and changing the receiver every time an increase of the boiling-point is noticed. This separation of volatile liquid, known as fractional distillation, is, however, not absolutely complete, because traces of substances having a higher boiling-point are simultaneously volatilized with the distilling substance. Fig. 39. Flasks arranged for fractional distillation For fractional distillation of small quantities of liquids as well as for the determination of boiling-points, flasks arranged like those shown in Fig. 39 may be used. Properties of hydrocarbons. There are no other two elements which combine together in so many proportions as carbon and hydro- gen. Several hundred hydrocarbons are known, many of which form either homologous series or are metameric or polymeric. Hydrocarbons occur either as gases, liquids, or solids. If the mole- cule contains not over 4 atoms of carbon, the compound is generally a gas at the ordinary temperature; if it contains from 4 to 10 or 12 298 CONSIDERATION OF CARBON COMPOUNDS. atoms of carbon, it is a liquid ; and if it contains a yet higher number of carbon atoms, it is generally a solid. Ail hydrocarbons may be volatilized without decomposition, all are colorless substances, and many have a peculiar and often characteristic odor; they are generally insoluble in water but soluble in alcohol, ether, disulphide of carbon, etc. In regard to chemical properties, it may be said that hydrocarbons are neutral substances, behaving rather indifferently toward most other chemical agents. Most of them are, however, oxidized by the oxygen of the air, by which process liquid hydrocarbons are often converted into solids. Hydrocarbons of the paraffin or methane series. The hydrocarbons having the general composition CuHn2+2 are known as paraffins, the name being derived from the higher members of the series which form the paraffin of commerce. The following table gives the composition, boiling-points, etc., of the first sixteen members of this series : Methyl hydride or methane, C H4 Ethyl hydride or ethane, C2 H6 Propyl hydride or propane, C3 H8 B. P. Sp. gr. gases. Butyl hydride or butane, C4 H10 1° C. Amyl hydride or pentane, C5 H12 38 0.628 Hexyl hydride or hexane, C6 Hu 70 0.669 Heptyl hydride or heptane, C7 H16 99 0.690 Octyl hydride or octane, C8 Hlg 125 0.726 Nonyl hydride or nonane, C9 H20 148 0.741 Decyl hydride or decane, C10H22 166 0.757 Undecyl hydride or undecane, CUH24 184 0.766 Dodecyl hydride or dodecane, C12H26 202 0.778 Tridecyl hydride or tridecane, C13H28 218 0.796 Tetradecyl hydride or tetradecane, C14H30 236 0.809 Pentadecyl hydride or pentadecane, C15H32 258 0.825 Hexadecyl hydride or hexadecane, C16H34 280 etc. The above table shows that the paraffins form an homologous series ; the first four members are gases, most of the others liquids, regularly increasing in specific gravity, boiling-point, viscidity, and vapor density, as their molecular weight becomes greater. The paraffins are saturated hydrocarbons, the constitution of which has been already explained ; they are incapable of uniting directly with monatomic elements or residues, but they easily yield substitu- tion-derivatives when subjected to the action of chlorine or bromine. HYDROCARBONS. 299 Most of the paraffins are known in two (or even more) modifications ; there are, therefore, other homologous series of hydrocarbons of the same composition as the above normal paraffins, which show some difference from the normal paraffins in boiling-points and other properties. In these isomeric paraffins the atoms are arranged differently from those in the normal hydrocarbons, which fact may be proven by the difference in decomposition which these substances suffer when acted upon by chemical agents. No isomeric hydrocarbons of the first three members of the paraffin series are known, which fact is in accordance with our present theories. Assuming that the quadrivalent carbon atoms exert their full valence, and that they are held together by one atomicity only, we can arrange the atoms in the compounds, CH4, C2H6, and C3H8, not otherwise than thus : /H c H C\H XH c h3 I c^h2 (Lh3 c=h3 c h3 In the next compound, butane, C4H10, we have two possibilities explaining the structure of the molecule, namely, these : c=h3 A=h2 I c=h2 I OEEH3 C=H, „ I C=H3—CH—CH=H3. Both these compounds are known, and termed normal butaue and isobutane, respectively. The next member, pentane, C5H12, shows three possibilities of constitution, thus: c-h3 I (’=H2 I c=h2 (J=H2 I c_h3 c=h3 c : H.,—C—H u I ceeh3 C—h3 I c^;h3—c—czh, I c=h3. These compounds also are known. With the higher members of the paraffins the number of possible isomeres rises rapidly according to the law of permuta- tion, so that we have of the seventh member 9, of the tenth 75, and of the thirteenth member 799, possible isomeric hydrocarbons. Methane, CH4 (Marsh-gas, Fire-damp). This hydrocarbon has been spoken of in Chapter 13, where it was stated that it is a color- less, combustible gas, which is formed by the decay of organic matter in the presence of moisture, during the formation of coal in the interior of the earth, and by the destructive distillation of various organic matters. Methane is of special interest, because it is the compound 300 CONSIDERATION OF CARBON COMPOUNDS. from which thousands of other substances are derived. It may be made by the action of inorganic substances upon one another; for instance, by passing a mixture of steam and carbon disulphide over copper heated to red heat, when the following change takes place: 6Cu + CS2 + 2H20 = 2Cu„S + 2CuO + CH4. Bearing in mind that carbon disulphide, as well as water, may be obtained by direct union of the elements, it is evident that methane may be formed indirectly, by means of the above method, from the elements carbon and hydrogen. Experiment 40. Use apparatus shown in Fig. 5, page 37, omitting the bent tube B. Mix in a mortar 20 grammes of sodium acetate with 20 grammes of potassium (or sodium) hydroxide and 30 grammes of calcium hydroxide; fill with this mixture the tube A, which should be made of glass fusing with difficulty, or of so-called “ combustion tubing;” apply heat and collect the gas over water. The decomposition takes place thus: NaC2H302 + NaOH = Na2C03 + CH4. Ignite the gas, and notice that its flame is but slightly luminous. Mix some of the gas in a wide-mouth cylinder, of not more than about 200 c.c. capacity, with an equal volume of air and ignite. Repeat this experiment with mixtures of one volume of methane with 2, 4, 6, 8, and 10 volumes of atmospheric air. Which mixture is most explosive, and why? How many volumes of oxygen and how many volumes of atmospheric air are needed for the complete com- bustion of one volume of methane ? Coal, Coal-oil, Petroleum. The name eoal-oil is applied to a mix- ture of the various liquid paraffins, containing often in solution the gaseous and solid members of the group, and also hydrocarbons belong- ing to other series. Coal-oil is produced in nature most likely by the decomposition of organic matter, possibly during the formation of coal. The various substances classed together under the name of coal con- sist principally of carbon, associated with smaller quantities of hydro- gen, oxygen, nitrogen, sulphur, and certain inorganic mineral matters which compose the ash. Coal is formed from buried vegetable matter by a process of decomposition which is partly a fermentation, partly a decay, and chiefly a slow destructive distillation, the heat for this latter process being derived from the interior of the earth, or by the decom- position itself. The principal constituent of the organic matter furnishing coal is wood (or woody fibre, cellulose), and a comparison of the composition HYDROCARBONS. 301 of this substance with the various kinds of coal gradually formed will help to illustrate the chemical change taking place : Wood .... Carbon. . 100 Hydrogen. 12.18 Oxygen. 83.07 Peat . 100 9.85 55.67 Lignite .... . 100 8.37 42.42 Bituminous coal . 100 6.12 21.23 Anthracite coal . 100 2.84 1.74 This table shows a progressive diminution in the proportions of hydrogen and oxygen during the passage from wood to anthracite. These two elements must, therefore, be eliminated in some form of combination which allows them to move, viz., as gases or liquids. The gases formed are chiefly carbon dioxide (which finds its way through the rocks and soils to the surface either in the gaseous state or after hav- ing been absorbed by water in the form of carbonic acid springs) and methane, known to coal-miners as fire-damp, frequently causing the formation of explosive gas mixtures in the coal mines, or escaping, like carbon dioxide, through fissures to the surface of the earth, where it may be ignited. While methane and other combustible gases are undoubtedly formed during the formation of coal, the gas mixture now generally termed natural gas (a mixture of methane, ethane, propane, hydrogen, and a few other gases), and used largely for heating and illuminating purposes, is most likely a product of the complete decomposition of vegetable and animal matter which has been precipitated from water, simultaneously with inorganic matter, during the formation of certain rocks, chiefly slate and limestone. The decomposition of this organic matter has been so complete that the gaseous decomposition-pro- ducts only are left, but no solid residue similar to coal. Petroleum, as has been stated, is chiefly a mixture of various hydro- carbons, the boiling-points of which lie between 0° and 300° C. (32° and 572° F.), or even higher. The crude oil is purified, by treating it with sulphuric acid, followed by other processes of refining, and finally by fractional distillation, in order to separate the members of low boiling-points from those of higher boiling-points. The hydrocarbons of low boiling points, chiefly a mixture of C5H12 and C6II14, are officinal, under the name of petroleum-ether or benzin, which name must not be confounded with benzene or benzol, C6H6. According to the U. S. P., benzin should have a specific gravity from 0.67 to 0.675, and a boiling-point of 50° to 60° C. (122° to 140° F.). Other similar liquids are sold in the market under the name of rhigoline, B. P. about 21° C. (70° F.) and gasoline, B. P. about 75° C. (167° F.); they are highly inflammable. 302 CONSIDERATION OF CARBON COMPOUNDS. The paraffins distilling between 150° and 250° C. (302° aud 408° F.) constitute the common illuminating oil, various kinds of which are sold as kerosene, paraffin oil, astral oil, mineral sperm oil, etc. The danger which arises in the use of coal-oil as an illuminating agent is caused by the use of oils which have not been sufficiently freed from the more volatile members of the series, which, when but slightly heated (or even at ordinary temperature), will vaporize, and, upon mixing with atmospheric air, form explosive mixtures. An oil to be safely used for illuminating purposes in common lamps should not give off inflammable vapors (or flash) below 49° C. (120° F.). Experiment 41. Various forms of apparatus are used for the exact determi- nation of the flashing-point; students may determine it approximately by operating as follows: Fill a cylinder (about one inch in diameter and six inches high) two-thirds with kerosene, suspend in the oil a thermometer, place the cylinder in a vessel with water (water-bath), keeping the level of the oil even with that of the water, and heat the latter slowly. Cover the cylinder loosely with a piece of pasteboard, and when the thermometer indicates a rise in temperature pass a small flame quickly over the mouth of the cylinder after having removed the pasteboard. Repeat this operation, from degree to degree, until a bluish flame is noticed running down to the surface of the oil. The temperature at which this takes place indicates the flashing-point. After the illuminating oil has been distilled off, a mixture of sub- stances passes over, which is used for lubricating purposes or furnishes, after having been purified by treatment with bone-black, the officinal article known ns petrolatum, petroleum, ointment, or vaseline. A mixture of the highest and solid members of the paraffin series distilling at a temperature about 350° C. (662° F.) is known as paraffin, a white, crystalline substance used for candles, etc.; it fuses at about 75° C. (167° F.). Illuminating gas is a mixture of gases obtained by the destructive' distillation of coal (or wood) in iron retorts, with subsequent purifica- tion of the gases generated. The constituents of coal have been men- tioned above. The products formed from it during its destructive distillation are very numerous; the following are the most important; Hydrogen .... H. Methane .... CH4. Ethene .... C2H4. Acetylene .... C2H2. Nitrogen . . . . N. Ammonia .... NH3. Carbonic acid . . . CO. Carbon dioxide . . . C02. Hydrosulphuric acid . . H2S. Hydrocyanic acid . . HCN. Gases HYDROCARBONS. 303 Benzene .... C6H6 80° Toluene .... C7H8 110 Aniline .... C6H5NH2 132 Acetic acid .... C2H402 117 Water H20 100 B. P. Liquids Coal-tar Carbolic acid . . . C6I160 188 Kresylic acid . . . C7H80 201 Naphthalene . . . C10H8 220 Anthracene.... C14H10 860 Paraffin .... C1BH,4 280 Solids Solid residue: Coke, chiefly carbon and inorganic matter The gases are purified by condensing ammonia (and some other gases) in water, carbon dioxide and hydrosulphurie acid in calcium hydroxide. The following is the composition of a purified illuminat- ing gas obtained from cannel-coal: Hydrogen ...... 46 volumes. Methane ...... 41 “ Ethene ...... 6 “ Carbonic oxide ..... 4 “ Carbon dioxide . . . . .2 “ Nitrogen ...... 1 volume. Experiment 42. Use apparatus shown in Fig. 5, page 37. Fill the combus- tion-tube A with sawdust (almost any other non-volatile organic matter may be used), apply heat and continue it as long as gases are evolved. Notice that by this process of destructive distillation are formed a gas (or gas mixture), which may be ignited, a dark, almost black liquid (tar), which condenses in the tube B, and that a residue is left which is chiefly carbon. The tarry liquid shows an acid reaction, due to acetic and other acids present. Coal-tar, obtained as a by-product in the manufacture of illuminat- ing gas, contains, as shown by the above table, many valuable sub- stances, such as benzene, aniline, carbolic acid, paraffin, etc., which are separated from each other by making use of the difference in their boiling-points and specific gravities, or of their solubility or insolu- bility in various liquids, or, finally, of their basic, acid, or neutral properties. Olefines. The hydrocarbons of the general formula CnH2*1 are termed olefines. To this series belong : Ethene or ethylene ..... C2H4. Propene or propylene .... C3Hg. Butene or butylene C4H8. Pentene or amylene ..... C5H10. Hexene or hexylene C6H12. 304 CONSIDERATION OF CARBON COMPOUNDS. Methene, CH2, the lowest term of this series, is not known. The hydrocarbons of this series are not only homologous, but also poly- meric with one another. Of special interest is the first known member of the series, ethene or olefiant gas, on account of its normal occurrence in illuminating gas, as well as in most common flames, the luminosity of which depends greatly on the quantity of this compound present in the burning gas. Benzene series or aromatic hydrocarbons. The members of a series of hydrocarbons having the general composition C„H2n_6, and all the derivatives of this group, including the alcohols, acids, etc., are the substances spoken of before as aromatic compounds, and will be con- sidered later. Volatile or essential oils. The term essential oil is more a pharma- ceutical than chemical term, and is used for a large number of liquids obtained from plants, and having in common the properties of being volatile, soluble in ether and alcohol, almost insoluble in water, and having a distinct and in most cases even highly characteristic odor. They stain paper as do fats or fat oils, from which they differ, how- ever, by the disappearance after some time of the stain produced, while fats leave a permanent stain. In their chemical composition essential oils differ widely; some are compound ethers, others aldehydes, but most of them are hydrocarbons or oxidized hydrocarbons, belonging to the benzene-derivatives, where they will be considered. 41. ALCOHOLS. Constitution of alcohols. The old term “alcohol” originally indi- cated but one substauce (ethyl alcohol), but is now applied to a large Questions.—391. How do hydrocarbons occur in nature, and by what pro- cesses are they formed in nature or artificially? 392. State the general physical and chemical properties of hydrocarbons. 393. What is the general composi- tion of the paraffins? 394. State the composition and properties of methane, and also the conditions under which it is formed in nature. 395. What is coal, what are its constituents, from what is it derived, and by what process has it been formed? 396. What is crude coal-oil, what is petroleum ether, and what is petrolatum? 397. How is illuminating gas manufactured, and what are its chief constituents? 398. Mention some of the important substances found in coal-tar. 399. Explain a method by which the flashing-point of coal-oil can be determined. 400. Which substances are termed volatile oils, and what are their properties ? ALCOHOLS. 305 group of substances which may be looked upon as being derived from hydrocarbons by replacement of one, two, or more hydrogen atoms by hydroxyl, OH. Any hydrocarbon may be converted into an alcohol radical by removal of one or more hydrogen atoms; methane, CH4, for instance, is converted into methyl, CH3, which, upon combining with hydroxyl, forms methyl alcohol, CH3OII. It has been shown before that the higher members of the paraffin series are capable of forming a number of isomeric compounds. Running parallel to the various series of hydrocarbons (and their isomeres) we have homologous series of alcohols. The isomeric alcohols also show properties different from one another, and yield different decomposition products. The isomeric alcohols are distinguished as normal or primary, secondary and tertiary alcohols; a normal alcohol is derived from a normal paraffin, and contains hydroxyl in the place of a hydrogen atom in a methyl group, the constitution of normal ethyl alcohol CH2.OH being, for instance, represented by the formula | CH3. If hydroxyl replaces but one atom of hydrogen in a hydrocarbon, the alcohol is termed monatomic; diatomic and triatomic alcohols are formed by replacement of two or three hydrogen atoms respectively. (Diatomic alcohols are also termed glycols.) As an instance of a diatomic alcohol may be mentioned ethylene alcohol, C2H4(OH)2, while glycerin, C3H5(OH)3, is a triatomic alcohol. Alcohols correspond in their composition to the hydroxides of inor- ganic substances ; both classes of compounds containing hydroxyl, OH, which, in the case of alcohols, is in combination with residues contain- ing carbon and hydrogen, in the case of inorganic hydroxides with metals, as, for instance, in potassium hydroxide, KOH. If we represent any unsaturated hydrocarbon by Al.R. (alcohol radical), the general formula of the alcohols will be : Monatomic alcohol. Diatomic alcohol. Triatomic alcohol. Al.Ri—OH A1 Rii Ai.Ku \QH /OH Al.Riii—OH \OH or A1 RH)H Al.Rii(OH)2 Al.Riii(OH)3 corresponding to KOH Caii(OH)2 Bi»i(OH)3. Occurrence in nature. Alcohols are not found in nature in a free or uncombined state, but generally in combination with acids as com- pound ethers. Some plants, for instance, contain compound ethers 306 CONSIDERATION OF CARBON COMPOUNDS. mixed with volatile oils. The triatomic alcohol glycerin is a normal constituent of all fats or fatty oils, and is therefore found in many plants and in most animals. Formation of alcohols. Alcohols are often produced by fermenta- tion (ethyl alcohol from sugar), sometimes by destructive distillation (methyl alcohol from wood): they are obtained from compound ethers (which are compounds of acids and alcohols) by treating them with the alkali hydroxides, when the acid enters into combination with the alkali, whilst the alcohols are liberated according to the general formula: £c!:> + K0H = id)0 + AIK0H- Alcohols may be obtained artificially by various processes, as, for instance, by treating hydrocarbons with chlorine, when the chloride of a hydrocarbon residue is formed, which may be decomposed by alkali hydroxides in order to replace the chlorine by hydroxyl, when an alcohol is formed. For instance: C2H6 + 2C1 = C2H5C1 + HC1. Ethane. C2H5C1 + KOH = KC1 + C2H5OH. Ethyl chloride. Ethyl chloride. Potassium hydroxide. Potassium chloride. Ethyl alcohol. Properties of alcohols. Alcohols are generally colorless, neutral liquids; some of the higher members are solids, none is gaseous at the ordinary temperature. Most alcohols are specifically lighter than water; the lower members are soluble in or mix with water in all proportions ; the higher members are less soluble, and, finally, insolu- ble. Most alcohols are volatile without decomposition ; some of the highest members, however, decompose before being volatilized Although alcohols are neutral substances, it is possible to replace the hydrogen of the hydroxyl by metals, as, for instance, CH3OH = methyl alcohol; CH3ONa = sodium methyl oxide or sodium me- thylate. The oxygen of alcohols may be replaced by sulphur, when com- pounds are formed known as hydrosulphides or mercaptans ; these bodies may be obtained by treating the chlorides of hydrocarbon residues with potassium sulphydrate : By replacement of the hydrogen of the hydroxyl in alcohols by alcohol radicals ethers are formed ; by replacing the same hydrogen with acid radicals compound ethers are produced. C2H5C1 + KSH = KC1 + c2h5sh. ALCOHOLS. 307 Monatomic normal alcohols of the general composition CnH2n +1OH or C H2n+20. B. P. Methyl alcohol .... . C HjOH 67° C. Ethyl “ .... . C2H5OH 78 Propyl “ .... . c3 h7 OH 97 Butyl “ .... . c4h9oh 115 Amyl “ .... . C5HuOH 132 Hexyl “ .... . C6II13OH 150 Heptyl “ .... . C7H15OH 168 Octyl “ . C8 H17OH 186 Nonyl “ .... . C9H19OH 204 Cetyl “ .... • C16H33OH 1 Fusing- Ceryl “ • c27h55oh 79 l [ point. Melissyl “ .... • CjqHjjOH 85-) 1 Methyl alcohol, CH3OH (.Methyl hydroxide, Methyl alcohol, Wood- spirit, Wood-naphtha). Methyl alcohol is one of the many products obtained by the destructive distillation of wood. When pure it is a thin, colorless liquid, similar in smell and taste to ethyl alcohol; crude wood-spirit, which contains many impurities, has an offensive odor and a nauseous, burning taste. Methyl alcohol mixes in all propor- tions with water; it dissolves resins and volatile oils as freely as ethyl alcohol, and is often substituted for the latter for various purposes in the arts and manufactures. Ethyl alcohol, C2H5OH = 46 (Common alcohol, Ethyl hydroxide, Ethylic alcohol), may be obtained from ethene, C2H4, by addition of the elements of water, which may be accomplished by agitating ethene with strong sulphuric acid, when direct combination takes place and ethyl sulphuric acid is formed : C2H4 + H2S04 = C2H5HS04. Ethene. Sulphuric acid. Ethyl sulphuric acid. Ethyl sulphuric acid mixed with water and distilled yields sul- phuric acid and ethyl alcohol: Ethyl alcohol may also be obtained, as already mentioned, by treat- ing ethyl chloride with potassium hydroxide : c2h5hso4 + h2o = h2so4 + c2h5oh. C2H5C1 + KOH = KOI + C2H5OH. While the above methods for obtaining alcohol are of scientific interest, there is but one mode of manufacturing it on a large scale, namely, by the fermentation of certain kinds of sugar, especially grape-sugar or glucose, C6H1206. A diluted solution of grape-sugar 308 CONSIDERATION OF CARBON COMPOUNDS. under the influence of certain ferments (yeast) suffers decomposition, yielding carbon dioxide and alcohol: C6H1206 = 2C02 + 2C2H5OH. Glucose. Carbon dioxide. Ethyl alcohol. Experiment 43. To a solution of 25 grammes of commercial glucose (grape- sugar) in 1000 c.c. of water add a little brewer’s yeast and introduce this mix- ture into a flask. Attach to the flask, by means of a perforated cork, a bent glass tube leading into clear lime-water, contained in a small flask. After standing (a warm place should be selected in winter for this operation) a few hours fermentation will commence, which can be noticed by the evolution of carbon dioxide, which, in passing through the lime-water, causes the precipi- tation of calcium carbonate. Fig. 40. Liebig’s condenser with distilling-flask. After fermentation ceases connect the flask with a condenser and distil over 50 to 100 c.c. of the liquid. Verify in the distilled portion the presence of alcohol by applying the tests mentioned below. For condensation of the dis- tilling vapors a Liebig’s condenser, represented in Fig. 40, may be used. This apparatus consists of a wide glass tube through which passes the narrow con- densing tube, connected with the boiling-flask a. A constant current of cold water is obtained by allowing water to flow into b, and to escape by c. A small flask is placed under d for collecting the distillate. By distilling the fermented liquid an alcohol is obtained containing large quantities of water; on distilling this dilute alcohol a second and a third time, collecting the first portions of the distilled liquid separately, an alcohol is obtained containing but little water. These ALCOHOLS. 309 last quantities of water, amounting to about 14 per cent., cannot be removed by simple distillation, but may be separated by mixing the alcohol with half its weight of calcium oxide, which combines with the water to form calcium hydroxide, from which the alcohol may now be separated by distillation. The alcohol thus obtained is known as, pure, absolute, or real alcohol. The alcohol of the U. S. P. contains 91 per cent, by weight, or 94 per cent, by volume of real alcohol, and has a specific gravity of 0 820 at 15.6° C. (60° F.) The diluted alcohol is made by mixing equal weights of water and alcohol, and has a specific gravity of 0.928. What is generally known as spirit of wine'or rectified spirit is an alcohol containing 84 per cent., and proof-spirit one containing 49 per cent, by weight of pure alcohol. Pure alcohol is a transparent, colorless, mobile, and volatile liquid, of a characteristic pungent and agreeable odor,1 and a burning taste; it boils at 78.3° C. (173° F.), has a specific gravity of 0.794, is of a neutral reaction, becomes syrupy at —110° C. (—166° F.), and solidi- fies at —130° C. (—202° F.); it burns with a non-luminous flame ; when mixed with water a contraction of volume occurs, and heat is liberated; the attraction of alcohol for water is so great that strong alcohol absorbs moisture from the air or abstracts it from membranes, tissues, and other similar substances immersed in it; to this property are due its coagulating action on albumin and its preservative action on animal substances. The solvent powers of alcohol are very exten- sive, both for inorganic and organic substances; of the latter it readily dissolves essential oils, resins, alkaloids, and many other bodies, for which reason it is used in the manufacture of the numerous officinal tinctures, extracts, and fluid extracts. Alcohol taken internally in a dilute form has intoxicating proper- ties ; pure alcohol acts poisonously; it lowers the temperature of the body from 0.5° to 2° C. (0.9° to 3.6° F.), although the sensation of warmth is experienced. Analytical reactions for ethyl alcohol. 1. Dissolve a small crystal of iodine in about 2 c.c. of alcohol; add to the cold solution potassium hydroxide until the brown color of the solution disappears; a yellow precipitate of iodoform, CHI3, forms. Many other alcohols, aldehyde, acetone, etc., show the same reaction. 1 It is stated that perfectly anhydrous alcohol has no odor. 310 CONSIDERATION OF CARBON COMPOUNDS. 2. Add to about 1 c.c. of alcohol the same volume of sulphuric acid ; heat to boiling and add gradually a little more alcohol: the odor of ethyl ether will be noticed distinctly on further heating. 3. Add to a mixture of equal volumes of alcohol and sulphuric acid, a crystal (or strong solution) of sodium acetate: acetic ether is formed and recognized by its odor. 4. To about 2 c.c. of potassium dichromate solution add 0.5 c.c. of sulphuric acid and 1 c.c. of alcohol: upon heating gently the liquid becomes green from the formation of chromic sulphate, while aldehyde is formed and may be recognized by its odor. Alcoholic liquors. Numerous substances containing sugar or starch (which may be converted into sugar) are used in the manufacture of the various alco- holic liquors, all of which contain more or less of ethyl alcohol, besides color- ing matter, ethers, compound ethers, and many other substances. White and red wines are obtained by the fermentation of the grape-juice; the so-called light wines contain from 10 to 12, the strong wines, such as port and sherry, from 19 to 25 per cent, of alcohol; if the grapes contain much sugar, only a portion of it is converted into alcohol, whilst another portion is left undecomposed; such wines are known as sweet wines. Effervescent wines, as champagne, are bottled before the fermentation is complete; the carbonic acid is disengaged under pressure and retained in solution in the liquid. Beer is prepared by fermentation of germinated grain (generally barley) to which much water and some hops have been added; the active principle of hops is lupulin, which confers on the beer a pleasant, bitter flavor, and the property of keeping without injury. Light beers have from 2 to 4, strong beers, as porter or stout, from 4 to 6 per cent, of alcohol. Spirits differ from either wines or beers in so far as the latter are not dis- tilled, and therefore contain also non-volatile organic and inorganic substances, such as salts, etc., not found in the spirits, which are distilled liquids contain- ing volatile compounds only. Moreover, the quantity of alcohol in spirits is very much larger, and varies from 45 to 55 per cent. Of distilled spirits may be mentioned: American whiskey, made from fermented rye or Indian corn; Irish whiskey, from potatoes; Scotch whiskey, from barley; brandy or cognac, by distilling French wines; rum, by fermenting and distilling molasses; arrack, from fermented rice; gin, from various grains flavored with juniper berries. Amyl alcohol, C5Hu0H. This alcohol is frequently formed in small quantities during the fermentation of corn, potatoes, and other substances. When the alcoholic liquors are distilled, amyl alcohol passes over toward the end of the distillation, generally accompanied by propyl, butyl, aud other alcohols, and by certain ethers and com- pound ethers. A mixture of these substances is known as fusel oil, and, from this liquid, amyl alcohol may be obtained in a pure state. It is an oily, colorless liquid, having a peculiar odor, and a burning, ALCOHOLS. 311 acrid taste; it is soluble in alcohol, but not in water. By oxidation of amyl alcohol, valerianic acid is obtained. Amylene hydrate, Ethyl-dimethyl-carhinol, ChHvlO, is an alcohol isomeric with the above amyl alcohol, but yielding only acetic acid on oxidation. It is a colorless liquid, having a pungent, ethereal odor, and a boiling-point of 100° C. (212° F.). Glycerin, Glycerinum, C3H5(0H)3 = 92. Glycerin is the triatomic or tri-acid alcohol of the residue glyceryl, C3H5, formed by removal of the three atoms of hydrogen from the saturated hydrocarbon pro- pane, C3H8, and by combination of the residue with 30H. Glycerin is a normal constituent of all fats, which are glycerin in which the three atoms of hydrogen of the hydroxyl have been replaced by residues of fat acids. When fats are treated with alkalies, these latter combine with the fat acids, whilst glycerin is liberated. Upon this decomposition, carried out on a large scale in the manufacture of soap, depends the mode of obtaining glycerin. Pure glycerin is a clear, colorless, odorless liquid of a syrupy con- sistence, oily to the touch, hygroscopic, very sweet, and neutral in re- action, soluble in water and alcohol in all proportions, but insoluble in ether, chloroform, benzol, and fixed oils; its specific gravity is 1.255 ; it cannot be distilled by itself without decomposition, but is volatil- ized in the presence of water, or when hot steam is allowed to pass through it. Glycerin is a good solvent for a large number of organic and inor- ganic substances; the solutions thereby obtained are often termed glycerites; frequently used are the glycerites of starch, carbolic acid, tannic acid, sodium biborate, etc. Analytical reactions. 1. A borax bead immersed for a few minutes in a solution of glycerin (made slightly alkaline with potassium hydroxide) imparts a green color to a non-luminous flame, owing to the liberation of boric acid. 2. Glycerin slightly warmed with an equal volume of sulphuric acid should not turn dark, but, on further heating, the characteristic, irritating odor of acrolein is noticed. 3. Fehling’s solution (see index) should not cause a red precipita- tion on heating, indicating the absence of glucose and dextrin. 312 CONSIDERATION OF CARBON COMPOUNDS. Nitro-glycerin, C3H5(N020)3 (Glyceryl tri-nitrate). When glycerin is treated with nitric acid, or, better, with a mixture of concentrated sulphuric and nitric acids, the radical N02 replaces hydrogen in the glycerin, forming either mono- or tri-nitro-glycerin, substances which belong to the compound ethers, the constitution of which will be explained later. The tri-nitro-glycerin is the common nitro-glycerin, a pale-yellow oily liquid, which is nearly insoluble in water, soluble in alcohol, crystallizes at —20° C. (—4° F.) in long needles, and explodes very violently by concussion; it may be burned in an open vessel, but explodes when heated over 250° C. (482° F.).' Dynamite is infusorial earth impregnated with nitro-glycerin. A % Phenols. The substances termed phenols are formed by replace- ment of hydrogen by hydroxyl in the aromatic hydrocarbons of the benzene series; they have the constitution of alcohols but are not alcohols in the sense in which this term is used. The more important substances belonging to this group will be considered later. 42. ALDEHYDES. HALOID DERIVATIVES. Aldehydes. The name aldehyde is derived from alcohol dehydro- genation, referring to its method of formation, viz., by the removal of hydrogen from alcohols, as, for instance : C2ir60 — 2H = c2h4o. This removal of hydrogen may be accomplished by various methods,, as, for instance, by oxidation of alcohols, when one atom of oxygen combines with two atoms of hydrogen, forming water, whilst an alde- hyde is formed at the same time. Aldehydes, when further oxidized, Ethyl alcohol. Acetic aldehyde. Questions.—401. What is the general constitution of alcohols, and what is the difference between monatomic, diatomic, and triatomic alcohols ? 402. How do alcohols occur in nature? 403. By what processes may alcohols be formed artificially, and how may they be separated from their combinations ? 404. State the general properties of alcohols. 405. Mention names and composition of the first five members of alcohols of the general composition CnH2n+iOH. 406. By what process is methyl alcohol obtained, under what other names is it known, and what are its properties? 407. Describe the manufacture of pure alcohol from sugar. 408. Give the alcoholic strength of the alcohol and diluted alcohol of the U. S. P., and also of spirit of wine, proof spirit, light wines, heavy wines, beers, and spirits. 409. What are the general properties of com- mon alcohol? 410. What is glycerin, how is it found in nature, how is it obtained, and what are its properties ? ALDEHYDES. HALOID DERIVATIVES. are converted into acids ; aldehydes are, consequently, the intermediate products between alcohols and acids, and are frequently looked upon as the hydrides of the acid radicals. The constitution of acetic acid may be represented by the formula C Hg.CO.OH ; the radical of acetic acid or acetyl is the group CH3.CO, and the hydride of acetyl is acetic aldehyde, CH3.COH. It is the group COH which is characteristic of, and found in, all aldehydes. Only a few aldehydes are of practical interest, as, for instance, acetic aldehyde, paraldehyde, and benzoic aldehyde, which latter substance will be more fully considered in con- nection with the aromatic substances. Acetic aldehyde, C2H40 or CH3C0H. Alcohol may be converted into aldehyde by the action of various oxidizing agents ; the one generally used is potassium dichromate, which oxidizes two hydrogen atoms of the alcohol molecule, converting it into aldehyde : c2h6o + 0 = c2h4o + h2o. Experiment 44. Place in a 500 c.c. flask, provided with a funnel-tube and connected with a Liebig’s condenser, 6 grammes of potassium dichromate. Pour upon this salt through the funnel-tube, very slowly, a previously pre- pared and cooled mixture of 5 c.c. of sulphuric acid, 24 c.c. of water and 6 c.c. of alcohol. Chemical action begins generally without application of heat, and often becomes so violent that the liquid boils up, for which reason a largsS; flask is used. The escaping vapors, which are a mixture of aldehyde, aldbholy' and water, are collected in a receiver kept cold by ice. • From mixfcire? pure aldehyde may be obtained by repeated distillation. Use the distillate for silvering a test-tube by adding some ammoniated silver nitrate. How much potassium dichromate is needed for the conversion of 5 grammes of«pure alcohol into aldehyde ? Aldehyde is a neutral, colorless liquid, having a strong aud charac- teristic odor; it mixes with water and alcohol in all proportions and boils at 21° C. (69.8° F.). The most characteristic chemical property" of aldehyde is its tendency to combine directly with a great number of substances; thus it combines with hydrogen to form alcohol, with oxygen to form acetic acid, with ammonia to form aldehyde-ammonia, C2H4O.NH3, a beautifully crystallizing substance, with hydrocyanic acid to form aldehyde hydrocyanide, C2H4O.HCN, and with many other substances. In the absence of such other substance it unites often with itself, forming polymeric modifications, such as paraldehyde and metaldehyde. Aldehyde is a strong reducing agent, which property is used in the silvering of glass, which is done by adding aldehyde to an ammoniacal 314 CONSIDERATION OF CARBON COMPOUNDS. solution of silver nitrate, when metallic silver is deposited on the walls of the vessel or upon substances immersed in the solution. Paraldehyde, C6H1203. When a few drops of concentrated sulphuric acid are added to aldehyde, this becomes hot and solidifies on cooling to 0° C. (32° F.). This solid crystalline mass of paraldehyde, which liquefies at 10.5° C. (51° F.), has been formed by the direct union of three molecules of common aldehyde. Paraldehyde is soluble in 8 parts of water, boils at 124° C. (253° F.), and is reconverted into common aldehyde by boiling it with dilute sulphuric or hydrochloric acid. Metaldehyde, (C2H40)a;, is another polymeric modification of aldehyde, ob- tained by a process similar to the one mentioned for paraldehyde, but at a lower temperature. It is a solid crystalline substance, insoluble in water, but slightly soluble in alcohol, ether, and chloroform. Trichloraldehyde, Chloral, C2HC130 or CC13.C0H (Trichloracetyl hydride). This substance may be looked upon as acetic aldehyde, C2H40, in which three atoms of hydrogen have been replaced by chlorine. It is made by passing a rapid stream of dry chlorine into pure alcohol to saturation, keepiug the alcohol cool during the first few hours, and warming it gradually until the boiling-point is reached. According to the quantity of alcohol operated on, the con- version requires several hours or even days. The crude liquid pro- duct separates into two layers ; the lower is removed and shaken with three times its volume of strong sulphuric acid and distilled, the dis- tillate is mixed with calcium oxide and again distilled; the portion passing over between 94° and 99° C. (201° and 210° F.) is collected. The decomposition taking place between alcohol and chlorine may be explained by the formation of aldehyde : and by the subsequent replacement of hydrogen by chlorine: C2H60 + 2CI = C2H40 + 2HC1, C2H40 + 6C1 = C2HC130 + 3HC1. The actual decomposition is, however, somewhat more complicated, numerous other products being formed at the same time. By treat- ment with sulphuric acid these other substances are removed. Chloral is a colorless, oily liquid, having a penetrating odor and an acrid, caustic taste ; its specific gravity is 1.5, and its B. P. 95° C. (203° F.). ALDEHYDES. HALOID DERIVATIVES. 315 Chloral hydrate, Chloral, U. S. P., C2HC1302.H0 = 165.2. When water is added to chloral the two substances combine, heat is dis- engaged, and the hydrate of chloral is formed, which is a crystalline, colorless substance, having an aromatic, penetrating odor, a bitter, caustic taste, and a neutral reaction; it is freely soluble in water, alcohol, and ether, also soluble in chloroform, carbon disulphide, benzene, fatty and essential oils, etc.; it liquefies when mixed with carbolic acid or with camphor; it melts at 58° C. (136° F.), and boils at 95° C. (203° F.), and also volatilizes slowly at ordinary temperature. Chloral, and its hydrate, are decomposed by weak alkalies into chloroform and a formate of the alkali metal : C2HC130 + KHO = KCH02 + OHCI3. Chloral. Potassium hydroxide. Potassium formate. Chloroform. This decomposition was believed to take place in the animal body, and especially in the blood, whenever chloral was given internally, but recent in- vestigations seem to contradict this assumption. There is no chemical antidote which may be used in cases of poisoning by chloral, and the treatment is, therefore, confined to the use of the stomach-pump and to the maintenance of respiration. Analytical reactions for chloral. 1. Chloral or chloral hydrate heated with potassium hydroxide is converted into potassium formate and chloroform, which latter may be recognized by its odor. (See explanation above.) 2. Heated with silver nitrate and ammonium hydroxide a silver- mirror is formed on the glass. 3. Heated with Fehling’s solution a red precipitate is formed. See also reactions 2 and 6 for chloroform below. Chloroform, Chloroformum, CHC13 = 119.2 (Trichlormethane, Di- chlormethyl chloride). When either chlorine, bromine, or iodine is allowed to act upon methane, CH4, a number of substitution products are formed. Thus, if methane is considered as methyl hydride, CH3H, the first product of substitution is methyl chloride, CH3C1; the second is monochlormethyl chloride, CH2C1CI; the third is dichlormethyl chloride or chloroform, CHC12C1; and the fourth is carbon tetrachloride, CC14. Similar products are formed by the action of iodine or bromine upon methane, or, in fact, upon any of the paraffins. Chloroform is, however, not obtained for commerce by the above process, but by the action of bleaching-powder and calcium hydroxide 316 CONSIDERATION OF CARBON COMPOUNDS. on alcohol. The three substances named, after being mixed with a considerable quantity of water, are heated in a retort until distillation commences; the crude product of distillation is an impure chloroform, which is purified by mixing it with sulphuric acid and allowing the mixture to stand ; the upper layer of chloroform is removed and treated with sodium carbonate (to remove any acids) and distilled over calcium oxide (to remove water). The explanation of the formation of chloroform by the above pro- cess has indirectly been given in connection with the consideration of chloral, where it has been shown that alcohol is converted by the action of chlorine first into aldehyde and subsequently into chloral, which, upon being treated with alkalies, is decomposed into an alkali formate and chloroform. The action of the chlorine of the calcium hypochlorite (which is the active principle in bleaching-powder) upon the alcohol is similar to that of free chlorine upon alcohol; in both cases aldehyde, and afterward chloral, are formed, which latter, in the manufacture of chloroform, is decomposed by the calcium hydroxide into calcium formate and chloroform. If the various intermediate steps of the decomposition are not con- sidered, the process may be represented by the following equation : Alcohol. 4C2H60 + 8Ca(CIO)2 = 2CHC13 + 3[Ca(CH02)2] + 5CaC!2 + 8H,0. Calcium hypochlorite. Chloroform. Calcium formate. Calcium chloride. Water. Chloroform is now made extensively by the action of bleaching- powder upon acetone; the reaction takes place thus : Acetone. 2CO(CH3)2 + 3Ca(C10)2 = 2CHCI3 + 2Ca(OH)2 + Ca(C2H302)2. Calcium hypochlorite. Chloroform. Calcium hydroxide. Calcium acetate. Pure chloroform is a heavy, colorless liquid, of a characteristic ethereal odor, a burning, sweet taste, and a neutral reaction ; it is but very sparingly soluble in water, but miscible with alcohol and ether in all proportions; the specific gravity of pure choloroform is 1.50, but a small quantity of alcohol (from one-half to one per cent.), allowed to be present by the U. S. P., causes the specific gravity to be about 1.488 ; boiling-point 62° C. (143° F.), but rapid evaporation takes place at all temperatures. Chloroform should be tested for excess of alcohol by specific gravity; for hydrochloric acid and chlorine by shaking it with water, which afterward should not give a precipitate with silver nitrate; for aldehyde by heating with solution of potassium hydroxide, which should not be colored brown; for empyreumatic and other organic compounds by shaking with an equal volume ALDEHYDES. HALOID DERIVATIVES. of pure sulphuric acid, which should remain colorless; or by evaporation when no residue should be left and no odor should be perceptible after the chloroform has been volatilized. Analytical reactions for chloroform. 1. Dip a strip of paper into chloroform and ignite. The flame has a green mantle and emits vapors of hydrochloric acid, rendered more visible upon the approach of a glass rod moistened with water of ammonia. 2. Add a drop of chloroform and a drop of aniline to some alcoholic solution of potassium hydroxide and heat gently: a peculiar, pene- trating, offensive odor of benzo-isonitril, C6II5NC, is noticed. (Chloral shows the same reaction.) CHC13 + 3K0H + C6H5.NH2 = C6H5NC + 3KC1 + 3H20. 3. Add some chloroform to Fehling’s solution and heat: red cuprous oxide is precipitated. 4. Vapors of chloroform, when passed through a glass tube heated to redness, are decomposed iuto carbon, chlorine, and hydrochloric acid. The two latter should be passed into water, and may be recog- nized by their action on silver nitrate (white precipitate of silver nitrate) and on mucilage of starch, to which potassium iodide has been added (blue iodized starch is formed). 5. Heat some chloroform with solution of potassium hydroxide and a little alcohol. Chloroform is decomposed into potassium chloride and formate : CHC13 + 4K0H = 3KC1 + KCH02 + 2H.fi. Divide solution into two portions. Acidulate one portion with nitric acid, boil, aud add silver nitrate: white precipitate of silver nitrate. To second portion add a little water of ammonia and a crystal of silver nitrate: a mirror of metallic silver will be formed after heating slightly. 6. Add to 1 c.c. of chloroform about 0.3 gramme of resorcin in solution, and 3 drops of solution of sodium hydroxide; boil strongly : a yellowish-red color is produced, and the liquid shows a beautiful yellow-green fluorescence. (Chloral shows the same reaction.) In cases of poisoning, chloroforni is generally to be sought for in the lungs and blood, which are placed in a flask connected with a tube of difficultly fusible glass. By heating the flask the chloroform is expelled and decomposed in the heated glass tube, as stated above in reaction 4. Another portion of 318 CONSIDERATION OF CARBON COMPOUNDS. chloroform should be distilled without decomposing it, and the distillate tested as above stated. What has been said above regarding antidotes to chloral holds good for chloroform also. Bromoform, CHBr3 (Dibromomethyl bromide). Obtained by gradu- ally adding bromine to a cold solution of potassium hydroxide in methyl alcohol until the color is no longer discharged, and rectifying over calcium chloride. Bromoform is a colorless liquid which has an aromatic odor and a sweet taste. Sp. gr. 2.13 ; B. P. 150° C. (302°F.); solidifies at —9° C. (15 8° F.) It is sparingly , soluble in water, soluble in alcohol and ether. Its physiological action is similar to that of chloroform. Iodoform, Iodoformum, CHI3 = 292.6 (Diiodomethyl iodide). This compound is analogous in its constitution to chloroform and bromo- form. It is made by heating together an aqueous solution of an alkali carbonate, iodine, and alcohol until the brown color of iodine has dis- appeared ; on cooling, iodoform is deposited in yellow scales, which are well washed with water and dried between filtering paper. Iodoform occurs in small, lemon-yellow, lustrous crystals, having a peculiar, penetrating odor, and an unpleasant, sweetish taste; it is nearly insoluble in water and acids, soluble in alcohol, ether, fatty and essen- tial oils. It contains 96.7 per cent, of iodine. Iodoform digested with an alcoholic solution of potassium hydroxide imparts, after acidulation with nitric acid, a blue color to starch solu- tion. Experiment 45. Dissolve 4 grammes of crystallized sodium carbonate in 6 c.c. of water; add to this solution 1 c.c. of alcohol; heat to about 70° C. (158° F.), and add gradually 1 gramme of iodine. A yellow crystalline deposit of iodoform separates. Ethyl bromide, C2H5Br (Hydrobromio ether). Obtained by the simultaneous action of phosphorus and bromine on ethyl alcohol. It is a colorless, ethereal liquid, which boils at 40° C. (104° F.), and has a sp. gr. of 1.473. Sulphonal, (CH3)2C(C2H5S02)2, Dimethyl-diethylsulphonyl-methane. It has been stated before that mercaptans are alcohols in which the oxygen is replaced by sulphur. Alcohol treated with oxidizing agents are converted into acids by exchanging two atoms of hydrogen for one atom of oxygen. Mercaptans behave differently; they combine MONOBASIC FATTY ACIDS. 319 directly with three atoms of oxygen, forming compounds known as sulphonic acids. Thus, ethyl mercaptan, C2H5HS, when treated with nitric acid, is converted into ethyl-sulphonic acid, C2H5HS03. The radical of this acid, known as ethylsulphonyl, C2HsS02, may, by indi- rect process, be caused to replace hydrogen in methane, CH4, twice, while the two remaining methane hydrogen atoms can be replaced by methyl. The compound thus obtained is the dimethyl-diethylsul- phonyl-methane, or sulphonal. The relations between methane and some of its derivatives, which have been considered in this chapter, may be shown graphically thus : H I II—C—H I H Methane. Cl H—C—Cl I Cl Chloroform. I H—C—I I I Iodoform. COH I H—C—PI I H Aldehyde. COH Cl—C -Cl I Cl Chloral. ch3 I CH3—C—c2h5so2 I C2H5S02. Sulphonal. Sulphonal is a white crystalline substance, haviug neither odor nor taste; it is soluble in 20 parts of boiling and 100 parts of cold water, soluble with difficulty in alcohol, but easily soluble in ether, benzene, and chloroform ; it fuses at 130° C. (266° F.), and volatilizes at about 300° C. (572° F.), with partial decomposition. 43. MONOBASIC FATTY ACIDS. General constitution of organic acids. When hydroxyl, OH, re- places hydrogen in hydrocarbons, alcohols are formed; when the univalent group, C02H, known as carboxyl, replaces hydrogen in hydrocarbons, acids are formed. Monatomic, diatomic, and triatomic alcohols are formed by introducing hydroxyl once, twice, or three Questions.—411. What is an aldehyde, and what are its relations to alco- hols and acids? 412. State the composition of acetic aldehyde. 413. Explain the action of chlorine upon alcohol. 414. Give the composition and properties of chloral and chloral hydrate. 415. What decomposition takes place when alkalies act upon chloral? 416. Describe the process of preparing and purify- ing chloroform. 417. What is the composition of chloroform and what are its properties? 418. How is chloroform tested for impurities? 419. By what tests may chloroform be recognized? 420. How is iodoform made!, and what are its properties ? • 320 CONSIDERATION OF CARBON COMPOUNDS. times respectively into hydrocarbon molecules; monobasic, dibasic, and tribasic acids are formed by substituting one, two, or three hydrogen atoms by carboxyl. For instance : Hydrocarbons. Monobasic acids. Dibasic acids. ch4 ch3.co2h CH /G02H °a12\co2h- Methane. Acetic acid. Malonic acid c2h6 c2h5.co2h c jj /C02H *\C02H‘ Ethane. Propionic acid. Succinic acid. The constitution of carboxyl is represented by 0=0—O—H, which shows that of the four affinities of the carbon atom, two are saturated by an atom of oxygen, one by hydroxyl, whilst one is unprovided for; any univalent hydrocarbon residue may attach itself to this unprovided affinity, when an acid is formed. Acids may be looked upon, therefore, as being composed of hydrocarbon residues and hydroxyl, united by the bivalent radical CO, termed carbonyl. By replacement of the hydrogen of the hydroxyl (or of the carboxyl, which is the same) by metals the various salts are formed. What is termed the acid radical is the group of the total number of atoms present in the molecule, with the exception of the hydroxyl. In acetic acid, C2H402, for instance, the radical is CH3CO, or C2H30, which group of atoms, known as acetyl, is characteristic of acetic acid, and of all acetates, and may often be transferred from one compound into another without decomposition. The difference between alcohol radicals and acid radicals may also be stated, by saying that the first contain carbon and hydrogen only, while acid radicals contain carbon, hydrogen, and oxygen. In a similar manner, as there are homologous series of alcohols corresponding to the various series of hydrocarbons, there are also homologous series of organic acids running parallel with the corre- sponding series of hydrocarbons or alcohols. Occurrence in nature. Organic acids are found and formed both in vegetables and animals, and are present either in the free state, or (and more generally) in combination with bases as salts, or with alcohols as compound ethers. Uncombined or as salts are found, for instance, citric, tartaric, and oxalic acids in plants, formic acid in some insects, uric acid in urine, etc.; as compound ethers are found many of the fatty acids in the various fats. Some organic acids are also found as products of the decomposition of organic matters in nature. MONOBASIC FATTY ACIDS. 321 Formation of acids. Many acids are produced by oxidation of alco- hols. As intermediate products are formed aldehydes, which may be looked upon (as stated in the last chapter) as alcohols from which two atoms of hydrogen have been removed. For instance: c2h5oh + o = c2h3oh + h2o. Ethyl alcohol. Acetic aldehyde. C2HjOH + O = C2H3O.OH. Acids are obtained from compound ethers by boiling them with alka- lies, when salts are formed, which may be decomposed by sulphuric or other acids. For instance : Acetic aldehyde. Acetic acid. C(?h)° + K0H = + C2H5OH. Ethyl acetate. Potassium hydroxide. Potassium acetate. Ethyl alcohol, 202H3K02 + H2S04 = 2C2H402 + K2S04. Potassium acetate. Sulphuric acid. Acetic acid. Potassium sulphate. Acids are formed also by destructive distillation (acetic acid); by fermentation (lactic acid); by putrefaction (butyric acid); by oxidation of many organic substances (oxalic acid by oxidation of starch), etc. Properties. Organic acids show the characteristics mentioned of inorganic acids, viz., when soluble, have an acid or sour taste, redden litmus, and contain hydrogen replaceable by metals, with the forma- tion of salts. Most organic acids, and especially the higher members, show these acid properties in a less marked degree than inorganic acids ; in fact, they become so weak that the acid properties can often scarcely be rec- ognized. As stated above, mono-, di-, and tri-basic organic acids are known, the latter two being capable of forming normal, acid, or double salts. Most organic acids are colorless, some of the lower and volatile acids have a characteristic odor, but most of them are odorless; most organic acids are solids, Some liquids, scarcely any gaseous at the ordi- nary temperature. Any salt formed by the union of an organic acid and a non-volatile metal (especially alkali metal) leaves the car- bonate of this metal after the salt has suffered combustion. It is for this reason that ashes contain most metals in the form of carbonates. Whilst the hydrogen of the hydroxyl may be replaced by metals or by other residues, the hydrogen of the acid radical may often be re- placed by chlorine, and the oxygen of the hydroxyl by sulphur. The 322 CONSIDERATION OF CARBON COMPOUNDS. acids formed by this last reaction are known as thio acids, for instance, thio-acetic acid, C2H4OS. When the hydrogen of the hydroxyl is replaced by a second acid radical (of the same kind as the one forming the acid) the so-called anhydrides are produced, which correspond to the inorganic anhydrides. For instance: HNOs or N02.0H Nitric acid. C2H402 or C2H,O.OH. Acetic acid. N02/U Nitric anhydride. c2h30/u- Acetic anhydride. Amido-acids are compounds obtained by replacement of a hydrogen atom by NH2; these compounds will be spoken of later in connection with amides. Fatty acids of the general composition, CnH2n02 or CnH2n+1C02H. Formic acid, H C02H Fusing- Boiling- point. •point. + 4° C. 100° C. Occurs in : Red ants and some plants, etc. Acetic acid, c h3co2h + 17 118 Vegetable and animal fluids. Propionic acid, c2 h5 co2h —21 140 Sweat, fluids of the stomach, etc. Butyric acid, c3 h7 co2h —20 162 Butter. Valerianic acid, C4 H9 C02H —16 185 Valerian root. Caproic acid, c5 huco2h — 2 205 Butter. (Enanthylic acid, C6 H13C02H —10 224 Castor oil. Caprylic acid, c7 h15co2h + 14 236 Butter; cocoanut oil. Pelargonic acid, c8 H17co2H 18 254 Leaves of geranium. Oapric acid, c9 h19co2h 30 270 Butter. Laurie acid, cuh23co2ii 43. Myristic acid, c13h27co2h 54 } Cocoanut oil. Palmitic acid, c15h31co2h 62 Palm oil, butter. Margaric acid, c16h33co2h 60 (Obtained artificially.) Stearic acid, c17h35co2h 70 Most solid animal fats. Araehidic acid, c19h39co2h 75 ) Behenic acid, o21h4Sco2h 76 Oils of certain plants. Hysenic acid, c24h49co2h 77 i Oerotic acid, C26H53C02H 80 Beeswax. Melissic acid, c29h59co2h 90 } The name fatty acids has been given to these acids on account of their frequent occurrence in fats, and also in allusion to the somewhat fatty appearance of the higher members of the series. The gradual change of properties which the members of an homol- ogous series show, is well marked in the series of fatty acids, thus: MONOBASIC FATTY ACIDS. 323 First member. Last member. Is liquid. Is solid. Volatilized at 100° C. Not volatilized without decomposition. Strongly acid. Scarcely acid. Strongly odoriferous. Odorless. Easily soluble in water. Insoluble in water. Produces no grease spot. Produces a grease spot. Forms salts easily soluble without ■ Forms salts which are insoluble or de- decomposition. composed by water. The intermediate members of the series show intermediate proper- ties, and this change in properties is in proportion to the gradual change in molecular weight. Formic acid, H.C02H or CHO.OH. This acid is found in the red ant and in other insects, which eject it when irritated. It is also con- tained in some plants, as, for instance, in the leaves of the stinging- nettle. It is formed by the oxidation of methyl alcohol: CH30 + 02 = CH202 + H20, Methyl alcohol. Formic acid. by the action of carbonic oxide on potassium hydroxide KOI! + CO = KCH02, Potassium formate. by the action of potassium hydroxide on chloroform : CHC13 + 4K0H = 3KC1 + 2H20 + KCH02, by heating equal parts of glycerin and oxalic acid, when the latter is split up into carbon dioxide and-formic acid, which may be separated from the glycerin by distillation : C2H204 = C02 + CH202. It is also a product of the decomposition of sugar, starch, etc. Formic acid is a colorless liquid having a penetrating odor, and a strongly acid taste; it produces blisters on the skin; it is a powerful deoxidizer, being, when thus acting, converted into carbon dioxide and water : Oxalic acid. Formic acid. ch2o2 + o = co2 + h2o. Acetic acid, H.C2H302, or C2H30.0H, or CH3.C02H = 60. The most important alcohol is ethyl alcohol, and the most largely used organic acid is acetic acid, obtained from ethyl alcohol by oxidation. Acetic acid is found in combination with alkali metals in the juices of many plants, also in the secretions of the glands, etc. 324 CONSIDERATION OF CARBON COMPOUNDS. Acetic acid is formed chiefly either by the oxidation of alcohol (and aldehyde) or by the destructive distillation of wood. It is produced commercially on a large scale as follows : A diluted alcohol (8 to 10 per cent.) is allowed to trickle down slowly through wood-shavings contained in high casks having perforated sides in order to allow a free circulation of the air; the temperature is kept at about 24° to 30° C. (75° to 86° F.), and the liquid having passed through the shavings is repeatedly poured back in order to cause complete oxida- tion. When the latter object has been accomplished the liquid is a diluted acetic acid. It appears that the conversion of alcohol into acetic acid is greatly facilitated by the presence of a microscopic organism (mycoderma aceti) commonly termed “ mother of vinegar.” This serves in some unexplained way to convey the atmospheric oxygen to the alcohol. The term “ acetic fermentation ” is often applied to this conversion, although it is not a true fermentation, since no splitting up of the alcohol molecule into other less complex compounds, but a process of slow oxidation, takes place. The second process for manufacturing acetic acid is the heating of wood to a red heat in iron retorts, when numerous products (gases, aqueous and tarry substances) are formed. The aqueous products contain, besides other substances, methyl alcohol and acetic acid. The liquid is neutralized with calcium hydroxide and distilled, when methyl alcohol, water, etc., evaporate and a solid residue is left, which is an impure calcium acetate. From this latter, acetic acid is obtained by distilling with sulphuric (or hydrochloric) acid, calcium sulphate (or chloride) being formed and left in the retort, whilst acetic acid distils over. Experiment 46. Add to 54 grammes of sodium acetate contained in a small flask which is connected with a Liebig’s condenser, 40 grammes of sulphuric acid. Apply heat and distil over about 35 c.c. Determine volumetrically the amount of pure acetic acid in this liquid. Pure acetic acid, or glacial acetic acid, is solid at or below 16° C. (61° F.); at higher temperatures it is a colorless liquid having a char- acteristic, penetrating odor, boiling at 118° C. (241° F.), and causing blisters on the skin; its specific gravity is 1.056; it is miscible with water, alcohol, and ether, is strongly acid, forming salts known as acetates, which are all soluble in water. Vinegar is dilute acetic acid (about 6 per cent.), containing often other substances, such as coloring matter, compound ethers, etc. MONOBASIC FATTY ACIDS. 325 Vinegar was formerly obtained exclusively by the oxidation of fer- mented fruit-juices (wine, cider, etc.), the various substances present in them imparting a pleasant taste and odor to the vinegar; to-day vinegar is ofteu made artificially by adding various coloring and odoriferous substances to dilute acetic acid. Vinegar should be tested for sulphuric and hydrochloric acids, which are sometimes fraudu- lently added. Acidum aceticum, Acidum aceticum dilutum, and Acidum aceticum glaciate are the three officinal forms of acetic acid. The first-named acid contains 36 per cent., the second 6 per cent., the third at least 99 per cent, of pure acetic acid. Acetic acid shows an exceptional behavior in regard to the specific gravity of its aqueous solutions. The highest specific gravity of 1.0748 belongs to an acid of 77 per cent., which is equal to an acid containing one molecule of water and one of acetic acid, or C2H402.H20. The addition of either acetic acid or of water causes the liquid to become lighter. For instance, the specific gravity of an acid contain- ing 95 per cent, is equal to that containing 56 per cent, of pure acid, both solutions having a specific gravity of 1.066. The specific gravity of dilute acetic acid cannot, therefore, be used as a means of determining the amount of pure acid; this is done by exactly neutralizing a weighed portion of the acid with an alkali; from the quantity of the latter used, the quantity of actual acid present may be easily calculated. (See also volumetric methods in Chapter 37.) Analytical reactions. (Sodium acetate, NaC2H302, may be used.) 1. Any acetate heated with sulphuric acid evolves acetic acid, which may be recognized by its odor. 2. Acetic acids or acetates heated with sulphuric acid and alcohol give a characteristic odor of acetic ether. 3. A solution containing acetic acid, or an acetate carefully neutral- ized, turns deep red on the addition of solution of ferric chloride, and forms, on boiling, a reddish-brown precipitate of an oxyacetate of iron. Potassium acetate, Potassii acetas, KC2H302 = 98. Sodium acetate, Sodii acetas, NaC2H302.3H20 = 136. Zinc acetate, Zinci acetas, Zn (C2H302),.2H20 = 219. These three salts may be obtained by neu- 326 CONSIDERATION OF CARBON COMPOUNDS. tralizing the respective carbonates with acetic acid and evaporating the solution; they are white salts, easily soluble in water. Ferric acetate, Fe2(C2H302)6. A 33 per cent, solution of this salt is the Liquor Jerri acetatis of the U. S. P. It is made by dissolving freshly precipitated ferric hydroxide in acetic acid; it is a dark, red- brown, transparent liquid of a specific gravity of 1.16. Lead acetate, Plumbi acetas, Pb(C2H302)23H20 = 378.4 {Sugar of lead), is made by dissolving lead oxide in diluted acetic acid. It forms colorless, shining, transparent crystals, easily soluble in water; on heating, it melts and then loses water of crystallization ; at yet higher temperatures it is decomposed ; it has a sweetish, astringent, afterward metallic taste. Commercial sugar of lead contains often an excess of lead oxide in the form of basic salts; such an article when dissolved in spring water gives generally a turbid solution, in conse- quence of the formation of lead carbonate ; the addition of a few drops of acetic acid renders the liquid clear by dissolving the precipi- tate. When a mixture of lead acetate and lead oxide is digested or boiled with water, the acetate combines with the oxide, forming a basic lead acetate, Pb(C2II302)2 + 2PbO, a 25 per cent, solution of which is the Liquor plumbi subacetatis, or Goulard's extract, whilst a solution con- taining about 1 per cent, is the Liquor plumbi subacetatis dilutus, or lead-water. Cupric acetate, Cupric acetas, Cu(C2H302)2H20 =199.2 (Acetate of copper). The commercial verdigris is a basic acetate of copper, Cu(C2H302)2Cu0, made by the action of dilute acetic acid and atmos- pheric air on metallic copper. By adding to this basic acetate more acetic acid, the neutral acetate is obtained, but this may be made directly also by dissolving cupric hydroxide or carbonate in acetic acid. It forms deep green, prismatic crystals, which are soluble in water. By boiling verdigris with arsenous oxide, cupric aceto-arsenite, 3CuAs204 + Cu(C2H302)2, is formed, which is the chief constituent of emerald green or Schweinfurt green, a substance often used as a color- ing matter. Paris green is of a similar composition, but less pure. Chlor-acetic acids. By treating acetic acid with chlorine, either one, two, or three hydrogen atoms may be replaced by this element, when either mono-, di-, or tri-chlor acetic acid is formed. Trichlor-acetic acid, C2C13H02, is a color- MONOBASIC FATTY ACIDS. 327 less, crystalline substance, which fuses at 55° C. (131° F.), and boils at 195° C. (383° F.). Acetone, C3HcO. This compound is obtained by the destructive distillation of acetates (and of a number of other substances). The decomposition which calcium acetate suffers may be shown by the equation : CELCOO\n CHAnn . p ha CH3COO/Ca = CH3/C0 + CaC03- Calcium acetate. Acetone. The above graphic formula of acetone shows this substance to be dimethyl carbonyl, or carbon monoxide whose two available affiuities have been satisfied by two methyl groups. Acetoue is the represen- tative of a series of compounds known as acetones or generally as ketones, the general composition of which may be assumed to be R—C—R , R representing in this case any univalent radical. O Acetone is a colorless liquid, boiling at 58° C. (136° F.), miscible with water, alcohol, and ether in all proportions. Butyric acid, HC4H702. Among the glycerides of butter those of butyric acid are found; they exist also in cod-liver oil, croton oil, and a few other fatty oils; some volatile oils contain compound ethers of butyric acid; free butyric acid occurs in sweat and in cheese. It may be obtained by a peculiar fermentation of lactic acid (which itself is a product of fermentation), and is also generated during the putre- faction of albuminious substances. Butyric acid is a colorless liquid, having a characteristic, unpleasant odor; it mixes with water in all proportions. Valerianic acid, HC5H902 ( Valeric acid). This acid occurs in vale- rian root and angelica root, from which it may be separated ; it is, however, generally obtained by oxidation of amyl alcohol by potas- sium dichromate and sulphuric acid. After oxidation has taken place the mixture is distilled, when valerianic acid with some valerian- ate of amyl distils over. The change of amyl alcohol into valerianic acid is analogous to the conversion of ethyl alcohol into acetic acid: C5HuOH + 20 = HC5H902 + H20. Pure valerianic acid is an oily, colorless liquid, having a penetrat- ing, highly characteristic odor; it is slightly soluble in water, but soluble in alcohol; it boils at 185° C. (365° F.). Amyl alcohol. Valerianic acid. 328 CONSIDERATION OF CARBON COMPOUNDS. Several of the salts of valerianic acid are officinal; they are the valerianate of iron, of ammonium, of zinc, and of quinine. The last named three compounds are white salts, while the ferric valerianate has a dark-red color ; the ammonium salt is easily soluble in water, the other three compounds are insoluble or nearly so. Oleic acid, Acidum oleicum, HC18H3302 = 282. As shown by its formula, oleic acid does not belong to the above-described series of fatty acids of the composition CnII2n02, but to a series having the general composition CnH2n_202. Oleic acid is a constituent of most fats, especially of fat oils. Thus, olive oil is mainly oleate of glyceril. By boiling olive oil with potas- sium hydroxide, potassium oleate is formed, which may be decom- posed by tartaric acid, when oleic acid is liberated. Oleic acid is a nearly colorless or yellowish, odorless, tasteless, neu- tral liquid, insoluble in water, soluble in alcohol, chloroform, oil of turpentine, and fat oils, crystallizing near the freezing-point of water; exposed to the air it decomposes and shows then an acid reaction. Lead oleate is soluble in ether, lead palmitate and lead stearate are not. The officinal oleate of mercury and oleate of veratrine are obtained by dissolving the yellow mercuric oxide or veratrine in oleic acid. Questions.—421. What is the constitution of organic acids, which group of atoms is found in all of them, and how does an alcohol radical differ from an acid radical ? 422. Give some processes by which organic acids are formed in nature or artificially. 423. Mention the general properties of organic acids. 424. Which series of acids is known as fatty acids, and why has this name been given to them? 425. Mention names,composition, and occurrence in nature of the first five members of the series of fatty acids. 426. By what processes may formic acid be obtained, and what are its properties? 427. Describe the processes of manufacturing acetic acids from alcohol and from wood. 428. What is vinegar, and what is glacial acetic acid ? Give tests for acetic acid and for acetates. 429. Describe the processes for making the acetates of potassium, zinc, iron, lead, and copper, and also of Goulard’s ex- tract and lead-water; state their composition and properties. 430. When and in what form of combination is oleic acid found in nature, and what are its properties ? DIBASIC AND TRIBASIC ORGANIC ACIDS. 329 44. DIBASIC AND TRIBASIC ORGANIC ACIDS. Dibasic acids of the general composition C H2n_204. Oxalic acid H2C204 or (C02H)2. Malonic acid H2C3H204 or C H2(C02H)2. Succinic acid H2C4H404 or C2H4(C02H)2. Pyrotartaric acid .... H2C5H604 or C3H6(C02H)2. Adipic acid H2C6H804 or C4H8(C02H)2. etc. Of these acids, only the first member is of general interest. Oxalic acid, H2C204.2H20. This acid may be looked upon as a direct combination of two carboxyl groups, C02H—C02H, both atoms of hydrogen being replaceable by metals. Oxalic acid is distributed largely in the vegetable kingdom in the form of potassium, sodium, or calcium salts. It may be obtained from vegetables, or by the oxidation of many organic substances, chiefly fats, sugars, starch, etc., by nitric acid or other strong oxidiz- ing agents. Experiment 47. Pour a mixture of 15 c.c. nitric acid and 35 c.c. of water upon 10 grammes of sugar contained in a 200 c.c. flask. Apply heat gently until the reaction begins. When red fumes cease to escape pour the solution into a porcelain dish and evaporate to about one-half its volume. Crystals of oxalic acid separate on cooling; use them for making the analytical reactions mentioned below. Oxalic acid is manufactured on a large scale by heating sawdust with potassium or sodium hydroxide to about 250° C. (482° F.), when the oxalate of these metals is formed ; by the addition of cal- cium hydroxide to the dissolved alkali oxalate, insoluble calcium oxalate is formed which is decomposed by sulphuric acid. Oxalic acid crystallizes in large, transparent, colorless prisms, con- taining two molecules of water; it is soluble in water and alcohol, and has poisonous properties. When heated slowly, it sublimes at a temperature of about 155° C. (811° F.); but if heated higher or with sulphuric acid it is decomposed into water, carbonic oxide, and carbon dioxide: H2C304 = h2o + CO + C02. Oxalic acid acts as a reducing agent, decolorizing solutions of the permanganates, and precipitating gold and platinum from their solu- tions : PtCl4 + 2H3C204 = Ft + 4C02 + 4HC1. 330 CONSIDERATION OF CARBON COMPOUNDS. Analytical reactions. (Sodium oxalate, Na2C204, may be used.) 1. Oxalic acid or oxalates when heated with strong sulphuric acid evolve carbonic oxide and carbon dioxide (see above). 2. Neutral solutions of oxalic acid give with calcium chloride a white precipitate of calcium oxalate, CaC204, which is insoluble in acetic, soluble in hydrochloric acid. 3. Silver nitrate produces a white precipitate of silver oxalate, Ag2C20, 4. A dry oxalate (containing a non-volatile metal) heated in a test- tube evolves carbonic oxide, whilst a carbonate is left which shows elfervescence with acids. Antidotes to oxalic acid. Calcium carbonate or lime-water should be admin- istered, but no alkalies as in cases of poisoning by mineral acids, because the alkali oxalates are soluble. Oxalates. The acid potassium oxalate, KHC204, or its combina- tion with oxalic acid, is known under the name of salt of sorrel. Cal- cium oxalate, CaC204, is, in small quantities, a normal constituent of urine. Ferrous oxalate, ferri oxalas, FeC204.H20, is made by adding potassium or ammonium oxalate to ferrous sulphate, when double decomposition takes place, and the ferrous oxalate is precipitated as a pale-yellow, crystalline, nearly insoluble powder. Dibasic acids with alcoholic hydroxyl. /OH Malic acid = C4H605 or C2H3^-C02H 'C02H .OH //OH Tartaric acid = C4H606 or C2ll2 svco2h C02H In the various acids heretofore considered, the hydrogen is derived either from the unsaturated hydrocarbon residue, or from the hydroxyl in the carboxyl. As shown by the graphic formulas of the above two acids, they contain also hydrogen in the hydroxyl form not in combination with CO. This hydrogen, whilst not replaceable by metals, may be replaced by alcohol radicals; in other words, it be- haves like the hydroxyl hydrogen in alcohols. In order to indicate DIBASIC AND TRIBASIC ORGANIC ACIDS. this difference in the function of the hydrogen, malic acid is said to be dibasic, but triatomic; tartaric acid is dibasic and tetratomic. A few other acids behave in a similar manner, as, for instance, lactic acid. Malic acid, H2C4II405, occurs in the juices of many fruits, as apples, currants, etc. Tartaric acid, Acidum tartaricum, H2C4H406 = 150. Frequently found in vegetables, and especially in fruits, sometimes free, generally as the potassium or calcium salt; grapes contain it chiefly as potas- sium acid tartrate, which is obtained in an impure state as a by-product in the manufacture of wine. During the fermentation of grape-juice, its sugar is converted into alcohol; potassium acid tartrate is less soluble in alcoholic fluids than in water, and therefore is deposited gradually, forming the crude tartar, or argot of commerce, a substance containing chiefly potassium acid tartrate, but also calcium tartrate, some coloring matter, and often traces of other substances. Crude tartar is the source of tartaric acid and its salts. Tartaric acid is obtained from potassium acid tartrate by neutraliz- ing with calcium carbonate, and decomposing the remaining neutral potassium tartrate by calcium chloride : 2(KHC4H406) + CaC03 = CaC4H406 + K2C4H406 -f H20 + C02. Potassium acid tartrate. Calcium carbonate Calcium tartrate. Potassium tartrate. Water. Carbon dioxide. K2C4H406 + CaCl2 = CaC4H406 + 2KC1. Potassium tartrate. Calcium chloride. Calcium tartrate. Potassium chloride. The whole of the tartaric acid is thus converted into calcium tar- trate, which is precipitated as au insoluble powder; this is collected, well washed, and decomposed by boiling with sulphuric acid, when calcium sulphate is formed as au almost insoluble residue, while tar- taric acid is left in solution, from which it is obtained by evaporation and crystallization : ChC4H406 + H2S04 = H2C4H406 + CaS04. Calcium tartrate. Sulphuric acid. Tartaric acid. Calcium sulrhate. Tartaric acid crystallizes in colorless, transparent prisms; it has a strongly acid, but not disagreeable taste; it is readily soluble in water and alcohol, and fuses at 135° C. (275° F.). There are three acids which are isomeric with common tartaric acid, differing from it in physical, but not in chemical properties. These acids are known as inactive tartaric acid, levotartaric acid, and racemic acid, whilst the common tartaric acid is termed dextrotartaric acid. Crude tartar sometimes contains racemic acid. 332 CONSIDERATION OF CARBON COMPOUNDS. Analytical reactions. (Potassium sodium tartrate, KNaC4H406, may be used.) 1. Neutral solutions of tartaric acid give with calcium chloride a white precipitate of calcium tartrate, which, after being quickly col- lected on a filter and washed, is soluble in potassium hydroxide; from this solution calcium tartrate is precipitated on boiling. (Calcium citrate is insoluble in potassium hydroxide.) 2. A strong solution of a tartrate, acidulated with acetic acid, gives a white precipitate of potassium acid tartrate on the addition of potas- sium acetate. (Precipitate forms slowly.) 3. A neutral solution of a tartrate gives with silver nitrate a white precipitate of silver tartrate, Ag2C4II406, which blackens on boiling, in consequence of the decomposition of the salt, with separation of silver. If, before boiling, a drop of ammonia water be added, a mirror of metallic silver will form upon the glass. 4. Sulphuric acid heated with tartrates chars them readily. 5. Tartrates, when heated, are decomposed (blacken), and evolve a somewhat characteristic odor, resembling that of burnt sugar. The above reaction, 3, can be used to advantage for silvering glass by operat- ing as follows: Dissolve 1 gramme of silver nitrate in 20 c.c. of water, add water of ammonia until the precipitate which forms is nearly redissolved, and dilute with water to 100 c.c. Make a second solution by dissolving 0.2 gramme of silver nitrate in 100 c.c. of boiling water, add 0.166 gramme of potassium sodium tartrate, boil until the precipitate becomes gray, and filter. Mix the two solutions cold and set aside for one hour, when a mirror of metallic silver will be found. Potassium acid tartrate, Potassii bitartras, KHC4H406 = 188 (Bi- tartrate of potassium, Cream of tartar). The formation of this salt in the crude state (argol) has been explained above. It is purified by dissolving in hot water and crystallizing, when it is obtained in color- less crystals, or as a white, somewhat gritty powder of a pleasant, acidulous taste; it is sparingly soluble in cold, easily soluble in hot water. Potassium tartrate, Potassii tartras, 2(K2C4H406).H20 = 470 Tar- trate of 'potassium). Obtained by saturating a solution of potassium acid tartrate with potassium carbonate : Potassium acid tartrate. 2KHC4H406 + K2C03 = 2K2C4H4Ofi + H20 + C02. Potassium carbonate. Potassium tartrate. DIBASIC AND TRIBASIC ORGANIC ACIDS. 333 Small, transparent or white crystals, or a white neutral powder, soluble in less than its own weight of water. Potassium sodium tartrate, Potassii et sodii tartras, KNaC4H406. 4H20 = 282 (Tartrate of 'potassium and sodium, Rochelle salt). If in the above-described process for making neutral potassium tartrate, sodium carbonate is substituted for potassium carbonate, the double tartrate of potassium and sodium is formed. It is a white powder, or occurs in colorless, transparent crystals which are easily soluble in water. Experiment 48. Add gradually 24 grammes of potassium acid tartrate to a hot solution of 20 grammes of crystallized sodium carbonate in 100 c.c, of water. Heat until complete solution has taken place, filter, evaporate to about one-half the volume, and set aside for the potassium sodium tartrate to crys- tallize. How much crystallized sodium tartrate is required for the conversion of 25 grammes of potassium acid tartrate into Rochelle salt? Seidlitz powders consist of a mixture of 120 grains of Rochelle salt with 40 grains of sodium bicarbonate (wrapped in blue paper), and 35 grains of tartaric acid (wrapped in white paper). When dissolved in water the tartaric acid acts upon the sodium bicarbonate, causing the formation of sodium tartrate, while the escaping carbon dioxide causes effervescence. Antimony potassium tartrate, Antimonii et potassii tartras, 2(KSbO. C4H406).H20 = 664 (Tartrate of antimony and 'potassium, Potassium antimonyl tartrate, Tartar emetic). This salt is made by dissolving freshly prepared antimonious oxide (while yet moist) in a solution of potassium acid tartrate. From the solution somewhat evaporated, tartar emetic separates in colorless, transparent rhombic crystals : 2KHC4H406 + Sb203 = 2K Sb0.C4H406 + H20. Potassium acid tartrate. Antimonious oxide. Tartar emetic. The fact that not antimony itself, but the group SbO replaces the hydrogen, has led to the assumption of the hypothetical radical SbO, termed antimony]. Tartar emetic is soluble in water, insoluble in alcohol; it has a sweet, afterward disagreeable metallic taste. Action of certain organic acids upon certain metallic oxides. The solution of a ferric salt (or certain other metallic salts) is precipitated by alkaline hydroxides, a salt of the alkali and ferric hydroxide being 334 formed. When a sufficient quantity of either tartaric, citric, oxalic, or various other organic acids has been added previously to the iron solution (or to certain other metallic solutions) no such precipitate is produced by the alkaline hydroxides, because organic salts or double salts are formed which are soluble, and from which the metallic hydroxides are not precipitated by alkaline hydroxides. Upon evaporation no crystals (of the organic salt) form, and in order to obtain the compounds in a dry state, the liquid, after being evaporated to the consistence of a syrup, is spread on glass plates which are exposed to a temperature not exceeding 60° C. (140° F.), when brown, green, or yellowish-green, amorphous, shining, transparent scales are formed, which are the scale compounds of the U. S. P. Instead of obtaining these compounds, as stated above, by adding the organic acids (or their salts) to the inorganic salts, they are more generally obtained by dissolving the freshly precipitated 'metallic hydroxide in the organic acid. The true chemical constitution of many of these scale compounds has as yet not been determined with certainty. Of officinal scale compounds containing tartaric acid may be men- tioned the tartrate of iron and ammonium, and the tartrate of iron and potassium. The first compound is obtained by dissolving freshly precipitated ferric hydroxide in a solution of ammonium acid tartrate, the second by dissolving ferric hydroxide in potassium acid tartrate. The clear solutions, after having been sufficiently evaporated, are dried, as mentioned above, on glass plates. CONSIDERATION OF CARBON COMPOUNDS. Citric acid, Acidum citricum, H3C6Hf)07.H20 = 210. Citric acid is a tribasic acid containing three atoms of hydrogen replaceable by metals; its constitution may be expressed by the graphic formula : OH /\CO,H OA \Wh 2h Citric acid is found in the juices of many fruits (strawberry, rasp- berry, currant, cherry, etc), and in other parts of plants. It is obtained from the juice of lemons by saturating it with calcium car- bonate and decomposing by sulphuric acid the calcium citrate thus formed. (100 parts of lemons yield about 5 parts of the acid.) It forms colorless crystals, easily soluble in water. DIBASIC AND TRIBASIC ORGANIC ACIDS. 335 Analytical reactions. (Potassium citrate, K3C6H507> may be used.) 1. Neutral solutions of citrates yield with calcium chloride on boil- ing (not in the cold) a white precipitate of calcium citrate, which is insoluble in potassium hydroxide, but soluble in cupric chloride. 2. Neutral solutions of citrates are precipitated white by silver nitrate. The precipitate does not blacken on boiling, as in the case of tartrates. 3. A neutral or alkaline solution of a citrate to which a few drops of a solution of potassium permanganate have been added, becomes green or reddish-green. Tartrates decolorize permanganate. Citrates. Potassium citrate, K3CfiH507.H2O, Lithium citrate, Li3C6H507, and Magnesium citrate, Mg3(C6H507)2.14H02, are color- less substances, easily soluble in water and obtained by dissolving the carbonates in citric acid. The effervescent citrates of potassium, lithium, and magnesium, are mixtures of citric acid with potassium bicarbonate, lithium carbonate, and magnesium carbonate respectively ; sugar is added to all, and some sodium bicarbonate to the two last preparations. Bismuth citrate, BiC6H507, is obtained by boiling a solution of citric acid with bismuth nitrate, when the latter is gradually converted into citrate whilst nitric acid is set free; the insoluble bismuth citrate is collected, washed, and dried ; it forms a white, amorphous powder, which is insoluble in water, but soluble in water of ammonia. Bismuth ammonium citrate is a scale compound obtained by dissolv- ing bismuth citrate in water of ammonia and evaporating the solution at a low temperature. Ferric citrate, Ferri citras (Fe22(C6H507).6H202?). Obtained in transparent, red scales, by dissolving ferric hydroxide in citric acid and evaporating the solution as mentioned heretofore. By mixing solu- tion of ferric citrate with either water of ammonia, or quinine, strych- nine, sodium phosphate, or sodium pyrophosphate, evaporating to the consistence of syrup and drying on glass plates, the following scale compounds are obtained respectively : Soluble citrate of iron (citrate of iron and ammonia), citrate of iron and quinine, citrate of iron and strychnine, soluble ferric phosphate, and soluble ferric pyrophosphate. Lactic acid, Acidum lacticum, HC3H503—90. This acid is the second member of a group of monobasic, diatomic acids which have 336 CONSIDERATION OF CARBON COMPOUNDS. the general composition CnH2a03, and which contain two hydroxyl groups, the hydrogen of one being capable of replacement by metals, the other by alcohols. The first member of this series is glycolic acid, HC2H303, a white, deliquescent, crystalline substance which is found in unripe grapes and in the leaves of the wild grape. Glycolic acid has been shown to be acetic acid, C2H402, in which one atom of hydrogen has been replaced by hydroxyl. The name hydroxy-acetic acid, has, therefore, been given to this compound. Lactic acid occurs in many plant-juices; it is formed from sugar by a peculiar fermentation known as “ lactic fermentation,” which causes the presence of this acid in sour milk and in many sour, fer- mented substances, as in ensilage, sauer-kraut, etc. The formation of lactic acid from sugar may be expressed by the equation : C6H,206 = 2(HC3H508). Sugar. Lactic acid. For practical purposes lactic acid is made by mixing a solution of sugar with milk, putrid cheese, and chalk, and digesting this mixture for several weeks at a temperature of about 30° C. (86° F.). The bacteria in the cheese act as a ferment, and the chalk neutralizes the acid generated during the fermentation. The calcium lactate thus obtained is purified by crystallization and decomposed by oxalic acid, which forms insoluble calcium oxalate. Lactic acid is a colorless, syrupy liquid, of strongly acid properties; it mixes in all proportions with water and alcohol. A lactic acid called sarco-lactic acid is found in meat-juice, and, therefore, as a constituent of meat-extract. This acid has the compo- sition and all the properties of the above ordinary lactic acid, with the exception that it acts differently on polarized light. Ferrous lactate, Ferri lactas, Fe(C3H503)2.3H20 = 287.9. Made by dissolving iron filings in diluted lactic acid ; hydrogen is liberated and the salt formed. It is a pale, greenish-white, crystalline sub- stance, soluble in water. Questions.—431. Name the more common organic acids found in vegeta- bles and especially in sour fruits. 432. What is the composition of oxalic acid, how is it manufactured, and what are its properties? 433. Explain the formation of crude tartar during the fermentation of grape-juice, and how is tartaric acid obtained from it ? 434. Give properties of and tests for tartaric acid. 435. State the composition and formation of cream of tartar, Rochelle salt, and tartar emetic. 436. What are Seidlitz powders, and what changes take place when they are dissolved? 437. Mention some officinal scale com- ETHERS. 337 45. ETHERS. Constitution. It has been shown that alcohols are hydrocarbon residues in combination with hydroxyl, OH, and that acids are hydro- carbon residues in combination with carboxyl, CO.OH; it has further been shown that carboxyl may be considered as being composed of CO, and hydroxyl, OH, and that the term acid radical is applied to that group of atoms in acids which embraces the hydrocarbon residue + CO. If we represent an alcohol radical by AIR, and an acid radical by AcR, the general formula of an alcohol is A1R.OH, or A®>0, and of an acid, AcR.OH, or Ethers are formed by replacement of the hydrogen of the hydroxyl— in alcohols by hydrocarbon residues (or alcohol radicals), and com- pound ethers or esters are formed by replacement of the hydrogen of the hydroxyl (or carboxyl) in acids by hydrocarbon residues. While alcohols correspond in their constitution to hydroxides, ethers corre- spond to oxides, and compound ethers to salts. For instance: Hydroxides. Oxides. Acids. Salts. KOH = k2o = |>0 HN03 = N^\o kno3 = n^\o Potassium hydroxide. Potassium oxide. Nitric acid. Potassium nitrate. C2Hs\() H/U c2h5\0 c2h5/u c2h2o\q H/U c2ho\ c2h5/u Ethyl alcohol. Ethyl ether. Acetic acid. Ethyl acetate, or acetic ether. O /\ A1R\ A1R/0 AcR\n H/U AcR\0 A1R/U> Alcohol. Ether. Acid. Compound ether. It is not necessary that the two hydrocarbon residues in an ether should be alike, as in the above ethyl ether, but they may be different, in which case the ethers are termed mixed ethers. For instance: CH3.C2H50 = gagj\0 C3H7.C5Hu.O = cJhJj/O- Methyl-ethyl ether. Propyl-amyl ether. In diatomic or triatomic alcohols, or in dibasic or tribasic acids, containing more than one atom of hydrogen derived from hydroxyl or carboxyl, these hydrogen atoms may be replaced by various other univalent, bivalent, or trivalent residues. This fact shows that the pounds of iron, and give a general outline of the mode of preparing them. 438. From what and by what process is citric acid obtained ? 439. Mention tests by which citric acid may be distinguished from tartaric acid. 440. From what, and by what process is lactic acid obtained ; what are its properties ? 338 CONSIDERATION OF CARBON COMPOUNDS. number of ethers or compound ethers which are capable of being formed is very large. Formation of ethers. Ethers may be formed by the action of the chloride or iodide of a hydrocarbon residue upon an alcohol, in which the hydroxyl hydrogen has been rephiced by a metal. For instance: Cy®5\o _L OKI C.2H5\q -M- r Na/U + — C2H5/u r iNHl- Sodium ethylate. Ethyl iodide. Ethyl ether. Sodium iodide. -1- C,n T 4- 'N'hT Na/U + (,asL _ ch3/u JNai> Sodium ethylate. Methyl iodide. Ethyl-methyl ether. Sodium iodide. Ethers are also formed by the action of sulphuric acid upon alco- hols ; the sulphuric acid removing water in this case, thus: 2(C2H5OH) = S255\o + H20. 'J2n 5/ Compound ethers are formed by the combination of acids with alcohols and elimination of water. (Presence of sulphuric acids facili- tates this action.) Ethyl alcohol. Ethyl ether. Water. \() i_ 2H30\q C2H/)\ . HO H/u ' H/U — C2H5/U + They are also formed by the action of hydrocarbon chlorides (or iodides) on salts. For instance : Ethyl alcohol. Acetic acid. Ethyl acetate. Water. C6HnCl + CII«>0 = \0 + KC1. Amyl chloride. Potassium formate. Amyl formate. Potassium chloride. Occurrence in nature. Many ethers are products of vegetable life and occur in some essential oils; wax contains the compound ether palmitate of melissyl, C30H61.O16H31O.O, and spermaceti, a solid sub- stance found in the head of the whale, is the palmitate of cethyl, C16H33.C16H310.0. The most important group of compound ethers are the fats and fatty oils, which are distributed widely in the vege- table, but even more so in the animal kingdom. General properties. The ethers and compound ethers of the lower members of the monatomic alcohols and fatty acids have generally a characteristic and pleasant odor. Fruit essences consist mainly of such compound ethers, and what is generally known as the “bouquet” or “flavor” of wine and other alcoholic liquors is due chiefly to ETHERS. ethers or compound ethers, which are formed during (and after) the fermentation by the action of the acids present upon the alcohol or the alcohols formed. The improvement which such alcoholic liquids undergo “by age” is caused by a continued chemical action between the substances named, All ethers are neutral substances; those formed by the lower alco- hols and acids are generally volatile liquids, those of the higher members are non-volatile solids. When compound ethers are heated with alkalies, the acid combines with the latter, whilst the alcohol is liberated. (The properties of the compound ethers, termed fats, will be considered further on.) Ethyl ether, iEther, (C2H5)20 = 74 [Ether, Sulphuric ether, Ethyl oxide). The name of the whole group of ethers is derived from this (ethyl-) ether, in the same way that common (ethyl-) alcohol has given its name to the group of alcohols. The name sulphuric ether was given at a time when its true composition was yet unknown, and for the reason that sulphuric acid was used in its manufacture. Ether is manufactured by heating to about 140° C. (284 F.) a mix- ture of 1 part of alcohol and 1.8 parts of concentrated sulphuric acid in a retort, which is so arranged that additonal quantities of alcohol may be allowed to flow into it, while the open end is connected with a tube, leading through a suitable cooler, in order to condense the highly volatile product of the distillation. Experiment 49. Mix 100 grammes of alcohol with 180 grammes of ordinary- sulphuric acid, allow to stand and pour the cooled mixture into a flask which is provided with a perforated cork through which pass a thermometer and a bent glass tube leading to a Liebig’s condenser. Apply heat and notice that the liquid commences to boil at about 140° C. (284° F.). Distil about 50 c.e., pour this liquid into a stoppered bottle and add an equal volume of water. Ethyl ether will separate in a distinct layer over the water, and may be removed by means of a pipette. Repeat the washing with water, add to the ether thus freed from alcohol a little calcium chloride and distil it from a dry flask, standing in a water-bath. The greatest care should be exercised and the neighborhood of flames avoided in working with ether, on account of its volatility and the inflammability of its vapors. The apparatus described above for etherification can be constructed so as to make the process continuous. This may be done by using with the boiling- flask a cork with a third aperture through which a glass tube passes into the liquid. The other end of the tube is connected by means of rubber tubing with a vessel filled with alcohol and standing somewhat above the flask. As soon as distillation commences alcohol is allowed to flow into the flask at a rate equal to that of the distillation, keeping the temperature at about 140° C. (284° F.). The flow of alcohol is regulated by a stop-cock. 340 CONSIDERATION OF CARBON COMPOUNDS. The action of sulphuric acid upon alcohol is not quite so simple as described above in connection with the general methods for obtaining ethers, where the final result only was given. An intermediate pro- duct, known as ethyl-sulphuric acid or sulpho-vinic acid, is formed, which, by acting upon another molecule of alcohol, forms sulphuric acid and ether, which latter is volatilized as soon as formed. The decomposition is shown by the equations : Alcohol. c2h5oh + h2so4 = c2h5hso4 + h2o. Sulphuric acid. Ethyl-sulphuric acid. Water. C2H5HS04 + C2H5OH = H2S04 + (C2h5)2o Ethyl-sulphuric acid. Alcohol. Sulphuric acid. Ether, The liberated sulphuric acid at once attacks another molecule of alcohol, again forming ethyl-sulphuric acid, which is again decomposed, etc. Theo- retically, a given quantity of sulphuric acid should be capable, therefore, of converting any quantity of alcohol into ether; practically, however, this is not the case, because secondary reactions take place simultaneously, and because the water which is constantly formed does not all distil with the ether, and, therefore, dilutes the acid to such an extent that it no longer acts upon the alcohol. Ether thus obtained is not pure, but contains water, alcohol, sulphurous and sulphuric acids, etc.; it is purified by mixing it with chloride and oxide of calcium, pouring off the clear liquid and distilling it. Pare ether is a very mobile, colorless, highly volatile liquid, of a refreshing, characteristic odor, a burning and sweetish taste, and a neutral reaction; it is soluble in alcohol, chloroform, liquid hydro- carbons, fixed and volatile oils, and dissolves in eight volumes of water. Specific gravity at 0° C. (32° F.), is 0.720; boiling-point 35° C. (95° F.). It is easily combustible and burns with a luminous flame. When inhaled, it causes intoxication and then loss of con- sciousness and sensation. The great volatility and combustibility of ether necessitate special care in the handling of this substance near fire or light. Stronger ether, JEther fortior of the U. S. P., contains about 94 per cent, of pure ether and 6 per cent, of alcohol, with a little water, while ether, aether of the U. S. P., contains 74 per cent, of pure ether and 26 per cent, of alcohol, with some water. Spiritus cetheris and Spiritus cetheris eompositus are mixtures of about one part of ether and two parts of alcohol, 3 per cent, of cer- tain ethereal oils.being added to the second preparation. Acetic ether, iEther aceticus, C2H5C2H302 = 88 (Ethyl acetate). Made by mixing dried sodium acetate with alcohol and sulphuric acid, ETHEES. 341 distilling and purifying the crude product by shaking with calcium chloride and rectifying: C2H5OH + NaC2H302 + H2S04 = C2H5C2H302 + NaHS04. Ethyl alcohol. Sodium acetate. Acetic ether. Sodium acid sulphate. Experiment 50. Add to a mixture of 40 grammes of pure alcohol and 100 grammes of concentrated sulphuric acid 60 grammes of sodium acetate. In- troduce this mixture into a boiling-flask, connect it with a Liebig’s condenser and distil about 50 c.c. Redistil the liquid from a flask, as represented in Fig. 39, page 297, and collect the portion which passes over at a temperature of 77° C. (170° F.); it is nearly pure ethyl acetate. Acetic ether is a colorless, neutral, and mobile liquid, of a strong ethereal and somewhat acetous odor, soluble in alcohol, ether, chloro- form, etc., in all proportions, and in 17 parts of water. Specific gravity 0.889. Boiling-point 76° C. (169° F.). Ethyl nitrite, C2H.N02 (Nitrous ether). Made by distilling a mix- ture of alcohol, sulphuric and nitric acids at a temperature of 80° C. (176° F.). By the deoxidizing action of alcohol on nitric acid, HN03, the latter is converted into nitrous acid, HN02, which in its turn acts on alcohol, the two substances combining with elimination of water, which is absorbed by the sulphuric acid : c2h5oh + hno2 = c2h5.no2 + h2o. Spirit of nitrous ether, Spiritus cetheris nitrosi, Sweet spirit of nitre. This is a mixture of 5 parts of the crude ethyl nitrite with 95 parts of alcohol. It is a clear, mobile, volatile, and inflammable liquid, of a pale straw color inclining slightly to green, a fragrant, ethereal odor, and a sharp, burning taste. It slightly reddens litmus paper, but evolves no carbon dioxide with carbonates. Amyl nitrite, Amyli nitris, C5HuN02 = 117 (Nitrite of amyl). Made by distilling equal volumes of pure amyl alcohol and nitric acid from a retort until an inserted thermometer shows a temperature of 100° C- (212° F.). The crude distillate is purified by agitating it with a solution of potassium carbonate and hydroxide, separating the upper layer of the liquid and redistilling it ; the liquid passing over between 96° C. and 100° C. (205° F. and 212° F.) is the amyl nitrite. It is a clear, pale-yellowish liquid, of an ethereal, fruity odor, an aromatic taste, and a neutral or slightly acid reaction. Specific gravity 0.872. Boiling-point 96° C. (205° F.). 342 CONSIDERATION OF CARBON COMPOUNDS. Fats and fat oils. All true fats are compound ethers of the tri- atomic alcohol glycerin, in which the three replaceable hydrogen atoms of the hydroxyl are replaced by three univalent radicals of the higher members of the fatty acids. For instance : /OH Glycerin = C3H5.(OH)3 or C3H5/OH \OH Stearic acid = C,8Ha5O.OH or C,»H„0\a / (^18^-35^ Stearin or tristearin = C3H5.(C18H350)3.03 or C3H5—(C18H350).0 \(C18H350).0 While all natural fats are glycerin in which the three hydrogen atoms are replaced, we may by artificial means introduce but one or two acid radicals, thus forming: /(C18H350)0 Monostearin = C3H \OH /(C18H350)0 Distearin = C3H5^-(C18H350)0 ' \(0H Fats are often termed glycerides; stearin being, for instance, the glyceride of stearic acid. The principal fats consist of mixtures of palmitin,C3H5.(C16H310)3.03, stearin, C3H5.(C18H350)3.03, and olein, C3H5.(C18H330)3.03. Stearin and palmitin are solids, olein is a liquid at ordinary temperature; the relative quantity of the three fats mentioned determines its solid or liquid condition. The liquid fats, containing generally olein as their chief costituent, are called fatty oils or fixed oils in contradistinction to volatile or essential oils. All fats, when in a pure state, are colorless, odorless, and tasteless substances, which stain paper permanently; they are insoluble in water, difficultly soluble in cold’alcohol, easily soluble in ether, disul- phide of carbon, benzene, etc. The taste and color of fats are due to foreign substances, often produced by a slight decomposition which has taken place in some of the fat. All fats are lighter than water, and all solid fats fuse below 100° C. (212° F.); fats can be distilled without change at about 300° C. (572° F.), but are decomposed at a higher temperature with the formation of numerous products, some of which have an extremely disagreeable odor, as, for instance, acrolein, C3H40, an aldehyde which in composition is equal to glycerin minus two molecules of water : C3H5(HO)3 - 2H20 = C3H40. F Some fats keep without change when pure; since they contain, however, impurities generally, such as albuminous matter, etc., they ETHERS. 343 suffer decomposition (a kind of fermentation aided by oxidation), which results in a liberation of the fatty acids, which impart their odor and taste to the fats, causing them to become what is generally termed rancid. Some fats, especially some oils, suffer oxidation, which renders them hard. These drying oils differ from other oils in being mixtures of olein with another class of glycerides, containing unsaturated acids with less hydrogen in relation to carbon than oleic acid. Drying oils are prevented from drying by albuminous impurities, which may be removed by treating the oil with 4 per cent, of concentrated sulphuric acid; the acid does not act on the fat, but quickly destroys the albu- minous matters, which, with the sulphuric acid, sink to the bottom, whilst the “refined” oil may be removed by decantation. Fats are largely distributed in the animal and vegetable kingdoms. They exist in plants chiefly in the seeds, while in animals they are found generally under the skin, around the intestines, and on the muscles. Human fat, beef tallow, mutton tallow, and lard, are mixtures of palmitin and stearin with some olein. Butter consists of the glycer- ides of butyric acid, caproic acid, caprylic acid, and capric acid, which are volatile with water vapors, and of myristic, palmitic, and stearic acids, which are not volatile. The principal non-drying vegetable oils (consisting chiefly of olein) are olive oil, cottonseed oil, cocoanut oil, palm oil, almond oil. Among the drying oils are of importance: linseed oil, castor oil, croton oil, hemp oil, cod-liver oil. Whenever fats are treated with alkaline hydroxides, or with a num- ber of other metallic oxides, decomposition takes place, the fatty acids combining with the metals, whilst glycerin is set free. Some of the substances thus formed are of great importance, as, for instance, the various kinds of soap. Soap. Any fat boiled with sodium or potassium hydroxide will form soap. Soft soap is potassium soap, hard soap is sodium soap. The better kinds of hard soap are made by boiling olive oil with sodium hydroxide : C3H5(C18H3302)s + 3NaOH = 3NaC18H3302 + C3H5(OH)3. Oleate of glyceryl (olive oil). Sodium hydroxide. Sodium oleate (hard soap). Glycerin, Experiment 51. Boil 50 grammes of olive oil with 60 c.c. of a 15 per cent, sodium hydroxide solution for about one hour. The soap which is thereby formed remains dissolved in or mixed with water and glycerin. Cause sepa- 344 CONSIDERATION OF CARBON COMPOUNDS. ration by adding a solution of 15 grammes of sodium chloride in 40 c.c. of water and boiling for a short while, when the soap, which is insoluble in the salt solution, rises to the surface and solidifies on cooling. Soaps are soluble in water and alcohol; they contain rarely less than 30 per cent., but sometimes as much as 70-80 per ceut. of water. Ammonia liniment, Linimentum ammonice, and lime liniment, Lini- mentum calcis, are obtained by mixing cottonseed oil with water of ammonia and lime-water, respectively. The oleate of ammonium or calcium is formed, and remains mixed with the liberated glycerin. Lead plaster, Emplastrum plumbi. Chiefly lead oleate, Pb(C18H33 02)2. Obtained by boiling lead oxide with olive oil and water for several hours, until a homogeneous mass is formed. Lead oleate differs from the oleates of the alkalies by its complete insolubility in water. Lanolin. This name has been given to the fat or fats which are found in sheep’s wool and are obtained by treating the wool with soap-water, and acidi- fying the wash liquor, when the fats separate unchanged. These fats differ from the fats spoken of above in so far as the alcohol present is not glycerin, but an alcohol, or rather two isomeric alcohols of the composition C26H43OH and known as cholesterin and iso-cholesterin. These alcohols, which are white, crystalline, fusible substances, when in combination with fatty acids form the compound ethers known as lanolin. Lanolin is a yellowish-white (or, when not sufficiently purified, a more or less brownish), fat-like substance, having the peculiar odor of sheep’s wool and fusing at a moderate temperature, forming an oily liquid. Unlike true fats, lanolin is capable of mixing with upward of 30 per cent, of water or aqueous solutions and yet retaining its fatty consistency; it is, moreover, much less liable to decompose than fats, and it is this property and its power to mix with aqueous solutions which have rendered lanolin a valuable agent in cer- tain pharmaceutical preparations. Questions.—441. Explain the constitution of simple, mixed, and compound ethers. To what inorganic compounds are they analogous ? 442. State the general processes for the formation of ethers and compound ethers. 443. What is the composition of ethyl ether? Explain the process of its manufac- ture in words and symbols, and state its properties. 444. How is acetic ether made, and what are its properties? 445. What is sweet spirit of nitre, and how is it made? 446. State the general composition of fats and the chief con- stituents of tallow, butter, and olive oil. 447. What is the solubility of fats in water, alcohol, and ether; how do heat and oxygen act upon them ; what is the cause of their becoming rancid? 448. Explain the composition and manufacture of soap, and state the difference between hard and soft soap. 449. How are ammonia liniment, lime liniment, and lead plaster made, and what is their composition ? 450. What is the source of lanolin ; what are its constituents and properties ? CARBOHYDRATES. 345 46. CARBOHYDRATES. Constitution. The term carbohydrates or carbhydrates is not well chosen, because it implies that these substances are carbon in combi- nation with water. Carbohydrates do contain hydrogen and oxygen in the proportion of two atoms of hydrogen to one atom of oxygen, or in the proportion to form water, but this does not exist as such in the carbohydrates. The true atomic structure of carbohydrates is as yet but little known. The compounds of the composition C6H1206 are now looked upon as the aldehyde of the hexatomic alcohol mannite, C6H14Og, the chief constituent of manna : C6Hu06 - 2H = C6H1206. Mannite itself is formed from the saturated hydrocarbon C6H14, by replacement of 6 atoms of hydrogen by 60H; its constitutional formula is, therefore, (C6H8)vi.(OH)6. Carbohydrates generally contain 6 atoms of carbon or a multiple of 6. Properties. Carbohydrates are either fermentable, or can, .in most cases, be converted into substances which are capable of fermentation. They are not volatile, but suffer decomposition when sufficiently heated ; they have neither acid nor basic properties, but are of a neu- tral reaction. Oxidizing agents convert them into saccharic and mucic acids and finally into oxalic acid. (Soluble carbohydrates have the property of bending the plane of polarized light.) Most carbohydrates are white, solid substances, and, with the ex- ception of a few, soluble in water. The members of the first two groups (glucoses and saccharoses) have a more or less sweet taste. Many of them, especially glucoses, are good reducing agents, as is shown by the fact that they deoxidize in alkaline solution salts (or oxides) of copper, bismuth, mercury, gold, etc., either to a lower state of oxidation or to the metallic state. Occurrence in nature. No other organic substances are found in such immense quantities in the vegetable kingdom as the members ot this group, cellulose being a chief constituent of all starch and various kinds of sugar of most plants. Carbohydrates are also found as products of animal life, as, for instance, the sugar in milk, in bees’ honey, etc. 346 CONSIDERATION OF CARBON COMPOUNDS. Groups of carbohydrates. r Glucoses. c6huo. Grape-sugar, Saccharoses. C12H22On. Cane-sugar, Amyloses. C6H10O5. Starch, Origin ■ Vegetable - Fruit-sugar, Mannitose, Melitose, Maltose, Dextrin, Gums, Animal I nosite. Milk-sugar. Cellulose, Glycogen. Grape-sugar, C6H]206 (Ordinary glucose, Dextrose). This substance is very abundantly diffused throughout the vegetable kingdom, and is generally accompanied by fruit-sugar. It is contained in large quan- tities in the juice of many fruits; the percentage of grape-sugar in the dried fig is about 65, in grape 10-20, in cherry 11, in mulberry 9, in strawberry 6, etc. Grape-sugar is found also in honey and in minute quantities in the normal blood (0.1 per cent, or less), and traces occur, perhaps, in normal urine, the quantity in both liquids rising, however, during certain diseases, as high as 5 per cent, or higher. Grape-sugar is produced in the plant from starch by the action of the vegetable aeids present; it may be obtained artificially from starch (and from many other carbohydrates) by heating with dilute mineral (sulphuric) acids, which convert starch first into dextrin and then into grape-sugar. Corn-starch is now largely used for that pur- pose, the excess of sulphuric acid being removed by treating the solu- tion with chalk; the filtered solution is either evaporated to a syrup and sold as “ glucose,” or evaporated to dryness, when the commercial “ grape-sugar ” is obtained. Experiment 52. Heat to boiling 100 c.c. of a 1 per cent, sulphuric acid and add to it very gradually and under constant stirring a mixture made by rub- bing together 25 grammes of starch and 25 grammes of water. Continue to boil until iodine no longer causes a blue color (which shows complete conver- sion of starch into either dextrin or glucose), and until 1 c.c. of the solution is no longer precipitated on the addition of 6 c.c. of alcohol (which shows the conversion of dextrin into sugar, dextrin being precipitated by alcohol). Apply to a portion of the glucose solution thus obtained, and neutralized by sodium carbonate, the tests mentioned below. To the remaining solution add a quantity of precipitated calcium carbonate sufficient to convert all sulphuric acid into calcium sulphate. Filter, evaporate the solution to a syrup and notice its sweet taste. Glucose is met with generally as a thick syrup which crystallizes with difficulty, combining during crystallization with one molecule of water; but anhydrous crystals, closely resembling those of cane- CARBOHYDRATES. 347 sugar, are also known. Glucose is soluble in its own weight of water and is less sweet than cane-sugar, the sweetness of glucose com- pared to that of cane-sugar being about 3 to 5 ; when heated to 170° C. (338° F.) it loses water, and is converted into glucosan, C6Hl0O5; by stronger heating it loses more water and forms caramel, a mixture of various substances; it turns the plane of polarized light to the right. Grape-sugar combines with various metallic oxides (alkalies, alka- line earths, etc.), and also with a number of other substances, forming a series of compounds known as glucosides. Grape-sugar may be recognized analytically: 1. By causing a bright-red precipitate of cuprous oxide, when boiled with a solution of cupric sulphate in sodium hydroxide, to which tartaric acid has been added. (A solution containing these three substances in definite proportions is known as Fehling’s solu- tion. See index.) 2. By precipitating metallic silver, bismuth, and mercury, when compounds of these metals are heated with it in the presence of caustic alkalies 3. By easily fermenting when yeast is added to the solution, alcohol and carbon dioxide being formed : C6H1206 = 2C2H5OH + 2C02. Fruit-sugar, C6H1206 {Levulose), occurs with glucose in sweet fruits and honey; it resembles glucose in most chemical and physical prop- erties, but does not crystallize from an aqueous solution ; it may, however, be obtained in white, silky needles from an alcoholic solu- tion ; it is met with generally as a thick syrup, is about as sweet as cane-sugar, and turns the plane of polarized light to the left; it is formed by the action of dilute mineral acids or ferments on cane- sugar, which latter takes up water and breaks up thus : CuHnOu + H20 = c6H12o6 + g6h12o6. Cane-sugar. Dextrose. Levulose. Mannitose, CGH1206. Obtained by the oxidation of mannite; it does not crystallize and resembles grape-sugar. Galactose, C6H1206, is formed together with dextrose when either milk-sugar or gum-arabic is boiled with dilute sulphuric acid. Galac- tose crystallizes, reduces an alkaline copper solution, but does not ferment with yeast. Inosite, C6H1206 (.Muscle-sugar), occurs in various muscular tissues, in the lungs, kidneys, liver, spleen, brain, and blood. Although 348 CONSIDERATION OF CARBON COMPOUNDS. identical in composition with grape-sugar, inosite differs from the latter in not being fermentable and by not precipitating cuprous oxide from alkaline copper-solutions. Cane-sugar, Saccharum, C12H22On == 342 (Ordinary saccharose, Common sugar, Beet-sugar). Cane-sugar is found in the juices of many plants, especially in that of the different grasses (sugar-cane), and also in the sap of several forest trees (maple), in the roots, stems, and other parts of various plants (sugar-beet), etc. Plants containing cane-sugar do not contain free organic acids, which latter would convert it into grape-sugar. Cane-sugar is manufactured from various plants containing it by crushing them between rollers, expressing the juice, heating and adding to it milk of lime, which precipitates vegetable albuminous matter. The clear liquid is evaporated to the consistency of a syrup, which is further purified (refined) by filtering it through bone-black and evaporating the solution in “vacuum pans” to the crystallizing- point; the mother-liquors are further evaporated, and yield lower grades of sugar; finally a syrup is left which is known as molasses. Cane-sugar forms white, hard, distinctly crystalline granules, but may be obtained also in well-formed, large, monoclinic prisms. It dissolves in 0.2 part of boiling, in 0.5 part of cold water, and in 175 parts of alcohol; when heated to 160° C. (320° F.) it fuses, and the liquid, on cooliug, forms an amorphous, transparent mass, known as barley sugar; at a higher temperature cane-sugar is decomposed, water is evolved, and a brown, almost tasteless substance is formed, which is known as caramel or burnt sugar. Oxidizing agents act energetically upon cane-sugar, which is a strong reducing agent. A mixture of cane-sugar and potassium chlorate will deflagrate when moistened with sulphuric acid ; potassium permanganate is readily deoxidized in acid solution; cane-sugar, however, does not affect an alkaline copper-solution, and does not ferment itself; but when heated with dilute acids or left in contact with yeast for some time, it is decomposed into dextrose and levulose, both of which are fermentable. Like dextrose, cane-sugar forms compounds with metals, metallic oxides, and salts, which compounds are known as sucrates. Experiment 53. Make a one per cent, cane-sugar solution; test it with Fehling’s solution and notice that no cuprous oxide is precipitated. Add to 50 c.c. of the cane-sugar solution 5 drops of hydrochloric acid and heat on a water-bath for half an hour. Again examine the liquid with Fehling’s solu- tion; a precipitate of cuprous oxide is now formed, proving the conversion of cane-sugar into glucose. CARBOHYDBATES. 349 Maltose, CAOn, is obtained by the action of diastase on starch. Diastase is a substance formed during the germination of various seeds (rye, wheat, barley, etc.), and it is for this reason that grain, used for alcoholic liquors, is allowed to germinate, during which pro- cess diastase is formed, which, acting upon the starch present, converts it into maltose and dextrin : 3(C6H10O5) + h2o = c12h22ou + C6H10O5. Starch. Maltose. Dextrin. Maltose is also formed by the action of dilute sulphuric acid upon starch, and is hence often present in commercial glucose; by further treatment with sulphuric acid it is converted into dextrose. Maltose crystallizes, reduces alkaline copper solutions, and ferments with yeast. Melitose, C12II22On, is the chief constituent of Australian manna. Milk-sugar, Saccharum lactis, C12H220u + H20 = 360 {Lactose). Found almost exclusively in the milk of the mammalia. Obtained by freeing milk from casein and fat and evaporating the remaining liquid (whey) to a small bulk, when the milk-sugar crystallizes on cooling. It forms white, hard, crystalline masses; it is soluble in 7 parts of water (at 15° C., 59° F.) and in 1 part of boiling water, insoluble in alcohol and ether; it is much harder than cane-sugar, and but faintly sweet; it is not easily brought into alcoholic fermentation by the action of yeast, but easily undergoes “ lactic fermentation ” when cheese is added. During this process milk-sugar is converted into lactic acid. Milk-sugar resembles grape-sugar in its action on alkaline solution of copper, from which it precipitates cuprous oxide. Starch, Amylum, C6H10O5 = 162. Starch is very widely distributed in the vegetable kingdom, and is found chiefly in the seeds of cereals and leguminosse, but also in the roots, stems, and seeds of nearly all plants. It is prepared from wheat, potatoes, rice, beans, sago, arrow-root, etc., by a mechanical operation. The vegetable matter containing the starch is comminuted by rasping or grinding, in order to open the cells in which it is deposited, and then steeped in water; the softened mass is then rubbed on a sieve under a current of water which washes out the starch, while cellular fibrous matter remains on the sieve; the starch deposits slowly from the washings, and is further purified by treating it with water. 350 CONSIDERATION OF CARBON COMPOUNDS. Starch forms white, amorphous, tasteless masses, which are pecu- liarly slippery to the touch, and easily converted into a powder; it is insoluble in cold water, alcohol, and ether; when boiled with water, it yields a white jelly (mucilage of starch, starch-paste) which cannot be looked upon as a true solution, but is a suspension of the swollen starch particles in water; by continued boiling with much water some starch passes into solution. Starch, when examined under the microscope, is seen to consist of granules differing in size, shape, and appearance, according to the plant from which the starch was obtained. Concentric layers, which are more or less characteristic of starch-granules, show that they are formed in the plant by a gradual deposition of starch matter. The most characteristic test for starch is the dark-blue color which iodine imparts to it (or better to the mucilage). This color is due to the formation of iodized starch, Amylum iodatum, U. S. P., an unstable dark-blue compound of the doubtful composition C6H9I05I. Starch is an important article of food, especially when associated, as in ordinary flour, with albuminous substances. Dextrin, CcH1305 (British gum). Obtained by boiling starch with diluted acids, or by subjecting starch to a dry heat of 175° C. (347° F.) or by the action of diastase (infusion of malt) upon hydrated starch. Malt is made by steeping barley in water until it germinates and then drying it. Dextrin is a colorless or slightly yellowish, amorphous powder, resembling gum-arabic in some respects; it is soluble in water, reduces alkaline copper solutions, and is colored light w in e-red by iodine. It is extensively used in mucilage as a substitute for gum-arabic. Gums. These are amorphous substances of vegetable origin, soluble in water or swelling up in it, forming thick, sticky masses; they are insoluble in alcohol, and are converted into glucose by boiling with dilute sulphuric acid. Gum-arabic consists chiefly of the calcium salt of arabic acid, C6H10O5.H2O. Other gums occur in the cherry tree, in linseed or flaxseed, in Irish moss, in marsh-mallow root, etc. Cellulose, C6H10O5, perhaps C18H30O15 (Plant fibre, Lignine). Cellulose constitutes the fundamental material of which the cellular membrane of vegetables is built up, and forms, therefore, the largest portion of the solid parts of every plant; it is well adapted to this purpose on CARBOHYDRATES. 351 account of its insolubility in water and most other solvents, its resist- ance to either alkaline or acid liquids, and its tough and flexible nature. Some parts of vegetables (cotton, hemp, and flax, for instance) are nearly pure cellulose. Pure cellulose is a white, translucent mass, insoluble in all the common solvents, but soluble in an ammoniacal solution of basic cupric carbonate; it is not colored blue by iodine. Treated with concentrated sulphuric acid it swells up, and gradually dissolves; water precipitates from such solutions a substance known as amyloid, which is an altered cellulose giving a blue color with iodine. Upon diluting the sulphuric acid solution with water and boiling it, the cellulose is gradually converted into dextrin and dextrose. Unsized paper (which is chiefly cellulose) dipped into a mixture of two volumes of sulphuric acid and one volume of water, forms, after being washed and dried, the so-called “parchment paper,” which possesses all the valuable properties of parchment. Pyroxylin, Pyroxylinum, C6H8(N02)205 (.Dinitro-cellulose, Soluble gun- cotton). By the action of nitric acid of various strengths on cellulose, three different substitution products may be obtained, which are dis- tinguished as mono-, di-, and triuitro-cellulose : C6H10O5 + HN03 = C6H9(N02)05 + h2o C6H10O5 + 2HN0S = C6H8(N02)205 + 2H20 C6H10O5 + 3HN03 = C6H7(N02)305 + 3H20. The trinitro cellulose is the highly explosive gun-cotton; an intimate mixture of gun-cotton and camphor is now extensively used under the name of celluloid. The dinitro-cellulose or pyroxylin is soluble in a mixture of ether and alcohol; this solution is known as collodion. Neither the mono- nor trinitro-cellulose is soluble in a mixture of ether aud alcohol. Experiment 54. Immerse 2 grammes of dry cotton for ten hours in a pre- viously cooled mixture of 20 grammes of nitric acid and 24 grammes of sul- phuric acid. Wash the pyroxylin thus obtained with cold water until the washings have no longer an acid reaction. Dissolve 1 part of the dry pyroxylin in a mixture of 18 parts of ether and 6 parts of alcohol. The solution obtained is collodion. Glycogen, C6H10O5. Found exclusively in animals; it occurs in the liver, the white blood-corpuscles, in many embryonic tissues, and in muscular tissue. Pure glycogen is a white, starch-like, amorphous 352 CONSIDERATION OF CARBON COMPOUNDS substance, soluble in water, insoluble in alcohol; by the action of dilute acids it is converted into glucose. Glucosides. This term is applied to a group of substances (chiefly of vegetable origin) which, by the action of acids, alkalies, or fer- ments, suffer decomposition in such a manner that one of the products formed is grape-sugar. Glucosides may, therefore, be looked upon as compound sugars, or sugar in combination with various other sub- stances. The following is a list of the more important glucosides, giving also their composition and the sources whence they are obtained : Amygdalin, c2„h27nou Bitter almonds, etc. Arbutin, Arbutus uva ursi. Cathartic acid, c180h192n4so82 ? Senna. Carminic acid, ? Cochineal. Colocynthin, Colocynthis. Digitalin, Digitalis. Gentiopicrin, Root of gentiana. Glycyrrhizin, Liquorice root. Helleborin, Root of hellebore. Indican, ? Indigo plant. Jalapin, CsiH^Oj, Jalap resin. Myronic acid, Seeds of black mustard. Salicin, Bark of willow. Scam monin, Resin scammony. Solanin, ? Various specimens of solanum. Tannins, ChH10O9 In many barks, leaves, etc. Digitalin. The leaves of digitalis purpurea contain a number of glucosides, mixtures of which in varying proportions form the offici- nal article sold under above name. Digitonin is an amorphous, yel- lowish substance, soluble in alcohol. Digitalein is a white, intensely bitter, amorphous substance. Digitoxin is a colorless, crystalline solid; it is the most poisonous of the constituents of digitalin and is found in the leaves only to the extent of 0.01 to 0.02 per cent.; it is not a glucoside. Digitalin, (C5H802).r, is a white, amorphous powder, solu- ble at ordinary temperature in about 1000 parts of water and in about 100 parts of alcohol of 50 .per cent. It is soluble in concentrated hydrochloric acid, forming a golden yellow solution. A similar yellow solution is obtained by dissolving it in concentrated sulphuric acid, the color gradually changing to blood-red. The yellow color of the sulphuric acid solution changes to a beautiful violet on the addition of a drop of nitric acid or ferric chloride. AMINES AND AMIDES. CYANOGEN COMPOUNDS. 353 Myronic acid, C10Hl9NS2O10, is found as the potassium salt, which is known as sinigrin, in black mustard seed. When treated with solution of myrosin, a substance also contained in mustard seed and acting as a ferment upon myronic acid or its salts, potassium myro- nate is converted into dextrose, allyl mustard oil, and potassium bisulphate : kc10h18ns2o10 = c6h12o6 + c3h5ncs + khso4. Potassium myronate. . Dextrose. Allyl mustard oil. Potassium bisulphate. The univalent radical allyl, C3H5i, is isomeric, but not identical with the trivalent radical glyceryl, C3H5 The triatomic alcohol glycerin, C3II5(OH)3, may, liowever, be converted into the monatomic allyl alcohol C3H5OH, by various processes. From allyl alcohol an artificial allyl mustard oil is manufactured. Allyl sulphide, (C3H5)2S, is the chief constituent of the oil of garlic. 47. AMINES AND AMIDES. CYANOGEN COMPOUNDS. Forms of nitrogen in organic compounds. Nitrogen may be present in organic compounds in three forms, viz., ammonia, cyanogen, nitric acid, or derivatives of these compounds. Substances containing nitrogen in the nitric acid form may be either organic salts of this acid (nitrates), or may have been formed by replacement of hydrogen atoms by the nitric acid radical NC)2. These latter compounds, termed nitro-compounds, such as nitro-cellulose, nitro-benzene, etc., do not occur in nature, but are obtained exclusively by artificial means, generally by treatment of the organic substance with concen- trated nitric acid; all these nitro-compounds are more or less explo- sive. Questions.—451. To which group of substances is the term “ carbohy- drates’’ applied? 452. State the general properties of carbohydrates. 453. Mention the three groups of carbohydrates, and the composition and charac- teristics of the members of each group. 454. Mention some fruits in which grape-sugar, and some plants in which cane-sugar is found. 455. What is the difference between grape-sugar and cane-sugar, and by what tests can they be distinguished? 456. From what source, and by what process, is milk-sugar obtained ? 457. What is starch, what are its properties, by what tests can it be recognized, and what substance is formed when diastase or dilute acids act upon it? 458. Where is cellulose found in nature, and what are its proper- ties? 459. What three compounds may be obtained by the action of nitric acid upon cellulose, and what are they used for? 460. What substances are termed glucosides? Mention some of the more important glucosides. 354 CONSIDERATION OF CARBON COMPOUNDS. Cyanogen compounds contain nitrogen in the form of cyanogen, CN, a radical the compounds of which will be considered hereafter. Organic compounds containing nitrogen in the ammonia form are known as amines or amides, organic bases or alkaloids. (Albuminous substances also contain nitrogen in the ammonia form.) Amines. Whenever the hydrogen of ammonia is replaced by alco- holic radicals (or hydrocarbon residues) compounds are formed which are termed amines. For instance : /H /C2H5 /C2H5 /C2H5 yC H3 NfH, NfH , Nr-C2H5, NfC2H5, N—CSH6. \h \h \h \ch xch Ammonia. NH3, N(C2H5)H2, N(C2H5)2H, N(C2h#)„ nch3.c2h5 c4h9. Ethylamine. Diethylamine. Triethyiamine. Methylethyl-butylamine. Amines resemble ammonia in their chemical properties; they are, like ammonia, basic substances; they combine with acids directly and without elimination of water, thus : NH3 + HC1 = NII4C1; N(C2H5)3 + HC1 = N(C2H6)3HC1 Triethylamine. from ammonia by replacement of Thus: Triethylamine chloride. Amides are substances derived hydrogen atoms by acid radicals. /H N^-H, \H Ammonia. /C2h3o n(h , \h Acetamide. /C2H,0 n(-c2h3o, \h Diacetamide. n=h2. 22 Carbamide or urea. Amides also resemble ammonia in their chemical properties; to a less extent, however, than amines, because the acid radicals have a tendency to neutralize the basic properties of ammonia : Formamide, N(CHO)H2, is a colorless liquid, obtained by heating ethyl formate with an alcoholic solution of ammonia. This compound is of interest because it combines with chloral, forming Chloralformavii.de (Chloralamide), N(CH0)H2.C2HC130, a substance recently used as a hypnotic. It is a color- less, odorless, crystalline substance, having a faintly bitter taste. It is soluble in 20 parts of cold water and in 1.5 parts of alcohol. By heating the aqueous solution to 60° C. (140° F.) it is decomposed into chloral and formamide or ammonium formate. Caustic alkalies liberate iodoform and ammonia. Amido-acids are acids in which hydrogen has been replaced by NH2. Thus, amido-acetic acid, also known as glycocoll or glycine, is represented by the formula C2H3(NH2)02 or ; it is a sub- stance which has both acid and basic properties, and is a product of AMINES AND AMIDES. CYANOGEN COMPOUNDS. 355 the decomposition of either glycocholic or hippuric acid by hydro- chloric acid. Amido-formic acid or carbamic acid, CII.NH2.02, is the acid which, in the form of the ammonium salt, is a constituent of the commercial ammonium carbonate. It is formed by the direct action of carbon dioxide upon ammonia : C02 + 2NH3 = c.nh4.nh2.o2. Formation of amines and amides. These substances are found as products of animal life (urea), of vegetable life (alkaloids), of destruc- tive distillation (aniline, pyridine), of putrefaction (ptomaines), and may also be produced synthetically—for instance, by the action of ammonia upon the chloride or iodide of an alcohol or acid radical: C2H5.I + NH3 = HI + NH2C2H5. Ethyl iodide. Ammonia. Hydriodic acid. Ethylamine. C2H30.C1 + 2NH3 = nh4ci + nh2.c2h3o. Acetyl chloride. Ammonia. Ammonium chloride. Acetamide, Amines may also be formed by the action of nascent hydrogen upon the cyanides of the alcoholic radicals : Methyl cyanide. CH3ON + 4H = NH2.C2H5. Ethylamine. Amines may in some cases be formed by the action of nascent hydrogen upon nitro-compounds; the manufacture of aniline depends on this decomposition : C6H5N02 + 6H = 2H20 + NH2.C6H5. Nitro-benzene. Hydrogen. Water. Phenylamine, or aniline. Occurrence of organic bases in nature. The various organic basic substances found in nature are either amines (compounds containing carbon, hydrogen, and nitrogen only), or amides (compounds contain- ing, besides the three elements named, also oxygen). But a small number of organic bases is found in the animal system, urea being the most important one. In plants organic bases are more frequently met with, and are grouped together under the name of alkaloids. While the constitution of many alkaloids has not yet been sufficiently ex- plained, we known that many of them are derivatives of aromatic compounds, for which reason the consideration of the whole group will be deferred until benzene and its derivatives are spoken of. The large number of basic substances found in putrefying matter and termed 'ptomaines will also be considered later on. 356 CONSIDERATION OF CARBON COMPOUNDS. Cyanogen compounds. Cyanogen itself does not occur in nature, but compounds of it are found in a few plants (amygdalin), and also in some animal fluids (saliva contains sodium sulphocyanate). Gases issuing from volcanoes (or from iron furnaces) sometimes contain cyanogen compounds. The univalent residue cyanogen, — C=N,. or CN, was the first compound radical distinctly proved to exist, and isolated by Gay- Lussac in 1814. The name cyanogen signifies “generating blue,” in allusion to the various blue colors (Prussiau and Turnbull’s blue) containing it. (The symbol Cy, sometimes used in place of CN, has been adopted merely to simplify the writing of formulas of cyanogen compounds). Cyanogen and its compounds show much resemblance to the halo- gens and their compounds, as indicated by the composition of the following substances: C1C1, HC1, KI, HC10, Chlorine, Hydrochloric acid. Potassium iodide. Hypochlorous acid. Cyanogen. CNCN, HBr, KCN, HCNO, Hydrobromic acid. Potassium cyanide. Cyanic acid, CNC1, HCN, AgCN, HCNS, Cyanogen chloride. Hydrocyanic acid. Silver cyanide. Sulphocyanic acid. Dicyanogen, (CN)2. The unsaturated radical CN does not exist as such in a free state, but combines whenever liberated with another CN, forming dicyanogen. It may be obtained by heating mercuric cyanide: Hg(CN)2 = Hg + 2CN. It is a colorless gas, having an odor of bitter almonds, and burning with a purple flame, forming carbon dioxide and nitrogen; it is solu- ble in water, and may be converted into a colorless liquid by pressure ; it acts as a poison, both to animal and vegetable life, even when present in but small proportions in the air. Hydrocyanic acid, HCN = 27 (Oyanhydric acid, Hydrogen cyanide, Prussic add). This compound is found in the water distilled from the disintegrated seeds or leaves of amygdalus, prunus, laurus, etc. It is also found among the products of the destructive distillation of coal, and is formed by a great number of chemical decompositions. For instance: passing ammonia over red-hot charcoal: 4NH3 + 3C = 2(NH4CN) + CH4. Ammonia. Carbon. Ammonium cyanide. Methane. AMINES AND AMIDES. CYANOGEN COMPOUNDS. 357 By the action of ammonia on chloroform CHC13 + SH3 = HCN + 3HC1. Chloroform. Hydrocyanic acid. Hydrochloric acid. By heating ammonium formate to 200° C. (392° F.): NH4CH02 = II CIST + 2H20. Ammonium formate. Hydrocyanic acid. Water, By the action of hydrosulphuric acid upon mercuric cyanide By the decomposition of alkali cyanides by diluted acids : Hg2CN + H2S = HgS + 2HCN. KCN + HC1 = KC1 + HCJST. By the action of hydrochloric acid upou silver cyanide : AgCN + HC1 = AgCl + HCJST. By distilling potassium ferrocyanide with diluted sulphuric acid : 2K4Fe(CN)6 + 6(H2S04) = K2Fe2(CN)6 + fiKHS04 + 6HCN. Potassium ferrocyanide. Sulphuric acid. Potassium ferrous ferrocyanide. Potassium acid sulphate. Hydrocyanic acid. Experiment 55. Place 20 grammes of potassium ferrocyanide and 40 c.c. of water into a boiling-flask of about 200 c.c. capacity; provide the flask with a funnel-tube and connect it with a suitable condenser, the exit of which should dip into 60 c.c. of diluted alcohol, contained in a receiver, which latter should be kept cold by ice during the operation. After having ascertained that all the joints are tight, add through the funnel-tube a previously prepared mixture of 15 grammes of sulphuric acid and 20 c.c. of water. Apply heat and slowly distil until there is little liquid left with the salts remaining in the flask. Determine the strength of the alcoholic solution of hydrocyanic acid thus prepared volumetrically and dilute it with water until it contains exactly two per cent, of HCN. Pure hydrocyanic acid is, at a temperature below 26° C. (78.8° F.), a colorless, mobile liquid, of a penetrating, characteristic odor resem- bling that of bitter almonds; it boils at 26.5° C. (80° F.) and crystal- lizes at —15° C. (5° F.). It is readily soluble in water, and a 2 per cent, solution is the diluted hydrocyanic, acid, Acidum hydrocyanicum dilutum. According to the U. S. P., this diluted acid is made either by the decomposition of potassium ferrocyanide by diluted sulphuric acid in a retort, the delivery-tube of which passes into a receiver containing a mixture of water and alcohol, by which the liberated gas is absorbed, this liquid being afterward diluted with a sufficient quantity of water to make a 2 per cent, solution, or it is made extemporaneously by the decomposition of 6 parts by weight of silver cyanide by 5 parts of 358 CONSIDERATION OF CARBON COMPOUNDS. hydrochloric acid, diluted with 55 parts of water, allowing the silver chloride to subside and pouring off the clear liquid. The diluted acid has the characteristic odor of bitter almonds, a slightly acid reaction, and is completely volatilized by heating. Whilst the pure acid is very readily decomposed by exposure to light, etc., the dilute acid is fairly stable. Potassium cyanide, Potassii cyanidum, KCN = 65 (Cyanide of potas- sium). The pure salt may be obtained by passing hydrocyanic acid into an alcoholic solution of potassium hydroxide. The commercial article, however, is a mixture of potassium cyanide with potassium cyanate. It is obtained by fusing potassium ferrocyanide with potas- sium carbonate in a crucible, when potassium cyanide and cyanate are formed, whilst carbon dioxide escapes, and metallic iron is set free and collects on the bottom of the crucible. The decomposition is as fol- lows : K4Fe(CN)6 + K2C03 == 5KCN + KCNO + Fe + C02. Potassium ferrocyanide. Potassium carbonate. Potassium cyanide. Potassium cyanate. Iron. Carbon dioxide. Potassium cyanide is a white, deliquescent salt, odorless when perfectly dry, but emitting the odor of hydrocyanic acid when moist. Potassium cyanides and other alkali cyanides show a tendency to combine with the cyanides of heavy metals, forming a number of double cyanides, such as the cyanides of sodium and silver, NaCN. AgCN, etc. Silver cyanide, Argenti cyanidum, AgCN = 133.7 (Cyanide of silver). A white powder, obtained by precipitating solution of potassium cyanide with silver nitrate. It is insoluble in water, slightly soluble in water of ammonia; evolves cyanogen when heated. Mercuric cyanide, Hydrargyri cyanidum Hg(CN)2 (Cyanide of mercury). A white, crystalline salt, obtained by dissolving mercuric oxide in hydrocyanic acid; it is soluble in water and alcohol and evolves cyanogen when heated. Analytical reactions for hydrocyanic acid. (Potassium cyanide, KCN, may be used.) 1. Hydrocyanic acid, or soluble cyanides, give with silver nitrate a white precipitate of silver cyanide, which is sparingly soluble in am- AMINES AND AMIDES. CYANOGEN COMPOUNDS. 359 monia, soluble in alkali cyanides or thiosulphates, but insoluble in diluted nitric acid. Concentrated nitric acid dissolves it with decom- position : HCN + AgN03 = AgCN + HN03. 2. Hydrocyanic acid mixed with ammonium hydric sulphide and evaporated to dryness forms sulphocyanic acid, which, upon being slightly acidulated with hydrochloric acid, gives with ferric chloride a blood-red color of ferric sulphocyanate. (Excess of ammonium sulphide must be avoided.) 3. Hydrocyanic acid, or soluble cyanides, give, when mixed with ferrous and ferric salts and potassium hydroxide, a greenish precipi- tate, which, upon being dissolved in hydrochloric acid, forms a pre- cipitate of Prussian blue, Fe4(FeC6N6)3. This reaction depends on the formation of potassium ferrocyanide by the action of the cyanogen upon both the potassium of the potassium hydroxide and the iron of the ferrous salt. In alkaline solutions, the blue precipitate does not form, for which reason hydrochloric acid is added. 4. Hydrocyanic acid heated with dilute solution of picric acid gives a deep-red color on cooling. In cases of poisoning, the matter under examination is distilled (if neces- sary after the addition of water) from a retort connected with a cooler. To the distilled liquid the above tests are applied. If the substance under ex- amination should have an alkaline or neutral reaction, the addition of some sulphuric acid may be necessary in order to liberate the hydrocyanic acid. The objectionable feature to this acidifying is the fact that non-poisonous potassium ferrocyanide might be present, which upon the addition of sulphuric acid would liberate hydrocyanic acid. In cases where the addition of an acid becomes necessary, a preliminary examination should, therefore, decide whether or not ferro- or ferricyanides are present. Antidotes. Hydrocyanic acid is a powerful poison both when inhaled or swallowed in the form of the acid or of soluble cyanides. As an antidote is recommended a mixture of ferrous sulphate and ferric chloride with either sodium carbonate or magnesia. The action of this mixture is explained in the above reaction 3, the object being to convert the soluble cyanide into an insoluble ferrocyanide of iron. In most cases of poisoning by hydrocyanic acid there is, however, no time for the action of such an antidote, in conse- quence of the rapidity of the action of the poison, and the treatment is chiefly directed to the maintenance of respiration by artificial means. Cyanic acid, HCNO, and Sulphocyanic acid, HCNS, are both color- less acid liquids, the salts of which are known as cyanates and sulpho- cyanates. These salts are obtained from alkali cyanides by treating 360 CONSIDERATION OF CARBON COMPOUNDS. them with Oxidizing agents or by boiling their solutions with sulphur, when either oxygen or sulphur is taken up by the alkali cyanide: KCN -f- O = KCNO = Potassium cyanate. KCN -)- S = KCNS = Potassium sulphocyanate. The acids themselves may be liberated from their salts by dilute mineral acids. Sulpliocyanates give with ferric salts a deep-red color, which is not affected by hydrochloric acid, but disappears on the addi- tion of mercuric chloride. Metallocyanides. Cyanogen not only combines with metals to form the true cyanides, which may be looked upon as derivatives of hydro- cyanic acid, but cyanogen also enters into combination with certain metals (chiefly iron), forming a number of complex radicals, which upon combining with hydrogen form acids, or with basic elements form salts. It is a characteristic feature of the compound cyanogen radicals, thus formed, that the analytical characters of the metals (iron, etc.) entering into the radical are completely hidden. Thus, the iron in ferro- or ferricyanides is not precipitated by either alkalies, soluble carbonates, ammonium sulphide, or any of the common reagents for iron, and its presence can only be demonstrated by these reagents after the organic nature of the compound has been destroyed by burning it (or otherwise), when ferric oxide is left, which may be dissolved in hydrochloric acid and tested for in the usual manner. The principal iron-cyanogen radicals are ferrocyanogen [Fe* (CN)6i]iv, and ferricycmogen [Fe2vi(CX)12']vi. These two radicals con- tain iron in the ferrous and ferric state respectively, and form, upon combining with hydrogen, acids which are known as hydroferrocyanic acid, II4Fe(CN)6 (tetrabasic), and hydroferricyanic acid, H6Fe2(CN)12 (hexabasic); the salts of these acids are termed ferrocyanides and ferricyanides. Potassium ferrocyanide, Potassii ferrocyanidum, K4Fe(CN)6.3H20= 421.9 (.Ferrocyanide of potassium, Yellow prussiate of potash). This salt is manufactured on a large scale by heating refuse animal matter (waste leather, horns, hoofs, etc.) with potassium carbonate and iron (filings, etc.). The fused mass is boiled with water, and from the solution thus formed the crystals separate on cooling. The nitrogen and carbon of the organic matter (heated as above stated) combine, forming cyanogen, which enters into combination first with potassium and afterward with iron. AMINES AND AMIDES. CYANOGEN COMPOUNDS. 361 Potassium ferrocyauide forms large, translucent, pale lemon-yellow, soft, odorless, non-poisonous, neutral crystals, easily dissolving in water. A nalytical reactions: 1. Ferrocyanides heated on platinum foil burn and leave a residue of (or containing) ferric oxide. 2. Ferrocyanides heated with concentrated sulphuric acid evolve carbonic oxide; with dilute sulphuric acid liberate hydrocyanic acid; with concentrated hydrochloric acid liberate hydroferrocyanic acid. 3. Soluble ferrocyanides give a blue precipitate with ferric salts (Plate I., 5): 3K4Fe(CN)6 + 2FeaCI6 = 12KC1 + Fe4 (FeC6N6)3 Potassium ferrocyanide. Ferric chloride Potassium chloride. Ferric ferro- cyanide. The blue precipitate of ferric ferrocyanide, or Prussian blue, is in- soluble in wrater and diluted acids, soluble in oxalic acid (blue ink)? and is decomposed by alkalies with separation of brown ferric hydroxide and formation of potassium ferrocyanide. The addition of an acid restores the blue precipitate. 4. Soluble ferrocyanides give with cupric solutions a brownish-red precipitate of cupric ferrocyanide. (Plate III., 5.) 5. Soluble ferrocyanides produce, with solutions of silver, lead, and zinc, white precipitates of the respective ferrocyanides. 6. Ferrocyanides give with ferrous salts a white precipitate of ferrous ferrocyanide, soon turning blue by absorption of oxygen- (Plate I., 4.) Potassium ferricyanide, X6Ee2(CN)12 (Bed prussiate of potash). Obtained by passing chlorine through solution of potassium ferro- cyanide : 2K4Fe(CN)6 + 2C1 = 2KC1 + K6Fe2(CN)12. Potassium ferrocyanide. Chlorine. Potassium chloride. Potassium ferricyanide. While apparently this decomposition consists merely in a removal of two atoms of potassium from two molecules of potassium ferro- cyanide, the change is actually more complete, as the atoms arrange themselves differently, the iron passing also from the ferrous to the ferric state. Potassium ferricyanide crystallizes in red prisms, soluble in water. It forms, with ferrous solutions, a blue precipitate of ferrous ferricy- anide, or Turnbull’s blue: K6Fe2(CN)12 + 3FeS04 = 3K2S04 + Fe3Fe2(CN)12. 362 CONSIDERATION OF CARBON COMPOUNDS. With ferric solutions no precipitate is produced by potassium ferri- cyanide, but the color is changed to a dark greenish-brown. Nitro-cyan-methane, CH2CN.N02 (Fulminic acid). This substance may be looked upon as a derivative of methane, CH4, in which two atoms of hydro- gen are replaced by cyanogen and N02 respectively. It is not known in the separate state, but its combinations with metals are well known, especially mercuric fulminate, which is manufactured and used as an explosive in percus- sion caps, etc. It is made by adding alcohol to a solution of mercury in nitric acid. Silver fulminate can be obtained by a similar process. 48. BENZENE SERIES. AROMATIC COMPOUNDS. General remarks. It has been stated before, that all organic com- pounds may be looked upon as derivatives of either methane, or benzene, C6H6, these derivatives being often spoken of as fatty and aromatic compounds respectively. The term aromatic compounds was given to these substances on accouut of the peculiar and fragrant odor possessed by many, though not by all of them. Benzene and methane derivatives differ considerably in many respects, and, as a general rule, aromatic compounds cannot be converted into fatty com- pounds, or the latter into aromatic compounds, without suffering the most vital decomposition of the molecule, and in most cases this trans- formation cannot be accomplished at all. On the average, aromatic compounds are richer in carbon than fatty compounds, containing of this element at least 6 atoms ; when decom- posed by various methods, aromatic compounds, in many cases, yield benzene as one of the products; most aromatic substances have anti- septic properties, and none of them serves as animal food, although Questions.—461. What are the three chief forms in which nitrogen enters into organic compounds? 462. What are amines and amides; in what re- spects do they resemble ammonia compounds ? 463. What is cyanogen, what is dicyanogen, and how is the latter obtained? 464. How does cyanogen occur in nature, and which non-metallic elements does it resemble in the con- stitution of various compounds? 465. Mention some reactions by which hydrocyanic acid is formed, aud state the two processes by which the officinal diluted acid is obtained. What strength and what properties has this acid? 466. State the composition of pure potassium cyanide and of the commercial article. How is the latter made ? 467. Give reactions for hydrocyanic acid and cyanides. 468. Explain the constitution and give the composition of ferro- and ferricyanides. 469. Give composition, mode of manufacture, and tests of potassium ferrocyanide. 470. What is red prussiate of potash, how is it obtained, and by what reactions can it be distinguished from the yellow prussiate ? BENZENE SERIES. AROMATIC COMPOUNDS. 363 quite a number are taken into the system in small quantities, as, for instance, some essential oils, caffeine, etc. While some aromatic compounds are products of vegetable life, many of them (like benzene itself) are obtained by destructive distil- lation, and are, therefore, contained in coal-tar, from which quite a number are separated by fractional distillation. The constitution of benzene is best explained by assuming that of the 4 X 6 = 24 affinities of the 6 carbon atoms, 18 affinities are lost by uniting the carbon atoms into a closed chain, while but 6 affinities are left unprovided for and may be saturated by other elements or groups of elements. The carbon chain of aromatic compounds and benzene may be graphically represented thus: 1 6\c^c/2 11 i 5'/C\c^'Cn'3 I 4 H I H. XL .H l \ XH I H It has been found that whenever one atom or one radical replaces hydrogen in benzene, the product obtained is the same, no matter by what method the change was brought about. Thus we know but one mono-brom-benzene, C6H5Br, one nitro-benzene, C6H5N02, etc. It is different when two or more atoms or radicals (of the same kiud) replace hydrogen in benzene, since it has been found that in this case often isomeric compounds are formed. For instance, we known three different substances which have been obtained by replacement of two hydrogen atoms in benzene by two hydroxyl groups. This would indicate that it makes a difference, as far as the properties of a compound are concerned, in which relative position the introduced radicals stand to one another, and while we have no proof whatever in regard to this position, yet we often repre- sent it graphically, as, for instance, in the following three cases, where the two groups OH replace hydrogen in different positions : OH h\c/c\c/oh (4 A H/ \ \h I II Ortho-position. 1.2. OH H\cAc/H i A H/ \qS \0H I H Meta-position. 1.3. OH II I I OH Para-position. 1.4. 364 CONSIDERATION OF CARBON COMPOUNDS. Designating the hydrogen atoms in benzene with numbers, thus : 1 2 3 4 5 6 C6 H H H H H H, the above 3 compounds show that in one case the hydrogen atoms 1 and 2, in the second 1 and 3, in the third 1 and 4 have been replaced by OH. The compounds formed in this way are distinguished as ortho-, meta-, and para-compounds. The molecular formula of the above three compounds is C6H602, apparently indicating benzene in combination with two atoms of oxygen or dioxybenzene; actually they are dihydroxy benzene. Ortho-di- 1 2 hydroxy benzene, C6H4OHOH, is known as pyro-catecliin, meta dx- 1 3 hydroxy benzene, C6H4OHOH, as resorcin, and para-di-hydroxy 1 4 benzene, C6II4OHOH, as hydroquinone. Benzene series of hydrocarbons. By replacing the hydrogen atoms in benzene by methyl, CH3, a series of hydrocarbons is formed having the general composition CnH2n_g To this benzene series belong : Benzene c6 h6 - B. P. 80° C. Toluene c7hs = c6h5ch3 110 Xylene H10 = C6H4(CH3)2 142 Cumene C9H12 = C6H3(CH3)s 151 Cymene C10HU = (C6H2(CH3)4? 175 Penta-methyl-benzene CuH16 = C6H(CH3)5 188 Hexa-methyl-benzene C12H18 = C6(CH3)6 202 This first four members of this series are found in coal-tar ; the fifth member, cymene, C10H14, occurs in the oil of thyme ; the last two have been obtained by synthetical processes. While but one toluene is known, the higher members form quite a number of isomeric com- pounds. Cymene, found in the oil of thyme, is, for instance, not the tetra-methy 1-benzene, but the para-methyl-propyl-benzene, C6H4.CH3. C3H7. This compound is of interest on account of its close relation to the terpenes and camphors, which will be spoken of later. Benzene, C6H6 [Benzol). When coal-tar is distilled, products are obtained which are either lighter or heavier than water, and by col- lecting the distillate in water a separation into so-called light oil (floating on the water) and heavy oil (sinking beneath the water) is accomplished. Benzene is found in the light oil and obtained from it by distillation after phenol has been removed by treatment with caustic soda and some basic substances by means of sulphuric acid. Pare BENZENE SERIES. AROMATIC COMPOUNDS 365 benzene may be obtained by heating benzoic acid with calcium hydroxide : C6H5.C02H + Ca(OH)2 = CaC03 + H20 + C6H6. Experiment 56. Mix 25 grammes of benzoic acid with 40 grammes of slaked lime and distil from a dry flask, connected with a condenser. Add to the dis- tilled fluid a little chloride of calcium and redistil from a small flask. The product obtained is pure benzene. Notice that it solidifies when placed in a freezing mixture of ice and common salt. Observe the analogy between Ex- periments 56 and 40. In one case a fatty acid is decomposed by an alkali with liberation of methane, in the other an aromatic acid with liberation of benzene, the carbonate of the decomposing hydroxide being formed in both cases. Pure benzene is a colorless, highly volatile liquid, having a peculiar, pleasant odor and a specific gravity of 0.884 ; it boils at 80.5° C. (177° F.) and solidifies at 0° C. (32° F.); it is an excellent solvent for fats, oils, resins, and many other organic substances. Nitro-benzene, C6H5.N02. When benzene is treated with concen- trated nitric acid, or with a mixture of nitric and sulphuric acids, nitro-benzene is formed : c6h6 + hno3 = c6h5no2 + h2o. Experiment 57. Mix 50 c.c. sulphuric acid with 25 c.c. nitric acid; allow to cool, place the vessel containing the mixture in water, and add gradually 5 c.c. of benzene, waiting after the addition of a few drops each time until the reac- tion is over. Shake well until all benzene is dissolved and pour the liquid into 300 c.c. of water. The yellow oil which sinks to the bottom is nitro-benzene. It may be purified by washing with water and redistilling, after removal of water and shaking with calcium chloride. Nitro-benzene is an almost colorless or yellowish oily liquid, which is insoluble in water, has a specific gravity of 1.2, a boiling-point of 205° C. (401° F.), a sweetish taste, highly poisonons properties, even when inhaled, and an odor resembling that of oil of bitter almond, for which it is substituted under the name of essence of mirbane. It is manufactured on a large scale and used chiefly in the preparation of aniline, for which see Index. Benzene-derivatives. The relation existing between methane- and benzene-derivatives may be shown by comparing the composition of a few derivatives : . Methane, ch4 Benzene, c6h6 Methyl, ch3 Benzyl, Phenyl, }c6h5 Ethane, Toluene, }c6h5.ch3 Methyl-methane, }ch3.ch3 Methyl-benzene, 366 CONSIDERATION OF CARBON COMPOUNDS Methyl-hydroxide, 1 Methyl alcohol, J - CHgOH Phenyl-hydroxide, Phenol, }c6h5.oh Glycerin, /OH C3H5^OH 'OH Pyrogallol, /OH c6h3^oh 'OH Acetic acid, ch3.co2h Benzoic acid, C6H3.C02H Acetic aldehyde, CH3.COH Benzoic aldehyde, c6h5.coh Ethyl-sulphonic acid, so su2\OH Benzene-sulphonic acid, so radicals of benzoic and acetic acid respectively, thus,N—C2H302. Hay, \H and especially aromatic herbs, contain benzoic acid, or compounds having a similar composition, and a portion of these compounds is eliminated in hippuric acid. Administration of benzoic acid increases the amount of hippuric acid in urine. When pure, hippuric acid crystallizes in transparent, colorless, odorless prisms, which have a bitter taste, and are sparingly soluble in water. EXAMINATION OF NORMAL AND ABNORMAL URINE. 449 Analytically, hippuric acid is characterized— 1. By giving a sublimate of benzoic acid, and an odor of hydro- cyanic acid, when heated in a dry test-tube. 2. By giving a brown precipitate with ferric chloride. 3. By giving off benzene and ammonia when heated with calcium hydroxide. 4. By evolving an intense odor of nitro-benzene when evaporated to dryness with a few drops of nitric acid. Other organic substances, such as kreatin, kreatinin, xanthin, lactic acid, mucus, coloring matters, etc., occur in such small quantities in normal urine that their detection, separation, and quantitative estima- tion are very difficult, and are almost exclusively attempted during scientific investigations. 56. EXAMINATION OF NORMAL AND ABNORMAL URINE. Points to be considered in the analysis of urine. They are : 1. Color, odor, general appearance—whether clear, smoky, cloudy, turbid, etc. 2. Reaction—whether acid, neutral, or alkaline to test-paper. 3. Specific gravity. 4. Total amount of organic and inorganic solids. 5. Total amount of inorganic matter (ash). 6. Determination of urea. 7. Determination of uric acid. 8. Determination of inorganic acids and bases. (Hydrochloric, sul- phuric, and phosphoric acids; sodium, potassium, calcium, magnesium, and iron.) 9. Determination of albumiu. 10. Determination of sugar. Questions.—541. What is urine, where and by what process is it formed in the animal body, and what is its function ? 542. Mention the general physi- cal and chemical properties of urine. 543. Give the average composition of human urine, and state by what conditions the composition is influenced. 544. State the composition and properties of urea. 545. By what process is urea formed in the animal body, and how can it be obtained artificially? 346. Describe a process by which urea may be estimated quantitatively in urine. 547. In what forms is uric acid found in urine, and what are its properties ? 548. Describe the murexid test. 549. How can uric acid be determined quantitatively in urine? 550. What is hippuric acid, and by what tests may it be recognized ? 450 PHYSIOLOGICAL CHEMISTRY. 11. Examination for bile. 12. Examination of any organic or inorganic sediment, either by chemical means or by the microscope. Samples of urine should always be drawn from the well-mixed and exactly measured quantity of the total urine discharged in twenty-four hours. Color. Normal urine is generally pale-yellow or reddish-yellow, but it may be as colorless as water, or as dark brownish-black as porter; a reddish and smoky tint generally indicates the presence of blood, and a brownish-green suggests the presence of the coloring matter of bile. (Plate VII., 1-3.) The true nature of the normal coloring matters of urine is as yet doubtful; the existence of at least two has, however, been demon- strated ; they have been named urobilin and indican or uroxanthin, and are, most likely, products of the decomposition of biliary matters. Indican, Potassium indoxylsulphate, K.C8H6N.S04. Normal urine contains but very little of this coloring matter, but its quantity is increased in a number of diseases. The most convenient test is made by mixing equal volumes of urine and hydrochloric acid in a test-tube and adding drop by drop a filtered solution of bleaching-powder until a green, violet, or blue color is noticed. Normal urine shows a green color only, while a violet-red or intense blue color indicates the pres- ence of indican. By shaking the contents of the test-tube with a little chloroform, the indican is absorbed by the latter, imparting to it a dis- tinct blue color. If the urine should contain albumin, this must be separated before applying the test. Abnormal coloring matters are chiefly those of blood, bile, and of certain vegetables; thus, rhubarb and senna leaves cause a reddish- yellow to deep-red color, especially in alkaline urine ; santonin pro- duces a bright-yellow color, changing to red or crimson on the addition of an alkali. Carbolic acid introduced into the system causes a dark, or even black discoloration of urine, while large doses of salicylic acid color it green. The coloring matters of blood may be recognized by adding to a few drops of urine a drop of freshly prepared tincture of guaiaciun, and agitating with a solution of ozonized ether (ethereal solution of hydro- gen dioxide); the latter is colored blue in case haemoglobin is present. In place of ozonized ether, oil of turpentine which has been in contact with atmospheric air for some weeks may be used, and the test made by allowing urine to flow down the test-tube containing a mixture of URINE. IPl-Atie -rTxT. 1 Urine Tints—Pale yellow, light yellow, yellow. 2 Urine Tints—Reddish yellow, yellowish red, red. 3 Urine Tints— Brownish red. red- dish brown, brownish black. Murexicl test for uric acid. 4 5 Trmnmer’a or Fehliny’a test for sugar. liotger’s bismuth test for sugar 6 7 flmelUt’s test for biliary coloring matters. Pettenkofer’s test for biliary acids. AIbuniinous substances show nearly the same reaction. 8 EXAMINATION OF NORMAL AND ABNORMAL URINE. 451 the oil and tincture; a blue coloration will slowly appear, if blood be present. The test has the serious disadvantage that protoplasm in almost any form will give the blue color. Detection of biliary coloring matter will be considered later. Odor. The normal odor of fresh urine is characteristic, and is some- times spoken of as aromatic; it is not known by what substance or substances this odor is caused. The ammoniacal and putrescent odor which urine acquires on standing, is due to the products of decomposi- tion formed, chiefly ammonia. A number of substances taken internally and separated by the kidneys from the blood, cause the urine to assume a characteristic odor; aromatic substances especially impart such odors; oil of turpen- tine gives an odor reminding of violets, and the odor of cubebs, copaiba, asparagus, garlic, valerian, and other substances is promptly transferred to the urine of persons using these drugs internally. A sweetish smell sometimes attends the presence of large quantities of sugar in urine. Reaction. This is generally acid in healthy urine which has been recently passed, but may become neutral or alkaline within a short period, by decomposition of urea and formation of ammonium carbon- ate. The acid reaction of urine is due chiefly to monosodium ortho- phosphate, HaH2P04, and perhaps also to the acid salts of uric acid. The acidity may be determined voiumetrically by the addition of deci-nor- mal solution of sodium or potassium hydroxide to 100 c.c. of urine. The acidity of urine is generally expressed as oxalic acid, of which 1 c.c. of nor- mal potash solution neutralizes 0.0063 gramme. If, for instance, 100 c.c. of urine require 15 c.c. of deci-normal potash solution, then the acidity of the 100 c.c. urine is 15 X 0.0063 = 0.0945; and for the total urine of the 24 hours—say 1800 c.c.—the acidity expressed as oxalic acid is, therefore, equal to 1.701 grammes. While urine shows an acid reaction generally, it may have a neutral or even alkaline reaction. In many cases this alkaline reaction points to decomposition of urea in the bladder, but it may be due also to the elimination of alkali carbonates, derived from food taken or drugs administered. Thus, the alkali tartrates, citrates, acetates, etc., have (after diges- tion) a tendency to neutralize uric acid, and an excess of them is eliminated as carbonate. 452 PHYSIOLOGICAL CHEMISTRY. To distinguish between the harmless alkaline reaction caused by fixed alkalies and the alkaline reaction produced by decomposition of urea, a piece of red litmus-paper may be used. If this, after having been moistened with the urine, remains blue on drying (by warming gently) the reaction is due to the fixed alkalies; if the red color reap- pears, the alkaline effect is due to ammonia compounds. Urine sometimes is amphoteric in its reaction, i. e., it colors red litmus-paper faintly blue, and blue litmus-paper slightly red. This condition is caused most likely by the simultaneous presence of monosbdium orthophosphate, NaH2P04,which has an acid, and of disodium orthophosphate, Na2HP04, which has an alkaline reaction. Experiment 71. Prepare normal soda solu- tion as directed on page 265, dilute it with 9 parts of water, and titrate with this deci-nor- mal solution 100 c.c. of urine, using litmus- paper as an indicator. Specific gravity. The normal specific gravity of an average amount of 1500 c.c. of urine passed in twenty-four hours is about 1.020, but it varies, even in health, from 1.012 to 1.030 or more. A specific gravity above 1.030 may indicate the pres- ence of sugar, larger quantities of which may cause the specific gravity to rise to 1.050. Albuminous urine is frequently of low specific gravity, 1.010 to 1.012. It should be remembered that the spe- cific gravity of urine considered separately from the quantity of urine passed in twenty-four hours is of no value, and that in some diseases (for instance in acute nephritis with albuminuria) the specific gravity of albuminous urine may be as high as 1.030, while a diabetic urine may have a specific gravity of 1.025, or less, in consequence of a large volume passed. The determination of the specific gravity of urine is generally accomplished by the urinometer, which is a small hydrometer indicat- ing specific gravity from zero (or 1000) to 60 (or 1060). As the tern- Fig. 42 Urinometer. EXAMINATION OF NORMAL AND ABNORMAL URINE. 453 perature influences the density of liquids, a urinometer can only give correct results at a certain degree of temperature, which is generally marked upon the instrument. Some of the urinometers manufactured and sold show incorrect gravities, even at the normal (or stated) temperature, and for this reason a urinometer should always be thoroughly tested before placing full confidence in the results obtained by it. (See Fig. 42.) Determination of total solids. An approximate determination of total solids may be deduced from the specific gravity of the urine, as it has been found that the last two figures of the specific gravity of urine, multiplied by 2.33, correspond to the number of grammes in 1000 c.c. of urine. If, for instance, 1450 c.c. of urine, of a specific gravity of 1.018, have been discharged in twenty-four hours, then the quantity of total solids in 1000 c.c. will be 18 X 2.33, or 41.94 grammes; and in 1450 c.c. 60.81 grammes. A more exact method of determining the total solids in urine is the evaporation of about 10 c.c. in a weighed platinum dish over a water- bath (or, better, under the receiver of an air-pump over sulphuric acid), until it is found that no more loss in weight ensues on continued exposure of the dish in the drying apparatus. By now reweighing the dish, plus contents, and deducting from the weight that of the empty dish, the weight of total solids is found. Determination of inorganic constituents. The platinum dish con- taining the known quantity of total solids is exposed to the action of a non-luminous flame, and the heat continued until all organic matter has been destroyed and expelled. By re weighing now, and deducting the weight of the platinum dish, plus ash, from the weight of the dish, plus total solids, the quantity of total organic matter is determined; and by deducting weight of dish from weight of dish plus ash, the total quantity of inorganic matter is found. Experiment 72. Determine total solids, water, total organic and inorganic matters in a specimen of urine by following the directions given above. Use 10 or 20 c.c. of urine for the analysis. The analysis of the ash is effected by the methods given in con- nection with the consideration of the various acid and basic constituents themselves. Chlorine is determined by precipitating the solution of the ash in nitric acid with silver nitrate, sulphuric acid by barium chloride, phosphoric acid by ammonium molybdate, calcium by ammo- 454 PHYSIOLOGICAL CHEMISTRY. ilium oxalate, potassium by platinic chloride, iron by potassium ferro- cyanide, etc. For the determination of many of the inorganic constituents, it is not necessary to destroy the organic matter as described above, but this determination can be effected directly. Thus, chlorine may be pre- cipitated directly from urine (slightly acidulated with nitric acid) by silver nitrate ; the precipitated silver chloride is collected upon a small filter, well washed, dried, and weighed in a porcelain crucible, after the filter (to which particles of silver chloride adhere) has been burned separately and its ash added to the contents of the crucible, which is moderately heated before weighing. Experiment 73. Prepare deci-normal solution of silver nitrate as directed on page 271, and use it for the quantitative estimation of chlorine as follows: Dilute 10 c.c. of urine with about 100 c.c. of water, neutralize, add a few drops of potassium chromate solution and then the silver solution from a Gay-Lussac burette until the appearance of a permanent reddish color indicates the end of the reaction. Each c.c. of silver solution used represents 0.00354 gramme of chlorine or 0.00584 gramme of sodium chloride. If the urine should be highly colored, destroy the coloring matter by boiling the 10 c.c. of urine previous to titration with a dilute solution of permanganate of potassium, which add drop by drop until a slight rose tint becomes permanent. Now add a trace of oxalic acid to destroy this excess of permanganate, filter off the precipitate of organic matter and oxide of manganese and use the clear neutralized solution for titration. Phosphoric acid is found in urine, in part (about two-thirds) com- bined with alkalies, and in part (about one-third) with lime and magnesia. These phosphates have in acid or neutral urine the com- position NaH2P04, Na2HP04, CaHP04, CaH4(P04)2, MgHPO,; in alkaline urine compounds of the composition Na3P04, Ca3(P04)2, Mg3(P04)2, MgNH4P04 may be present. By adding any alkali the phosphates of calcium and magnesium (generally termed earthy phosphates) are precipitated, the phosphates of sodium or possibly potassium remain dissolved. The so-called earthy phosphates (phosphates of calcium and magne- sium) may be approximately determined by adding a few drops of an alkaline hydroxide to about 50 c.c. of urine, heating to the boiling- point, collecting on a filter, washing, igniting, and weighing in a platinum crucible. Experiment 74. Add to 50 c.c. of urine a few drops of calcium chloride solu- tion and then water of ammonia. Phosphoric acid is completely precipitated, chiefly as tricalcium phosphate, Ca32P04, containing, however, a very small quantity of magnesium ammonium phosphate. Collect the precipitate on a 455 EXAMINATION OF NORMAL AND ABNORMAL URINE. filter, wash well, dry, ignite and weigh it. Calculate the phosphoric acid from the tricalcium phosphate, without reference to the small amount of magnesium phosphate. Experiment 75. Add to 100 c.c. clear urine 5 c.c. hydrochloric acid; boil, and then add barium chloride to complete precipitation. Set aside for one hour, filter, wash well, dry, ignite and weigh. Calculate from the weight of barium sulphate, thus obtained, the percentage of sulphuric acid present in the urine examined. The methods for estimating urea and uric acid have been described in the preceding chapter. Detection of albumin. Serum-albumin and serum-globulin are the forms most frequently present in urine, but peptones and other albu- minoids are also met with. The different methods by which the presence of albumin in urine is demonstrated are based upon the coagu- lation of albumin. This coagulation may be accomplished by heat, by nitric acid, by picric acid, by potassium ferrocyanide in the presence of acetic acid, and also by either metaphosphoric or trichloracetic acid. The urine used for any of these tests must be perfectly clear; if it be not clear, it must be rendered so by processes which vary according to the nature of the substance causing the turbidity. In most cases filtration through good filter-paper may be sufficient; but if this does not accomplish the desired result, it may become necessary to use other means. Thus, if earthy, amorphous phosphates be present (which, especially in alkaline urine, are apt to pass through the best filter- paper), they may be removed by adding to the urine about a fourth part of potassium hydroxide solution, warming the mixture, and filtering. If the turbidity be caused by urates, the urine will generally become clear by passing the test-tube once or twice through a flame. The clear urine is then tested by either (or all) of the following methods : a. Coagulation by heat. A test-tube is filled about one-half with the urine, to which, if not distinctly acid to test-paper, a few drops of acetic acid are added.1 (In case potassium hydroxide has been added in order to precipitate the phosphates, enough of acetic acid must be added to cause a distinct acid reaction.) The test-tube is then held over the flame in such a manner that the heat acts upon the upper half of the urine only, heating this portion gradually to near the boil- ing-point. By thus operating, two strata of fluid are obtained for 1 If acetic acid alone causes a precipitate, this may be due to mucin, which should be fdtered off before heating. 456 PHYSIOLOGICAL CHEMISTRY. comparison, and by holding the test-tube against the light, or against a black background, any difference in the appearance of the upper and lower strata may be noticed. Any cloudiness or opacity seen may be due to albumin, but may also be caused by earthy phosphates, because these are often precipitated by heating, and have been mistaken for albumin quite frequently. The reason why phosphates are often precipitated by heating of urine is this: Urine, showing a slightly acid or neutral reaction, contains dicalcium and di- magnesium phosphates, which salts upon heating are decomposed into soluble monophosphates and insoluble triphosphates, thus : 4CaHP04 = Ca(H2P04)2 + Ca8(P04)2. Dicalcium phosphate. Monocalcium phosphate. Tricalcium phosphate. To decide the nature of the precipitate a few drops (10 to 15) of nitric acid are allowed to flow gently down the side of the tube into the urine. The precipitate will readily disappear when caused by phosphates, but will be permanent when albumin is present. Instead of heating, as above described, merely the upper half of the urine, the total quantity of the urine (acidulated by a few drops of acetic acid) may be heated, and the test-tube set aside for several hours (after having added 10 to 15 drops of nitric acid), in order to allow the albumin to subside, when it can be more distinctly seen and its quantity noticed. b. Nitric acid test. A test-tube is filled to the depth of about half an inch with colorless nitric acid, and an equal quantity of urine is allowed to flow down the side of the test-tube in such a manner that the specifically lighter urine forms a distinct and separate layer over the nitric acid. (If the urine be allowed to flow from a pipette, as shown in Fig. 43, the formation of the two strata is easily accom- plished.) In case albumin is present, a white band or zone of varying thickness (according to the quantity of albumin present) appears at the point of contact. If the urine be highly concentrated, a similar white zone is formed between the acid and urine, due to the separation of insoluble acid urates; the difference between the separated urates and albumin is that the latter forms a sharply defined zone, whilst the urates diffuse into the urine above. Moreover, the urates dissolve on the application of heat. Finally, the separation of acid urates may be avoided by dilut- ing the urine with an equal volume of water and placing this diluted urine upon the nitric acid. c. Picric acid test for albumin. This test has the advantage that EXAMINATION OF NORMAL AND ABNORMAL URINE. 457 neither phosphates nor urates can be mistaken for albumin. It con- sists in slowly dropping urine into a test-tube filled to about one- fourth with a highly colored solution of picric acid in water. In the presence of albumin a white cloud or sharply defined white turbidity is formed, and on warming the liquid the albumin collects into balls which rise to the surface of the liquid. d. Potassium ferrocyanide test. 5 to 10 c.c. of cold uriue are acidu- lated with 5 to 10 drops of acetic acid, and to the mixture are added a few drops of solution of potassium ferrocyanide. In the presence of even traces of albumin a turbidity is caused. This test is extremely Fig. 43. Nitric acic test for urine. delicate, especially when modified so as to allow a few c.c. of diluted acetic acid, to which a few drops of potassium ferrocyanide solution had been added, to flow down the side of the test-tube containing the urine. A decided turbidity at the point of contact of the two liquids shows albumin. In case the addition of acetic acid to the cold urine should cause a turbidity (which may be due to mucin) it must be filtered before add- ing the potassium ferrocyanide. e. Metaphosphoric add (glacial phosphoric acid) or trichlor-acetic acid may be used for the detection of albumin by dropping a fragment of either substance into a few c.c. of urine contained in a test-tube. As the acids dissolve, a cloudy ring forms in the presence of albumin, which is not dissolved on warming. 458 PHYSIOLOGICAL CHEMISTRY. In the above methods the manipulations and precautions are min- utely described, in order to detect small quantities or even traces of albumin. When albumin is abundantly present, there is no difficulty whatever in its detection, as heat will precipitate it from an acid, neutral, or sometimes even alkaline urine; the precipitate should, however, always be tested by the addition of a few drops of nitric acid, and the previous addition of a few drops of acetic acid is also advisable. A neutral urine should never be acidified by nitric acid (instead of acetic acid), because a drop or two of nitric acid may in some cases prevent the coagulation of albumin by heat, though a larger quantity (10 to 20 drops) has no such effect. Quantitative estimation of albumin. The average amount of albu- min present in acute cases of albuminuria is 0.1 to 0.5 per cent., rarely over 1 per cent., though it may rise to 4 per cent. An approximate method for the comparative estimation of albumin is to precipitate it (with the precautions above given) in a graduated test-tube by heat and setting aside for twelve (or better for twenty-four) hours. At the end of that time the proportion of the coagulated albumin which has collected at the bottom of the fluid is noticed. If the albumin occupy one-fourth, one-sixth, one-tenth of the height of the liquid, there is said to be one-fourth, one-sixth, or one-tenth of albumin in the urine. If, however, at the end of twelve or twenty-four hours scarcely any albumin has collected at the bottom, there is said to be a trace. The volumes of coagulated albumin indicate the following quantities of dry- albumin : Slight turbidity indicates about 0.01 per cent. of the tube is filled ....... 0.05 “ “ “ 0.10 “ i “ “ 0.25 “ i “ “ 0.50 “ 4 “ “ 1.00 “ Complete coagulation 2 to 3 “ A better method of exactly estimating the amount of albumin is its complete separation and weighing, as described below. Experiment 76. Acidify 100 c.c. of clear albuminous urine with acetic acid; heat to the boiling-point in a water-bath for half an hour, and filter through a small filter, previously dried at 110° C. (230° F.) and weighed; wash with boil- ing water to which a little ammonia water has been added (to remove uric acid and urates), then with pure water until the filtrate is not rendered turbid any EXAMINATION OF NORMAL AND ABNORMAL URINE. 459 longer by silver nitrate, next with pure alcohol, and finally with ether. Dry and filter contents at flO0 C. (230° F.) and weigh. As it may happen that the precipitated albumin encloses earthy phosphates,, it is well to burn filter with contents in a platinum crucible, and to deduct the weight of the remaining inorganic residue (less the weight of the filter ash), from that of the albumin. Peptones may be recognized by first precipitating all albumin by means of boiling the urine acidified by acetic acid, filtering and pour- ing the cold filtrate carefully upon some Fehling’s solution contained in a test-tube. At the junction of the two fluids a phosphatic cloud will appear, above which there will be seen a rosy-tinted layer in the presence of peptones. In case the albumin should not have been separated completely, the color will be more violet. Blood. The presence of blood in urine manifests itself generally,, unless the amount be too slight, by a blood-red or brownish color with a bluish, smoky, or greenish tint, and deposits a red or reddish- brown sediment after standing. As a general rule, all constituents of blood, including the corpuscles, are present, but in some cases haemo- globin alone has been found. The tests for blood depend either on the microscope or on chemical changes. By the microscope is examined the deposit which forms on standing; almost unaltered blood corpuscles may be found, or they may be much swollen, decolorized, aud deformed; the corpuscles are generally accompanied by blood and fibrin casts. Whenever blood is present, there are necessarily also albuminoids, which are precipitated by acidulating with acetic acid and boiling, when a brownish coagulum of albumin and hsematin are precipitated. Haemoglobin is also tested for by means of adding to the urine a few drops of freshly prepared tincture of guaiacum, a little ozonized ether, and shaking well. If haemoglobin is present, the ether assumes a blue color. Detection of sugar. The sugar found in urine is almost exclusively glucose, C6H1206. Traces of sugar, or as much as 0.01 per cent., are said to occur normally in urine, and are of no significance; moreover, it is as yet doubtful whether these traces of sugar are actually present in normal urine. A large amount of sugar is often indicated by a high specific gravity of the urine, which then varies from 1.030 to 1.050 ; the quantities found vary from mere traces to 10 per cent., the latter quantity, however, being a very rare occurrence, while 3 to 5 460 PHYSIOLOGICAL CHEMISTRY. per cent, is often found in the urine of persons suffering from diabetes mellitus. There are many tests by which sugar can be detected. They depend chiefly on the following properties of sugar, viz.: 1, to act as a deoxid- izing or reducing agent upon certain metallic oxides (copper, bismuth, silver, mercury) in the presence of alkalies ; 2, to produce a yellow or brown color, when in contact with alkalies, slowly in the cold, rapidly on heating; 3, to give a deep red color with picrates in alkaline solu- tion, and a number of different colors with certain phenols in the presence of sulphuric acid; 4, to ferment with yeast; 5, to unite with phenyl-hydrazine to a crystalline compound; 6, to have the power of rotation to the right of the plane of polarization. The tests for sugar should always be preceded by tests for albumin, which latter, if present, should be removed by coagulation and filtra- tion. Earthy phosphates also interfere with the copper tests some- times, because they are precipitated by the alkali, and this precipitate may be mistaken either for precipitated cuprous oxide, when no sugar is present, or it may cover the precipitated cuprous oxide to such an extent that this is not recognized, when sugar is present. To avoid these errors, it is well to render slightly alkaline the urine by a few drops of potash solution, filter after a few minutes, and use this urine for the tests. Trommel'’s test. A few drops (2-4) of a 5 per cent, solution of cupric sulphate are added to about 5 to 8 c.c. of urine in a test-tube and then an equal volume of potassium (or sodium) hydroxide solu- tion is added. The alkaline hydroxide precipitates both earthy phos- phates and cupric hydroxide, the latter, however, dissolving (especially if sugar be present) in the excess of the alkali, producing a beautiful blue transparent liquid. (If no sugar is present, the color is less blue, but more of a greenish hue.) The liquid is now heated, when, if sugar be present, a yellow precipitate of cuprous hydroxide is formed which subsequently loses its water and becomes the red cuprous oxide, which falls to the bottom or adheres to the sides of the test-tube. (Plate VII., 5.) As various organic substances (other than sugar) have a tendency to reduce cuprous oxide at a temperature of 100° C. (212° F.), it is well to set aside a test-tube prepared as above (without heating it) for from six to twenty-four hours. If sugar be present, the formation of cuprous hydroxide will gradually take place, whilst most other organic matters do not act upon cupric oxide at ordinary temperature. In drawing conclusions from the above test, it should be remem- EXAMINATION OF NORMAL AND ABNORMAL URINE. 461 berecl that a change of color does not indicate sugar ; that a precipitate of earthy phosphates must not be mistaken for cuprous oxide; and that substances other than sugar may deoxidize cupric oxide at the temperature of 100° C. (212° F.). Fehling’s test differs from Trummer’s test in merely using a pre- viously mixed reagent instead of producing this reagent, as it were, in the urine by adding to it cupric sulphate and an alkaline hydroxide successively. This reagent, known as Fehling’s solution, or as alkaline cupric tartrate volumetric solution, is made by mixing exactly equal volumes of the below-mentioned copper solution and the Rochelle salt solution at the time required. Copper solution: Crystallized cupric sulphate .... . 34.64 grammes, Water, sufficient quantity to make . . 500 c.c. Rochelle salt solution: Potassium sodium tartrate .... . 173 grammes. Potassium hydroxide ..... . 125 “ Water, sufficient quantity to make . . 500 c.c. Both solutions are preserved in small well-stoppered bottles, and mixed only at the time needed, because the mixture is apt to decom- pose when kept some time. The addition of sodium-potassium tartrate in Fehling’s solution prevents the precipitation of cupric hydroxide by the alkaline hydroxide. This action is analogous to the formation of the soluble scale compounds of iron, where the precipitation of ferric hydroxide is also prevented by tartaric or other organic acids. While Fehling’s solution is used chiefly for quantitative determina- tions, it can also be used to advantage for qualitative tests. This is done by heating about 10 c.c. of Fehling’s solution in a test-tube, and adding drop by drop the suspected urine; if the latter contains larger quantities of sugar a yellow or red precipitate of cuprous hydroxide and oxide will be produced very readily; if but small quantities are present, an equal volume of urine may be added to the solution, and the boiling repeated several times before the reaction takes place. Botger’s bismuth test consists in adding to a mixture of equal volumes of urine and potassium (or sodium) hydroxide solution a few grains of subnitrate of bismuth and boiling for half a minute. If sugar be present, a gray or dark-brown, finally black, precipitate of bismuthous oxide, Bi202, or of metallic bismuth is formed. If but very little sugar is present, the undecomposed excess of bismuthic nitrate (or bismuthic hydroxide) mixes with the metallic bismuth, 462 PHYSIOLOGICAL CHEMISTRY. imparting to it a gray color; the test should then be repeated with a smaller amount of the bismuth salt. (Plate VII., 6.) The above test may be somewhat modified by using a bismuth solu- tion, instead of the powder. The solution known as Nylander’s reagent is made by dissolving 2 grammes of bismuth subnitrate, 4 grammes of Rochelle salt, and 10 grammes of sodium hydroxide in 90 c.c. of water, and filtering. One-half c.c. of this solution is heated with about 5 c.c. of urine, when, in the presence of sugar, a brown or black precipitate will form after a few minutes. If the urine contains hydrogen disulphide (sometimes produced by decom- position of certain urinary constituents), black bismuth sulphide will be formed, which may be mistaken for metallic bismuth; albumin itself may be the cause of the formation of alkaline sulphides: the previous complete sepa- ration of albumin is therefore indispensable. Moore’s or Heller’s test is made by heating urine with about one- fourth of its volume of solution of potassium hydroxide. In the presence of sugar the color of the mixture will deepen to a dark yellow or brown, and the depth of color is a fair indication of the quantity of sugar present. In case but a slight change takes place in color, it is well to compare it with that of an unchanged specimen of the urine. The fermentation test is based upon the decomposition of sugar by the action of yeast with generation of carbon dioxide. The test is made by adding to about 50 or 100 c.c. of urine (contained in a large test-tube or small flask) a few grammes of common yeast. The vessel containing the urine is provided with a perforated cork, through which is passed one limb of a bent glass tube, long enough to reach nearly to the bottom of the vessel, which should be completely filled with urine. Under the second limb of the bent glass tube is placed a beaker. The apparatus thus prepared, is placed iu a room having a temper- ature of about 22°—28° C. (72°-82° F.). If sugar be present, fer- mentation will commence within twelve hours, and will manifest itself by the formation of carbon dioxide, which will force a portion of the fluid through the bent tube into the beaker placed there for its recep- tion. The disadvantages of this process are the length of time required for its per- formance, the unreliability of the ferment, and the fact that small quantities of sugar (less than 0.5 per cent.) evolve so little carbon dioxide that a doubt may be felt as to the presence of sugar at all. Picric acid test for sugar. It has been mentioned above that picric acid serves as an excellent reagent for albumin ; in the presence of EXAMINATION OF NORMAL AND ABNORMAL URINE. 463 alkalies it may also be used to advantage as a reagent for sugar. Urine is mixed with a few drops of a saturated aqueous solution of picric acid, a little caustic potash is added and gently heated; a marked reddish or reddish-brown coloration, due to the formation of picramic acid, H.C6H2.XH2.(N02)20, indicates sugar. A reddish color which appears without heating the mixture and which disappears completely within twenty minutes indicates the presence of kreatinin. A portion of the reddish liquid heated will turn more intensely red if sugar is also present. Molisch’s test. This is made by adding to urine a few drops of a 10 per cent, alcoholic solution of either thymol, menthol, or alpha-naph- thol. Into the inclined test-tube about 2 c.c. of concentrated sulphuric acid are then poured so as to form a layer below the urine. At the zone of contact a color is produced which is red with thymol and menthol, violet with greenish borders with alpha-naphthol. Traces of glucose are shown by these tests. Phenyl-hydrazine test is made by heating to boiling a mixture of equal volumes of urine and potassium hydroxide solution, to which a few drops of phenyl-hydrazine have been added. In the presence of sugar the mixture assumes an intense yellow or orange color. Upon supersaturating the cooled mixture with acetic acid a precipitate of golden yellow, needle-shaped crystals of phenyl-dextros-azon is formed. The test has the advantage that glucose is the only substance likely to occur in urine, which forms these crystals. Quantitative estimation of sugar. By far the best method ig the decomposition of a copper solution of a known strength, and Fehling’s solution prapared as stated above, answers this purpose well. 1000 c.c. of Fehling’s solution, containing 34.64 grammes of crys- tallized cupric sulphate, CuS04.5H20, are decomposed by 5 grammes of grape-sugar, or 1 c.c. of solution by 0.005 of grape-sugar. To make the quantitative determination, operate as follows : 10 c.c. of Fehling’s solution are poured into a porcelain dish of about 200 c.c. capacity, placed over a flame. The copper solution is diluted with about 40 c.c. of water, and heated to boiling; to the boiling liquid, urine (which has been previously diluted with 9 parts of water) is added from a burette very gradually, until the blue color of the solu- tion has disappeared, and there remains, upon subsidence of the cuprous oxide, an almost colorless, clear liquid. A filtered portion of this liquid, acidified with hydrochloric acid, should not give a reddish-brown pre- cipitate with potassium ferrocyanide (a precipitate would show that all 464 PHYSIOLOGICAL CHEMISTRY. copper had not been precipitated, and that more urine was needed), whilst a second portion of the filtered fluid should not produce a red precipitate on boiling with a few drops of Fehling’s solution (a pre- cipitate would indicate that too much urine had been added, in which case the operation has to be repeated). The calculation of the amount of sugar present is easily made. 10 c.c. of Fehling’s solution are decomposed by 0.05 gramme of sugar; this quantity must, therefore, be contained in the number of c.c. of urine used. Suppose 80 c.c. of urine, diluted with 9 parts of water, or 3 c.c. of pure urine, have been required to decompose the 10 c.c. of Fehling’s solution, then 3 c.c. of urine contain of grape-sugar 0.05 gramme, or 100 c.c. of urine 1.666 grammes, according to the proportion : 3 : 0.05 : : 100 : x x = 1.666. If the urine contains but very little sugar, it may be used directly without diluting it, or instead of diluting it with 9 parts of water, it may be diluted with 4 volumes or with an equal volume of water. Gerard’s determination by cyano-cupric solution. One disadvantage of the above process is the difficulty of determining readily the exact point at which the cupric solution has been converted into a cuprous compound. This difficulty is overcome to a great extent by making use of the power of potassium cyanide to form with cupric solutions a colorless compound which, while not acted upon by sugar solutions, is capable of dissolving cuprous oxide. If, therefore, a certain quantity of gelding’s solution is exactly decolorized by potassium cyanide, and to this solution is added more Fehling’s solution and the mixture heated with sugar, the Fehling’s solution will be decolorized, while the cuprous oxide remains in solution. Three solutions are required for this test: Solution No. 1. Cupric sulphate crystallized 69.28 grammes. Water, sufficient quantity to make . . . . 500 c.c. Solution No. 2. Sodium tartrate, crystallized 175 grammes. Sodium hydroxide 76.56 “ Water, sufficient quantity to make .... 500 c.c. Solution No. 8. Potassium cyanide (98 per cent.) .... 33 grammes. Water, sufficient quantity to make . . . . 500 c.c. If the potassium cyanide should contain less than 98 per cent, of this salt, solution No. 8 must be tested as follows : Five c.c. of each of EXAMINATION OF NORMAL AND ABNORMAL URINE. 465 the three solutions are mixed with 50 c.c. of water, then boiled. To the boiling liquid a solution of grape-sugar is added until the blue color is discharged. If any precipitate is formed, more cyanide must be added to No. 3 until the point is reached when, on boiling equal volumes of the mixed solutions with grape-sugar, no precipitate is formed. The determination is then made by adding from a burette the urine (diluted if necessary) to the boiling liquid made by mixing 10 c.c. each of solutions Nos. 1, 2, and 3, in the order given, and diluted with about 100 c.c. of water. The decomposition is complete when all blue color has disappeared. The calculation is made as described above under Fehling’s solution. It should be mentioned that the decolorized copper solution reassumes its blue color by oxida- tion, especially when boiling is stopped. For accurate observation it is usual to make a second and more rapid estimation, so as to check error that may arise from too long exposure of liquid to atmospheric oxygen. Determination by fermentation. The fermentation test above de- scribed can be used for quantitative determination of sugar, provided the quantity present is not less than 0.5 per cent., when the results are fairly accurate. The determination is made by observing carefully the specific gravity of the urine at the same temperature before and after fermentation. The decomposition of the sugar causes the specific gravity to become less, and every degree of the urinometer indicates 0.219 per cent, of sugar. If, for instance, urine showed a specific gravity of 1032 before, and 1022 after fermentation, the quantity of sugar present is 10 times 0.219, or 2.19 per cent. The yeast to be used for the experiment should be well washed upon a filter with pure water, and the urine quickly filtered before taking its specific gravity after fermentation has taken place. Experiment 77. Determine the amount of sugar in urine by the methods described above. If no suitable urine is to be had, add some glucose to urine and use this solution. Detection of bile. The presence of bile in urine is generally indi- cated by a decided color, which varies from a deep brownish-red to a dark brown ; the foam of such urine (produced by shaking) has a distinct yellow color, and a piece of filtering-paper or a piece of linen dipped into the urine assumes a yellow color, which does not dis- appear on drying. The further detection of bile depends upon the reactions of the biliary coloring matters or biliary acids; it frequently happens, however, that the pigments are present, whilst the acids are not. 466 PHYSIOLOGICAL CHEMISTRY. Gmelin’s test for biliary coloring matters has been considered al- ready, and may be applied to urine either by allowing a small quan- tity of nitric acid, containing some nitrous acid, to flow down the sides of a test-tube (containing the urine) in such a manner that the two fluids do not mix, or by placing upon a porcelain plate a few drops of the urine, near it a few drops of nitric acid, to which one drop of sul- phuric acid has been added, and allowing the two liquids to approach gradually. In both cases (if -bile pigment is present) a play of color is seen at the point of union between the two fluids, the colors changing from green to blue, violet-red, and yellow or yellowish-green. While the appearance of the green at the beginning is indispensable to prove the presence of bile, the presence of all the other colors is not essential. (Plate VII., 7.) The above test may be made in a somewhat modified form by mix- ing the urine with a concentrated solution of sodium nitrate, and pouring down the sides of the test-tube concentrated sulphuric acid in such a manner as to form two distinct layers; the colors are seen at the point of contact as above. If the urine be very dark in color, it should be diluted with water before applying the above tests. Ultzmann’s test for bile pigment is made by mixing 10 c.c. of urine with 3 or 4 c.c. of potassium hydroxide solution (1 in 3 of water), and supersaturating with hydrochloric acid ; the mixture assumes a beauti- ful emerald-green color after some time. Pette7ikofer’s test for biliary acids is made by dissolving a few grains of cane-sugar in urine contained in a test-tube, and allowing some concentrated sulphuric acid to trickle down the side of the inclined test-tube ; a purple band is seen at the upper margin of the acid, and on slightly shaking the liquid becomes at first turbid, then clear, and almost simultaneously it turns yellow, then pale cherry-red, dark car- mine-red, and finally a beautiful purple violet. The temperature must not be allowed to rise much above 38° C. (100° F.). (Plate VII., 8.) As many substances (other than biliary acids) show a similar reaction, it is often necessary to separate the bile acids by the process described in connection with the consideration of bile itself. In case the quantity of biliary constituents is so small that they cannot be noticed by the tests mentioned, the urine should be shaken with about one-fourth of its volume of chloroform, which dissolves the biliary matters. Some of this solution is dropped upon blotting paper, and after evaporation a drop of red fuming nitric acid is placed in the EXAMINATION OF NORMAL AND ABNORMAL URINE. 467 centre of the remaining stain, when concentric color rings appear. The second portion of chloroform solution is evaporated and the residue used for making the reactions as described above. Diazo-reaction. Some abnormal constituent (which has not been isolated yet) is found in the urine of persons suffering from typhoid fever, measles, sepsis, as also in some cases of phthisis, and possibly in other diseases. The presence of this unknown substance is indicated by a very characteristic reaction with diazo-benzol-sulphonic acid, which substance is produced by the action of nitrous acid on sulph- anilic acid. The test is made as follows: 1 gramme of sulphanilic acid is dissolved in a mixture of 350 c.c. of water and 15 c.c. of nitric acid. A second solution is made by dissolving 0.5 gramme of sodium nitrite in 100 c.c. of water. 5 c.c. of urine are mixed with an equal volume of the sulphanilic acid solution, then are added 3 to 4 drops of the sodium nitrite solution and finally 10 drops of ammonia water. Normal urine, thus treated, shows a deep yellow or orange color; the precipitated phosphates as well as the foam are colorless. In case the unknown substance is present the urine assumes gradually a carmine- red color; the foam produced on shaking is distinctly red and the phosphates show a greenish or violet color. As many phenol-deriva- tives give a similar color reaction with diazo-benzol-sulphonic acid, the presence of such substances in the urine, and derived from medi- cation, may lead to erroneous conclusions. Sugar shows the intense red color, especially on heating, when the test is made as above, but with the difference that potassium hydroxide is used in place of ammonia water. Urinary deposits (sediments). Normal urine is always clear, but occasionally, and particularly in abnormal conditions, it is turbid. Urine may be turbid when passed, and this indicates an excess of mucus, or the presence of renal epithelium, pus, blood, chyle, semen, bile, or phosphate or urate of sodium in excess, etc. A turbidity subsequent to the passage of the urine is generally due to the precipitation of phosphates or urates, or it may result from fermentation or decomposition. Either of the substances named will form a deposit on standing. When such a deposit is to be examined, a few ounces of the urine should be set aside for several hours in a tall, narrow, cylindrical glass; when the sediment has collected at the bottom, the supernatant 468 PHYSIOLOGICAL CHEMISTRY. urine may be decanted, or the sediment may be taken out by means of a pipette for examination. Sediments are either organized or unorganized. To the first belong : mucus, blood, pus, urinary casts, epithelium, spermatozoids, fungi, infusorise, etc.; to the second belong: uric acid, urates, calcium oxa- late, phosphate, or carbonate, magnesium-ammonium phosphate, cystin, hippuric acid, etc. The chemical examination of any urinary sediment should always be preceded by a microscopical examination, which latter is in many cases the only way of determining the nature of the sediment, espe- cially of the organized substances. Most of the unorganized and either crystalline or amorphous sediments may be easily recognized by chemical means. Urates of ammonium, calcium, and sodium dissolve on heating the urine, and are reprecipitated on cooling. The murexid test is used in addition. Phosphates of calcium or ammonium-magnesium dissolve in acetic acid, and ammonium molybdate dissolved in nitric acid produces a yellow precipitate on heating. Calcium oxalate is insoluble in acetic, but soluble in hydrochloric acid, from which solution it is reprecipitated on neutralizing with ammonia. Uric acid is not dissolved by heat, nor by acetic or hydrochloric acid, but dissolves on the addition of caustic potash and burns on platinum foil without leaving a residue; it is recognized by the murexid test. Cystin is insoluble in water and alcohol, but soluble in mineral acids and in caustic alkalies; from either solution it is reprecipitated by neutralizing. Cystin contains 26 per cent, of sulphur, which causes the formation of black sulphide of lead when cystin is boiled with caustic potash to which a few drops of solution of lead acetate have been added. Urinary calculi are solid deposits of larger or smaller size formed from the urine within the tracts (kidneys, ureter, bladder, and urethra). The chemical composition of the calculi is generally that of either of the above-named unorganized sediments, and their nature can easily be determined by using the following method : Make a section through the centre of the calculus, scrape some of the substance off, powder it finely, and heat some of it on platinum foil. It may either burn away completely (uric acid, urate of ammo- EXAMINATION OF NORMAL AND ABNORMAL URINE. 469 ilium, cystin, xanthin) or may be partially combustible (urates or oxalates), or may be incombustible (chiefly phosphates). A slight blackening occurs generally, even in heating a calculus consisting of incombustible matter, and is due to the presence of traces of organic urinary constituents. If completely combustible, digest a little of the powder with dilute hydrochloric acid ; cystin and xanthin are dissolved, uric acid remains undissolved. Apply murexid test for uric acid, the above-mentioned lead test for cystin, and for xanthin test by dissolving a little of the calculus in nitric acid and evaporating to dryness, when in the presence of xanthin a bright-yellow residue will be left, which becomes violet-red when treated with caustic potash. In case uric acid has been found, it may be in combination with ammonia, which may be verified by heating the powder with a little caustic potash, when ammonia gas is liberated, which may be recognized by its action on red litmus-paper, odor, etc. If partially combustible or incombustible, digest some of the powder with dilute hydrochloric acid. If it dissolves completely, uric acid is not present. If a residue be left, apply the murexid test. To a portion of the solution add ammonium molybdate and heat; a yellow precipitate indicates phosphoric acid. To another portion add am- monia water and then excess of acetic acid; a white pulverulent residue indicates calcium oxalate, which can be verified by igniting some of the calculus and adding a drop of acid, when effervescence will be noticed, the oxalate having been converted into a carbonate by the ignition ; the solution thus obtained can be tested for calcium by the addition of water of ammonia and ammonium oxalate. In case phosphoric acid has been found, this is present either as a calcium or magnesium-ammonium salt. To distinguish between them, neutralize the solution of the powder in hydrochloric acid with ammonia, add acetic acid and ammonium oxalate; a white precipitate indicates calcium ; if no precipitate is produced, supersaturate with ammonia, when the crystalline magnesium-ammonium phosphate will gradually form. Most common are calcnli of uric acid; often met with are those of urates, phosphates, and oxalates; rarely, however, those of xanthin and cystin. Microscopical examination of urinary sediments. The chemical examination of any urinary sediment should always be preceded by a microscopical examination, which latter is in many cases the only 470 PHYSIOLOGICAL CHEMISTRY. way of determining the nature of the sediment, especially of the organized substances. Fig. 44, A-O, shows the principal sediments found in, or produced from, urine, as seen with a magnifying power of 200 diameters. A. Uric acid occurs in many different forms, mostly in rhombic plates, with rounded obtuse angles, often joined into rosettes. Uric acid is found almost invariably colored red or reddish-brown, which generally distinguishes it from other sediments. The crystals or clusters of crystals are often large enough to be seen by the naked eye, and are then known by the terms “ sand,” “ gravel,” or u red- pepper grains.” B. Ammonium acid urate is found, generally associated with amor- phous or crystalline phosphates, in urine which has become alkaline. The crystalline globules are generally covered with spinous excres- cences, which give them the characteristic “ thorn-apple ” appearance. C. Sodium urate forms generally a part in the pulverulent, heavy, variously tinted deposit of the mixed urates known as “ brickdust ” or “ lateritious ” sediment. It occurs either in fine amorphous granules which cannot be distinguished microscopically from other amorphous sediments or in a crystalline form as shown in the figure. D. Urea nitrate crystallizes readily in large six-sided plates on the addition of nitric acid to urine. E. 1, Leucin, or amido-caproic acid, C6Hn(NH2)02; and 2, Tyrosin, c9huno3, are but rarely met with in urinary deposits. Leucin is found either as rounded lumps, showing but little crystalline struc- ture, or as spherical masses, exhibiting fine radial striation. Tyrosin appears generally in fine, long, silky needles, forming bundles or rosettes. F. Cystin occurs occasionally as a grayish, crystalline deposit, forming transparent six-sided plates; it also occurs in calculi. The latter may be recognized by the above-mentioned chemical properties or by dissolving a little in hydrochloric acid and neutralizing with ammonia, when cystin is reprecipitated and shows the characteristic six-sided plates under the microscope. G. Magnesium-ammonium phosphate, or triple phosphate, MgNII4- P04.6H20, is found generally in triangular prisms with bevelled ends, as shown in 1, but sometimes also in star-shaped, feathery crystals, represented in 2. EXAMINATION OF NORMAL AND ABNORMAL URINE. 471 Fig. 44. Urinary sediments. 472 PHYSIOLOGICAL CHEMISTRY. H. Calcium 'phosphate, Ca3(P04)2, is most frequently found in amor- phous globules, but also crystallized either in prisms, 1, or in “wedge- shaped ” crystals, 2. I. Calcium oxalate, CaC204, occurs either in quadratic octohedra with brilliant refraction, 1, or sometimes in the shape of “dumb- bells,” 2. J. Blood corpuscles appear under the microscope as reddish, circular disks, sometimes laid together in strings. If seen in profile, they appear biconcave. 1 shows the corpuscles in a fresh condition ; 2, as generally seen in urine. K. Mucus and pus are often difficult to distinguish from one another under the microscope, as they both appear as little granular globules, varying somewhat in appearance with the reaction of the urine. Pus is rendered slimy, ropy, viscid, and tenacious by the addition of caustic potash. 1 shows globules of mucus, 2 of pus, and 3 of pus treated with acetic acid, which clears up the granular globules with the appearance of a nucleus. L. Hcemin crystals. The formation of these crystals often serves to recognize blood and is accomplished by mixing the latter on a glass slide with a trace of sodium chloride aud a drop of glacial acetic acid and warming gently, when the characteristic crystals will appear. By repeating the process several times, larger and better-developed crystals are obtained. M. 1, Hyaline casts; 2, Granular casts. Urinary casts are tube- like cylinders, often found together with blood and pus corpuscles, or holding in their substance or walls epithelial cells, mucous corpuscles, and fat globules. Hyaline casts are distinguished by their trans- parent appearance, while granular casts show a more or less granular surface. N. Epithelial casts and cells. According to the origin (vagina, urethra, bladder, etc.) of these bodies, they differ somewhat, and it is difficult to recoguize with certaiuty the source whence they are derived. O. 1, Waxy casts ; 2, Casts ivith blood corpuscles ; 3, Casts with fat globules. Waxy casts resemble hyaline casts, but are less transparent. Casts containing blood corpuscles or fat globules are generally easily recognized. In addition to the above-mentioned urinary deposits there may also be found various kinds of fungi, vibrionae, spermatozoids, hair, or EXAMINATION OF NORMAL AND ABNORMAL URINE. 473 even such foreign matters as fibres of cotton, wool, or silk, with the characteristic appearance of which the student should familiarize him- self thoroughly. Questions.—551. What points are to be considered, and what substances determined, in the analysis of normal and abnormal urine ? 552. What is the color of urine, and what are the chief causes influencing the color? 553. What is the specific gravity of healthy urine, how is it determined, and how is the total amount of solids approximately calculated from the specific gravity? 554. Describe the different tests by which albumin may be recognized, and state the precautions necessary in making these tests. 555. How may the quan- tity of albumin in urine approximately and how accurately be determined ? 556. Describe the various tests for sugar. On what principles are they based? 557. How is sugar determined quantitatively ? 558. By what tests are biliary pigments and acids recognized in urine? 559. What is the nature of urinary sediments, and by what means are they recognized ? 560. What are urinary calculi generally composed of, and by what simple tests can their nature be determined ? APPENDIX. TABLE OF WEIGHTS AND MEASURES. Measures of length. 1 millimeter = 0.001 meter = 0.0394 inch. 1 centimeter = 0.01 meter = 0.3937 inch. 1 decimeter = 0.1 meter = 3.9371 inches. 1 meter = 39.3708 inches. 1 decameter = 10 meters = 32.8089 feet. 1 hectometer = 100 meters = 328.089 feet. 1 kilometer = 1000 meters = 0.6214 mile. 1 yard or 36 inches =. 0.9144 meter. 1 inch == 25.4 millimeters. Measures of capacity. 1 milliliter = 1 c.c. = 0.001 liter = 0.0021 U. S. pint. 1 centiliter = 10 c.c. = 0.01 liter = 0.0211 U. S. pint. 1 deciliter = 100 c.c. = 0.1 liter — 0 2113 U. S. pint. 1 liter = 1000 c.c. = 1.0567 IT. S. quart. 1 decaliter = 10 liters = 2.6418 U. S. gallons. 1 hectoliter = 100 liters = 26.418 U. S. gallons, 1 kiloliter = 1000 liters = 264.18 U. S. gallons, 1 U. S. gallon = 3785 3 c.c. 1 imperial gallon = 4543 5 c.c. 1 minim = 0.06 c.c. 1 fluidrachm = 3 70 c c. 1 fluidounce = 29 57 c.c. 1 liter == 33 81 fluidounces. Weights. 1 milligram = 0.001 gramme == 0.015 grain Tro}’. 1 centigram = 0.01 gramme == 0.154 grain Troy. 1 decigram = 0.1 gramme = 1.543 grain Troy. 1 gramme = 15.432 grains Troy. 1 decagram = 10 grammes = 154.324 grains Troy. I hectogram = 100 grammes = 0.268 pound Troy. I kilogram = 1000 grammes = 2.679 pounds Troy. 1 grain Troy = 0.0648 gramme. 1 drachm Troy = 3.888 grammes. 1 ounce Troy = 31.103 grammes. 1 ounce avoirdupois = 28.350 grammes. 1 pound avoirdupois ‘ = 453.592 grammes. 476 APPENDIX. Commercial weights and measures of the U. S. A. 1 pound avoirdupois = 16 ounces. 1 ounce = 437.5 grains. 1 gallon = 231 cubic inches. 1 gallon = 4 quarts = 8 pints. 1 pint of water weighs 7291.2 grains at a temperature of 15.6° 1 drachm = 60 grains. 1 ounce = 8 drachms = 480 grains. Troy weight. TABLE OF ELEMENTS. Symbol. Atomic weight. Aluminum . A1 27.04 Antimony . . Sb 119.6 Arsenic As 74.9 Barium . Ba 136.9 Beryllium1 . Be 9.03 Bismuth . . Bi 208.9 Boron . B 10.9 Bromine . . Br 79.76 Cadmium . . Cd 111.5 Caesium . Cs 132 7 Calcium . . Ca 39.91 Carbon . C 11.97 Cerium . Ce 139.9 Chlorine . . Cl 35.37 Chromium . Cr 52.0 Cobalt . Co 58 6 Columbium2 . Cb 93.7 Copper . Cu 63.18 Didymium5 . Di 142.0 Erbium . Er 166.0 Fluorine . . F 19 0 Gallium . Ga 69.9 Germanium . Ge 72 3 Gold . . Au 196.7 Hydrogen . . H 1.0 Indium . In 113.6 Iodine . I 126.53 Iridium . Ir 192.5 Iron . . Fe 55.88 Lanthanum . La 138.2 Lead. . Pb 206.4 Lithium . . Li 7.01 Magnesium • Mg 24.3 Manganese . Mn 54.8 Mercury • Hg 199.8 Symbol. Atomic weight. Molybdenum . Mo 95.9 Nickel . Ni 58.6 Nitrogen . . N 14.01 Osmium . . Os 190.3 Oxygen . 0 15.96 Palladium . . Pd 106.35 Phosphorus . P 30.96 Platinum . . Pt 194.3 Potassium . K 39.03 Rhodium . . Rh 102.9 Rubidium . . Rb 85.2 Ruthenium . Ru 101.4 Samarium. . Sm 149.62 Scandium . . Sc 43.97 Selenium . . Se 78.87 Silicon . Si 283 Silver • Ag 107.66 Sodium . Na 23.0 Strontium. . Sr 87.3. Sulphur • S 31.98 Tantalum . . Ta 182.0 Tellurium . . Te 125.0 Terbium . Tb 159.1 Thallium . . T1 203.7 Thorium . . Th 231.9 Tin . . Sn 118.8 Titanium . . Ti 48.0 Tungsten . . W 183.6 Uranium . . U 238.8 Vanadium . V 51.1 Ytterbium . Yb 172.6 Yttrium . . Yt 88.9 Zinc . . Zn 65.1 Zirconium. . Zr 90.4 1 Also called glucinum. 2 Also called niobium. 3 Composed of neo- and praseo-didymium. INDEX. { BSORPTION, 34 A Acetanilid, 380 Acetic acid 323 analytical reactions of, 325 aldehyde, 313 ether, 340 Acetone, 327 Acetylene, 296 Acid, acetic, 323 adipic, 329 aniline-para-sulphonic, 379 arabic, 350 arachidic, 322 arsenic, 206 arsenous, 205 behenic, 322 benzoic, 373 boric, 98 bromic, 123 butyric, 327 camphoric, 372 capric, 322 caproic, 322 caprylic, 322 carbamic, 355 carbazotic, 369 carbolic, 367 carbonic, 95 carminic, 352 cathartic, 352 cerotic, 322 chloric, 121 cholesterin, 432 cholic, 431 chromic, 176 citric, 334 copaivic, 371 cyanic, 359 diazo-benzol-sulphonic, 379 dithionic, 106 fluoric, 1 6 formic, 323 fulminic, 362 gallic, 375 glacial acetic, 324 phosphoric, 113 glycocholic, 431 glycolic, 336 hippuric, 448 hysenic, 322 hydriodic, 125 hydrobromic, 123 Acid, hydrochloric, 119 hydrocyanic, 356 hydroferricyanic, 360 hydroferrocyanic, 360 hydrofluoric, 126 hydrosulphuric, 106, 234 hypobromic, 123 hypochlorous, 121 hyponitrous, 89 hypophosphorous, 115 hyposulphurous, 106 kinic, 391 lactic, 335 lauric, 322 malic, 331 malonic, 329 manganic, 174 margaric, 322 meconic, 390 melissic, 322 metaphosphoric, 113 muriatic, 119 myristic, 322 myronic, 353 nicotinic, 382 nitric, 89 nitro-hydrochloric, 120 nitro-muriatic, 120 nitrous, 89 cenanthylic, 322 oleic, 328 orthophosphoric, 113 oxalic, 329 palmitic, 322 pelargonic, 322 pentath ionic, 106 perchloric, 121 permanganic, 174 phenol-sulphonic, 368 phospho-molybdic, 385 phosphoric, 113 phosphorous, 112 phthalic, 375 picric, 369 propionic, 322 prussic, 356 pyrogallic, 376 pyrophosphoric, 113 pyrosulphuric, 106 pyrotartaric, 329 rosolic, 263 salicylic, 374 480 INDEX. Acid, sarco-lactic, 336 silicic, 98 stannic, 217 stearic, 322 succinic, 329 sulphanilic, 379 sulphocarbolic, 368 sulphocyanic, 359 sulphonic, 319 sulphuric, 103 sulphurous, 102 sylvic, 371 tannic, 376 tartaric, 331 taurocholic, 431 tetrathionic, 106 thiosulphuric, 106 trithionic, 106 uric, 447 valerianic, 327 Acidimetry, 264 Acids, amido-, 354 aromatic, 366 biliary, 431 definitions of, 58 detection of, 241 fatty, 322 organic, 319 sulphonic, 319 thio-, 322 Aconitine, 396 Acrolein, 343 Actinic waves, 27 Adhesion, 32 ./Ether, 339 Affinity, chemical, 39 Agate, 97 Air, composition of, 85 Alabaster, 152 Albumin, 405 in urine, 455 Albuminates, 406 Albuminous substances, 403 analytical reactions of, 404 Albumoses, 408 Alcohol, 307 absolute, 309 amyl, 310 analytical reactions of, 309 butyl, 307 diluted, 309 ethyl, 307 methyl, 307 real, 309 Alcoholic liquors, 310 Alcohols, 304 monatomic, 307 Aldehyde, acetic, 313 benz-, 374 par-, 314 Aldehydes, 312 Alkali metals, 133 Alkalimetry, 265 Alkaline earths, 151 Alkaloids, 383 antidotes to, 385 cadaveric, 398 detection of, 385 Allotropic modifications, 66 Alloy, definition of, 132 Allyl-mustard oil, 353 sulphide, 353 Alnm, 159 Aluminum, 158 and ammonium sulphate, 159 analytical reactions of, 166 chloride, 161 hydroxide, 160 oxide, 161 sulphate, 161 Amalgam, 132 ammonium-, 145 tin-, 214 Amber, 371 Amides, 354 Amido-acetic acid, 354 -acids, 354 -formic acid, 355 Amine, diphenyl-, 379 ethyl-, 400 methyl-, 400 propyl-, 400 Amines, 354 Ammonia, 85 liniment, 344 water, 87 ,, Ammoniated mercury, 201 Ammonio-copper compounds, 188 Ammonium, 145 amalgam, 145 analytical reactions of, 147 benzoate, 373 bromide, 147 carbamate, 146 carbonate, 146 chloride, 145 hydroxide, 87 iodide, 147 molybdate, 219 nitrate, 146 phosphate, 146 sulphate, 146 sulphide, 147 sulphydrate, 147 valerianate, 328 Amorphism, 19 Amorphous phosphorus, 110 Amygdalin, 374 Amyl alcohol, 310 nitrite, 34 L Amylene hydrate, 311 Amyloid, 351 substance, 409 Amylopsin, 410 Analysis, definition of, 57 elementary, 279 gas-, 273 gravimetric, 256 INDEX. 481 Analysis, organic, 278 proximate, 278 qualitative, 221 quantitative, 256 ultimate, 279 urinary, 449 volumetric, 259 Analytical chemistry, 221 reactions of acetates, 325 alcohol, 309 albuminous substances, 404 aluminum, 162 ammonium, 147 antimony, 216 arsenic, 208 atropine, 395 barium, 157 benzoates, 373 bile, 431 bismuth, 191 blood, 425 borates, 99 bromides, 124 brucine, 394 calcium, 155 carbolic acid, 368 carbon, 92 carbonates, 95 cerium, 163 chloral, 315 chlorates, 122 chlorides, 120 chloroform, 316 cholesterin, 4 2 chromates, 178 chromium, 178 citrates, 335 cobalt, 183 cocaine, 396 codeine, 390 copper, 189 cyanides, 358 ferric salts, 172 ferricyanides, 361 ferrocyanides, 361 ferrous salts, 17 2 fluorides, 126 gastric juice, 426 srlvcerin, 311 gold, 218 grape-sugar, 347 hippuric acid, 448 hydriodic acid, 125 hydrobromic acid, 123 hydrochloric acid, 120 hydrocyanic acid, 358 hypochlorites, 122 hypophosphites, 115 iodides, 125 iron, 172 lead, 186 lithium, 148 magnesium, 151 manganese, 175 Analytical reactions of mercury, 201 morphine, 389 nickel, 183 nitrates, 91 nitrites, 89 oxalates, 330 phosphates, 114 phosphites, 112 platinum, 220 potassium, 139 quinine, 392 santonin, 377 silicates, 98 silver, 194 sodium, 144 starch, 350 strontium, 156 strychnine, 394 sugar, 347 sulphates, 105! sulphides, 107 sulphites, 103 tannic acid, 376 tartrates, 332 thiosulphates, 106 tin, 217 urates, 448 urea, 445 veratrine, 397 zinc, 182 Anilid, 380 Aniline, 378 dyes, 379 Animal charcoal, 154 cryptolites, 410 fluids and tissues, 421 food, 414 Anisidin, 369 j Anthracite coal, 301 | Antidotes to acids, 91 alkalies, 135 alkaloids, 385 antimony, 216 arsenic, 213 barium, 157 carbolic acid, 368 copper, 188 cyanides, 359 hydrocyanic acid, 359 lead, 186 mercury, 202 nitric acid, 91 oxalic acid, 330 phosphorus, 111 silver, 194 sulphuric acid, 105 zinc, 181 Antifebrine, 380 Antimonious chloride, 215 oxide, 215 Antimony, 213 analytical reactions of, 216 and potassium tartrate, 333 antidotes to, 216 482 INDEX. Antimony, black, 213 butter, 215 chloride, 215 crude, 213 oxide, 215 pentasulphide, 215 potassium tartrate, 215, 333 sulphide, 214 sulphurated, 214 trisulphide, 214 Antipyrine, 380 Antiseptics, 292 Apatite, 152 Apomorphine, 389 Aqua regia, 120 Arbutin, 352 Argentum, 192 Argol, 331 Aromatic compounds, 362 Arrack, 310 Arsenates, 206 Arsenic, 204 acid, 206 analytical reactions of, 208 antidotes to, 213 detection of, in case of poisoning, 212 oxide, 206 sulphides, 207 Arsenetted hydrogen, 206 Arsen ites, 205 Arsenous acid, 205 anhydride, 205 iodide, 207 oxide, 205 Arsine, 206 Artiads, 47 Asbestos, 149 Ash, bone-, 154 soda-, 141 Asphalt, 371 Atmospheric air, 85 pressure, 31 Atom, definition of, 40 { Atomic theory, 40 weights, determination of, 41, 48 Atoms, 39 quantivalence of, 45 Atropine, 394 analytical reactions of, 395 Auric chloride, 218 sulphide, 218 Auripigment, 204 Aurum, 217 Avogadro’s law, 24 Balsam, copaiva, 371 Balsams, 371 Barite, 157 Barium, 157 analytical reactions of, 157 antidotes to, 157 cai'bonate, 157 chloride, 157 chromate, 178 Barium dioxide, 157 oxide, 157 sulphate, 157 Barometer, 30 Basalt, 97, 159 Bases, definition of, 58 Beer, 310 Beet-sugar, 348 Bell-metal, 187 Benzaldehyde, 374 Benzene, 364 series, 362 Benzin, 301 Benzol, 364 Benzoic acid, 373 sulphinide, 381 Benzyl-glycocol, 448 Berberine, 398 Beryllium, 61 Bettendorff’s test, 210 Bicarbonate of potassium, 136 sodium, 142 Bichloride of mercury, 1981 Bichromate of potassium, 176 Bile, 430 detection of, in urine, 465 Biliary acids, 431 calculi. 432 pigments, 430 Bilirubin, 430 Biliverdin, 430 Bismuth, 190 analytical reactions of, 191 carbonate, 191 citrate, 335 hydroxide, 192 iodide, 191 nitrate, 190 oxide, 190 «' ;s, 190 onate, 191 ite, 190 x e, 190 191 BiSfeathyl, 190 carbonate, 191 iodide, 191 nitrate, 190 Bisulphide of carbon, 108 Biuret, 445 reaction, 404 Black antimony, 213 -ash, 141 -lead, 92 oxide of copper, 187 manganese, 173 mercury, 196 -wash, 196 Bleaching-powder, 155 Blood, 423 corpuscles, 423 detection of, 425, 459 -fibrin, 407 -serum, 424 INDEX. 483 Blood-stains, examination of, 425 Blue mass, 196 pill, 196 Prussian, 361 -stone, 188 Turnbull’s, 361 vitriol, 188 Bone, 433 -ash, 154 -black, 154 -oil, 381 Boric acid, 98 analytical reactions of, 99 Borax, 144 bead, 229 Boron, 98 Botger’s bismuth-test, 461 Brain, 435 Brandy, 310 Brass, 187 Brittleness, 20 Bromates, 123 Bromides, analytical reactions of, 124 Bromine, 122 Bromoform, 318 Bronze, 187 Brucine, 394 Burettes, 260 Butter, 438 -milk, 438 of antimony, 215 nADAVERIC alkaloids, 398 \J Cadaverine, 401 Cadmium, 182 iodide, 182 sulphate, 182 sulphide, 182 Caesium, 61 Caffeine, 398 ' Calamine, 179 Calcined magnesia, 150 Calcium, 151 analytical reactions of, 155 bromide, 155 carbonate, 153 chloride, 155 fluoride, 125 hydroxide, 152 hypochlorite, 155 hypophosphite, 154 oxalate, 330 oxide, 152 phosphate, 154 sulphate, 153 superphosphate, 154 tartrate, 331 Calc-spar, 152 Calculi, biliary, 432 urinary, 468 Calomel, 197 Camphor, 372 -mint, 373 Camphor, monobromated, 372 Cane-sugar, 348 Caoutchouc, 371 Capillary attraction, 33 Caramel, 347 Carbamide, 355, 443 Carbazotic acid, 369 Carbohydrates, 345 Carbolic acid, 367 analytical reactions of, 368 antidotes to, 368 Carbon, 92 bisulphide, 108 dioxide, 93 disulphide, 108 monoxide, 95 Carbonate, analytical reactions of, 95 Carbonic acid, 95 oxide, 95 Carbonyl, 320 Carboxyl, 320 Casein, 407 vegetable, 407 Cast-iron, 165 Casts, urinary, 458 Caustic, 193 lunar, 193 potash, 135 Celestite, 156 Celluloid, 351 Cellulose, 350 nitro-, 351 Cement, 434 Centigrade thermometer, 27 Cerebrin, 435 Cerite, 163 Cerium, 163 oxalate, 163 + Chains, 284 ' jf Chalk, 152 Charcoal, 92 ■ animal, 154 Cheese, 438 Chemical action, definition of, 39 affinity, 39 divisibility, 37 equations, 67 formulas, 41 reactions, 57 symbol, definition of, 41 Chemistry, analytical, 221 definition of, 40 how to study it, 69 organic, 274 physiological, 412 Chili saltpetre, 145 Chloral, 314 amide, 354 formamide, 354 hydrate, 315 Chlorates, analytical reactions of, 122 Chloric acid, 121 oxides, 121 Chlorides, analytical reactions of, 120 484 INDEX. Chlorinated lime, 155 Chlorine, 116 acids, 121 oxides, 121 -water, 119 Chloroform, 315 Chlorous oxide, 121 tetroxide, 121 Choke-damp, 96 Cholesterin, 344, 432 Cholic acid, 431 Choline, 401 Chromates, analytical reactions of, 178 Chrome-alum, 177 -iron ore, 175 -yellow, 186 Chromic acid, 176 hydroxide, 177 oxide, 177 Chromite, 175 Chromium, 175 chloride, 177 sulphate, 177 Chyle, 418, 425 Chyme, 417 Cinchona alkaloids, 391 Cinchonidine, 393 Cinchonine, 393 sulphate, 393 Cinnabar, 195, 201 Citrates, analytical reactions of, 335 Citric acid, 334 Clay, 161 Clot, 424 Coagulation, 403 Coal, 300 -oil, 300 -tar, 303 Cobalt, 179 Cocaine, 395 Codamine. 387 Codeine, 390 Cognac, 310 Cohesion, 18 Coke, 303 Colchicine, 387 Collagen, 411 Collidine, 382 Collodion, 351 Colloids, 35 Colocynthin, 352 Colophony, 371 Columbium, 61 Combustion, 76 Common salt, 140 Compound radicals, 60 Compounds, decomposition of, 53 definition of, 38 Congo paper, 429 Coniine, 387 Copaiva balsam, 371 Copper, 186 acetate, 326 ammonio-sulphate, 188 Copper, analytical reactions of, 189 antidotes to, 188 arsenite, 189 black oxide, 187 -glance, 187 hydroxide, 187 oxide, 187 pyrites, 187 sulphate, 188 sulphide, 187, 189 Copperas, 170 Corrosive chloride of mercury, 198 sublimate, 198 Corundum, 159 Cream, 437 of tartar, 332 Creamometer, 439 Creasote, 368 Crude antimony, 213 sulphur, 100 tartar, 331 Cryptolites, 410 Cryptopine, 387 Crystallin, 406 Crystallization, 19 Crystalloids, 35 Cubic nitre, 143 Cumene, 364 Cupric acetate, 326 arsenite, 189 ferrocyanide, 189 hydroxide, 187 oxide, 187 sulphate, 188 sulphide, 187, 189 Cuprous oxide, 187 Cuprum, 186 Curd, 437 Cyanhydric acid, 356 Cyanic acid, 359 Cyanides, analytical reactions of, 358 antidotes to, 359 Cyanogen, 356 Cymene, 370 Cystin, 468 DALTON’S atomic theory, 43 Decay, 289 Decomposition by electricity, 54 heat, 37, 53 light, 54 various modes of, 53 Decrepitation, 227 Deflagration, 228 Dentine, 434 ] Deodorizers, 292 Detection of impurities in officinal prep- arations, 248 | Deposits, urinary, 467 | Derivatives, 287 | Desiccator, 258 Destructive distillation, 288 | Dextrin, 350 INDEX. 485 Dextrose, 346 Diacetic ether, 380 Dialysis, 35 Dialyzed iron, 169 Diamond, 92 Diastase, 349 Diazo-reaetion, 467 Dibasic acids, 58, 329 Dicyanogen, 356 Didymium, 61 Diffusion, 34 Digitalein, 352 Digitalin, 352 Digitonin, 352 Digitoxin, 352 Digestion, 417 Dimorphism, 19 Disinfectants, 292 Distillation, 26 destructive, 288 dry, 288 fractional, 296 Disulphide of carbon, 108 Divisibility, 21 chemical, 37 Dolomite, 149 Donovan’s solution, 208 Double salts, 60 Dried alum, 160 Drinking water, 81 Dry distillation, 288 Drying-oven, 257 Ductility, 20 Dynamite, 312 PARTUS, 158 Lj alkaline, 151 Ebonite, 372 Ecgonine, 395 Elasticity, 20 Electricity, 54 Elementary analysis, 279 Element, definition of, 38 Elements, 61 derivation of names of, 71, 128 natural groups of, 62 non-metallic, 70 metallic, 128 relative importance of, 60 time of discovery of, 72, 128 valence of, 72 Emerald green, 326 Emery, 159 Emetine, 387 Empirical formulas, 281 Emulsine, 373 Enamel, 434 Energy, 18 Enzymes, 409 Epithelium, 434 Epsom salt, 150 Equations, chemical, 67 Equivalence, 47 Erbium, 61 Eserine, 387 Essential oils, 304 Esters, 337 Ethane, 298 Ethene, 96, 304 Ether, 339 acetic, 340 diacetic, 380 ethyl-, 339 nitrous, 341 sulphuric, 339 Ethers, 337 Ethyl, 307 acetate, 340 alcohol, 307 bromide, 318 ether, 339 hydroxide, 307 hydride, 298 nitrite, 341 oxide, 339 Ethylene, 303 Ethylic alcohol, 307 Extension, 17 FAHRENHEIT’S thermometer, 27 -T Fatty acids, 322 Fats, 342 Feathers, 424 Feces, 433 Fehling’s solution, 461 Feldspar, 134, 159 Fermentation, 290 Ferric acetate, 326 chloride, 168 citrate, 335 hydrate, 167 hydroxide, 107 hypophosphite, 171 nitrate, 171 oxide, 167 salts, analytical reactions of, 172 sulphate, 170 sulphocyanate, 172 tartrate, 334 valerianate, 328 Ferricyanogen, 360 Ferricyanides, analytical reactions of, 361 Ferrocyanides, analytical reactions of, 361 Ferrocyanogen, 360 Ferrous bromide, 169 carbonate, 171 chloride, 168 -ferric oxide, 167 hydroxide, 166 iodide, 169 lactate, 336 oxalate, 330 oxide, 166 phosphate, 171 salts, analytical reactions of, 172 sulphate, 170 486 INDEX Ferrous sulphide, 169 Ferrum, 164 Fibrin, 407 vegetable, 408 Fibrinogen, 406 Fire-damp, 96, 299 Flame, structure of, 97 -tests, 229 Fleitmann’s test, 210 Flowers of sulphur, 100 Fluorine, 125 Fluorspar, 125 Food, absorption of, 418 animal, 414 nitrogenous, 415 plant, 412 Force, definition of, 18 Formamide, 354 Formic acid, 323 Formulas, chemical, 41 empirical, 281 graphic, 283 molecular, 281 rational, 283 Fowler’s solution, 206 Fractional distillation, 296 Fruit-sugar, 347 Fusel oil, 310 Fusibility of metals, 129 Galena, 184 argentiferous, 192 Gallic acid, 375 Gallium, 61 Gall-stones, 432 Galvanized iron, 180 Gas-analysis, 273 definition of, 20 -furnace, 279 illuminating, 302 Gasoline, 301 Gastric juice, 417, 426 Gay Lussac’s burette, 261 law, 44 Gelatin, 411 Gelatinized starch, 350 Gelatinoids, 411 Gerard’s solution, 464 Germanium, 61 German silver, 187 Gin, 310 Glacial acetic acid, 324 phosphoric acid, 113 Glass, 162 Glauber’s salt, 142 Globulin, 405 Glucinum, 61 Glucose, 346 Glucosides, 352 Glue, 433 Gluten, 408 Glycerides, 342 Glycerin, 311 Glycerites, 311 Glycine, 354 Glycocoll, 354, 431 Glycocolic acid, 443 Glycogen, 351 Glycols, 305 Glycozone, 83 Glycyrrhizin, 352 Gmelin’s test, 430 Gold, 217 and sodium chloride, 218 analytical reactions of, 218 chloride, 218 coin, 218 sulphide, 218 Golden sulphuret of antimony, 215 Goulard’s extract, 326 Graham’s law, 36 Granite, 97, 159 Grape-sugar, 346 Graphic formulas, 283 Graphite, 92 Gravimetric methods, 256 Gravitation, 28 Green vitriol, 170 Gum, 350 -arabic, 350 British, 350 -resins, 371 Gun-cotton, 351 -metal, 187 Gunpowder, 137 Gutta-percha, 372 Gypsum, 152 Hair, 434 Hsematin, 409 Hsemato-crystallin, 409 Hsemine, 410 crystals, 472 Haemoglobin, 409 Halogens, 116 Haloids, 116 Hardness, 19 Heat, 24 action upon compounds, 53 matter, 21 organic substances, 288 decomposition by, 37, 55 latent, 25 specific, 27 Heavy magnesia, 149 spar, 157 Helleborin, 352 Hematite, 165 Hepar, 105, 137 Hippuric acid, 448 Homologous series, 285 Hoofs, 434 Hornblende, 159 Horns, 434 Humus, 420 Hydrargyrum, 195 INDEX. Hydrastine, 397 Hydriodic acid, 125 Hydrobromic acid, 123 Hydrocarbons, 96, 296 Hydrochloric acid, 119 analytical reactions of, 120 Hydrocyanic acid, 356 analytical reactions of, 358 antidotes to, 359 Hydroferricyanic acid, 360 Hydroferrocyanic acid, 360 Hydrofluoric acid, 126 Hydrogen, 78 arsenide, 206 arsenetted, 206 dioxide, 83 fluoride, 126 peroxide, 83 phosphide, 116 phosphoretted, 116 sulphide, 106, 234 sulphuretted, 106 Hydrometers, 29 Hydrosulphuric acid, 106 Hydroxyl, 284 Hygrine, 396 Hyoscine, 395 Hyoscyamine, 395 Hypobromites, 123 Hypochlorites, tests for, 122 Hypochlorous acid, 121 oxide, 121 Hvponitrous acid, 89 Hypophosphites, tests for, 115 Hypophosphorous acid, 115 Hyposulphurous acid, 106 ICHTHYOL, 369 I Illuminating gas, 302 oil, 302 Impurities, detection of, 248 Indestructibility, 36 India-rubber, 371 Indican, 352, 450 Indicators, 262 Indium, 61 Indol, 433 Inosite, 347 Iodides, analytical reactions of, 125 Iodimetry, 270 Iodine, 124 tests for it, 125 tincture of, 124 Iodized starch, 350 Iodoform, 318 Iodol, 381 Iridium, 61 Iron, 164 acetate, 326 analytical reactions of, 172 bromide, 169 carbonate, 171 cast-, 165 Iron chlorides, 168 citrate, 335 dialyzed, 169 galvanized, 180 hydroxides, 167 hypophosphite, 171 iodide, 169 lactate, 336 nitrate, 171 ores, 165 oxalates, 330 oxides, 166 phosphates, 171 pig, 165 pyrites, 165 reduced, 166 scale compounds of, 334 sulphates, 170 sulphide, 169 tannate, 172 tartrate, 334 trioxide, 168 wrought-, 165 Isomerism, 287 Isomorphism, 19 KAIRINE, 383 Kalium, 133 Kelp, 124 Keratin, 434 Ketones, 327 Kreatin, 434 T ACTIC acid, 335 lj Lactometer, 439 Lactoscope, 444 Lactose, 349 Lanolin, 344 Lanthanum, 61 Lapis infernalis, 193 lazuli, 162 Lardacein, 409 Latent heat. 25 Laudamine, 387 Laudanosine, 387 Laughing-gas, 88 Laurinol, 372 Law, Avogadro’s, 24 Charles’s, 25 of chemical combination by volume, 44 by weight, 42 of the conservation of energy, 36 of constancy of composition, 42 of diffusion of gases, 36 Dulong and Petit, 51 of equivalents, 45 Gay Lussac’s, 44 Graham’s, 36 Mariotte’s, 20 MendelejefPs, 62 of multiple proportions, 43 Newton’s 28 488 INDEX Law, periodic, 62 of Raoult, 52 Lead, 184 acetate, 326 analytical reactions of, 186 antidotes to, 186 carbonate, 185 chloride, 186 chromate, 186 iodide, 185 nitrate, 185 oleate, 344 oxide, 184 phosphate, 186 plaster, 344 sugar of, 326 -water, 326 white, 185 Lecithin, 432 Legumin, 407 Leucomaines, 402 Levulose, 347 Liebig’s condenser, 308 Light, decomposition by, 54 magnesia, 150 Lignine, 350 Lignite, 301 Lime, acid phosphate of, 154 chloride of, 155 chlorinated, 155 -kiln, 152 liniment, 344 quick-, 152 slaked, 153 superphosphate of, 154 -water, 153 Limestone, 152 Liniments, 344 Liquefaction of solids, 230 Liquids, definition of, 20 Litharge, 185 Lithium, 144 benzoate, 373 bromide, 144 carbonate, 144 citrate, 335 salicylate, 375 Litmus, 58 solution, 263 Lunar caustic, 193 Lutidine, 382 Lymph, 426 Magnesia, 149 calcined, 150 Magnesite, 149 Magnesium, 148 analytical reactions of, 151 carbonate, 149 citrate, 335 oxide, 150 sulphate, 150 sulphite, 151 Magnetic iron ore, 165 Malachite, 187 Malic acid, 331 Malleability, 20 Maltose, 349 Manganates, 174 Manganese, 173 analytical reactions of, 175 black oxide of, 173 dioxide, 173 oxides of, 173 Manganous carbonate, 175 hydroxide, 175 oxide, 173 sulphate, 174 Mannitose, 347 Marble, 152 Mariotte’s law, 20 Marsh-gas, 96, 299 Marsh’s test, 210 Mass, 17 -action, 56 Massicot, 185 Mastication, 417 Matter, definition of, 17 Mayer’s reagent, 385 Meconic acid, 390 Meconidine, 387 Meerschaum, 149 Melissic acid, 322 Melitose, 349 Melting-points of metals, 129 Mendelejeff’s law, 62 Menthol, 373 Mercaptans, 306 Mercurial ointment, 196 plaster, 196 Mercuric chloride, 198 cyanide, 358 fulminate, 362 iodide, 199 nitrate, 200 oxide, 196 salts, analytical reactions of, 201 sulphate, 199 sulphide, 200 Mercurous chloride, 197 chromate, 178 iodide, 198 nitrate, 200 oxide, 196 salts, analytical reactions of, 201 sulphate, 200 sulphide; 200 Mercury, 195 ammoniated, 201 analytical reactions of, 201 antidotes to, 202 basic sulphate, 199 carbonates, 201 chlorides, 197. 198 iodides, 198, 199 nitrates, 209 oleate, 328 INDEX. 489 Mercury, oxides, 196 sulphates, 199 sulphides, 200 with chalk, 196 Metaldehyde, 314 Metallic elements, 128 Metallo-cyanides, 360 Metalloids, 71 Metals, 128 classification of, 132 derivation of names, 128 melting-points, 129 separation of, 232 specific gravity, 130 valence, 131 Metamerism, 287 Metaphosphoric acid, 113 Methane, 96, 299 series, 298 Methyl alcohol, 307 amine, 400 hydride, 298 hydroxide, 307 orange, 263 Mica, 159 Microcosmic salt, 225 Milk, 435 adulterations of, 439 analysis of, 440 -casein, 438 of sulphur, 101 -sugar, 349 Millon’s reagent, 404 Mineral waters, 81 Minium, 185 Mint-camphor, 373 Mispickel, 204 Molasses, 348 Molecular motion, 24, 27 theory, 22 weight, 41, 51 Molecule, definition of, 22 Molybdenum, 219 Molybdic acid, 219 oxide, 219 Monobasic acids, 58 Monsel’s solution, 170 Morphine, 389 acetate, 389 analytical reactions of, 389 hydrochlorate, 389 sulphate, 389 Muscle-sugar, 347 Muscles, 434 Mucilage of starch, 270, 350 Mucin, 434 Mucus, 434 Murexid test, 448 Muriatic acid, 119 Myosin, 406 Myronic acid, 353 Myrosin, 353 Nails, 434 Naphthalene, 376 Naphthol, 377 beta-, 377 dinitro-, 377 Narceine, 390 [ Narcotine, 390 Nascent state, 57 Natrium, 140 Neutral substances, 58 Newton’s law, 28 Nickel, 179 Nicotine, 388 Niobium, 61 Nitrates, analytical reactions of, 91 Nitre, 136 Nitric acid, 89 Nitro-benzene+ 365 -cellulose, 351 -cyan-methane, 362 -glycerin, 312 Nitrogen, 84 determination, 280 oxides, 87 | Nitro-hydrochlorie acid, 120 -muriatic acid, 120 | Nitrous acid, 89 ether, 341 oxide, 88 ! Nomenclature, 66 I Non-metallic elements, 71 | Nordhausen sulphuric acid, 106 | Normal solutions, 261 | Nutrition of animals, 416 Nylander’s reagent, 462 OIL, almond, 343 bitter almond, 374 bone-, 381 castor, 343 cod-liver, 343 cotton-seed, 343 heavy, 364 illuminating, 302 juniper, 287 lemon, 287 light, 364 linseed, 343 olive, 343 turpentine, 371. vitriol, 103 wintergreen, 374 Oils, essential, 304 fat, 342 Olefiant gas, 304 Olefines, 303 Oleic acid, 328 Oleo-resins, 371 Olive oil, 343 Opium, 388 -alkaloids, 387 490 INDEX. Opium, denarcotized, 388 Organic analysis, 278 chemistry, 275 substances, classification of, 294 decomposition of, 288 formation of, in plants, 413 Orpiment, 204 Orthophosphoric acid, 113 Osmium, 61 Osmose, 35 Ossein, 411 Oxalates, analytical reactions of, 330 Oxalic acid, 329 antidotes to, 330 Oxide, definition of, 76 Oxidimetry, 267 Oxygen, 73 Ozone, 77 PALLADIUM, 61 1 Palmitic acid, 322 Palmitin, 342 Pancreatic juice, 432 Papaverine, 387 Paper, 351 Paraglobulin, 405 Paraffin, 302 Paraldehyde, 314’ Paris green, 326 Pearl-white, 191 Peat, 301 Pepsin, 410 saccharated, 410 Peptone, 408 Perchloric acid, 121 Periodic law, 62 Perissads, 47 Permanganates, 174 Petrolatum, 302 Petroleum, 300 -ether, 301 ointment, 302 Pettenkofer’s test, 431, 466 Phenacetine, 369 Phenetidin, 369 Phenol, 367 methyl-propyl, 373 phthalein, 263, 375 -sulphonic acid, 368 trinitro-, 369 Phenyl-acetamide, 380 -amine, 378 -hydrazine, 380 salicylate, 375 Phosphates, analytical reactions of, 114 Phosphine, 116 Phosphites, analytical reactions of, 112 Phospho-molybdic acid, 385 Phosplioretted hydrogen, 116 Phosphoric acid, 113 Phosphorous acid, 112 Phosphorus, 108 antidotes to, 111 Phosphorus, detection of, 111 determination in organic compounds, 281 red or amorphous, 110 Phthalic acid, 375 Physiological chemistry, 412 Physostigmine, 387 Picoline, 382 Picric acid, 369 Pilocarpine, 387 Pioskop, 440 Piperine, 387 Pipettes, 259 Plant-fibre, 350 -food, 412 Plaster-of-Paris, 153 Platinic chloride, 219 Platinum, 218 and ammonium chloride, 219 and potassium chloride, 219 Plumbago, 92 Plumbum, 184 Polymerism, 287 Polymorphism, 19 Porcelain, 162 Porosity, 32 Port wine, 310 Porter, 310 Potash, 134 caustic, 135 Potassium, 133 acetate, 325 acid carbonate, 136 acid oxalate, 330 acid tartrate, 332 analytical reactions of, 139 bicarbonate, 136 bichromate, 176 bitartrate, 332 bromide, 138 carbonate, 135 chlorate, 137 chromate, 176 citrate, 335 cyanate, 360 cyanide, 358 dichromate, 176 ferricyanide, 361 ferrocyanide, 360 hydrate, 135 hydroxide, 135 hypophosphite, 138 iodide, 138 manganate, 174 nitrate, 136 oxalate, 330 permanganate, 174 prussiate, 360 sodium tartrate, 333 sulphate, 137 sulphite, 137 sulphocyanate, 360 sulphurated, 137 tartrate, 332 INDEX. 491 Preliminary examination, 226 table for, 231 Proof-spirit, 309 Propionic acid, 322 Propyl alcohol, 310 Proteids, 403 bacterial, 402 Protopine, 387 Prussian blue, 361 Prussiate of potash, red, 361 yellow, 360 Prussic acid, 356 Pseudo-morphine, 387 Ptomaines, 398 Ptyalin, 426 Putrefaction, 290 Pyridine, 382 Pyrites, copper, 187 iron, 165 Pyrogallol, 376 Pyrolusite, 173 Pyrophosphoric acid, 113 Pyroxylin, 351 Pyrrole, 381 QUANTIVALENCE, 45 Quartz, 97 Quicksilver, 195 Quinidine, 393 Quinine, 391 acid sulphate, 392 analytical reactions of, 392 citrate of iron and, 392 sulphate, 392 valerianate, 328 Quinizine, 380 Quinoline, 382 RADICAL, definition of, 60, 284 Reactions, 57 analytical, 58 synthetical, 58 Reagents, list of, 224 Realgar, 204 Red iodide of mercury, 199 lead, 185 oxide of copper, 187 oxide of mercury, 197 phosphorus, 110 precipitate, 197 Reinsch’s test, 209 Residue, definition of, 60, 284 Resin, 371 Resorcin, 370 Respiration, 94, 419 Rhodium, 61 Rochelle salt, 333 Rock-crystal, 97 -salt, 140 Rosaniline, 379 Rosin., 372 Rosolic acid, 263 Rubber, 372 Rubidium, 61 Ruby, 159 Rum, 310 Ruthenium, 61 O ACCHARINE, 381 ■ 0 Saccharose, 348 Salicin, 374 Salicylic acid, 374 Saliva, 426 j Salol, 375 . I Sal-ammoniac, 145 sodse, 141 Salt cake, 141 common, 140 : Saltpetre, 136 Chili, 143 Salts, definition of, 59 tables of solubility, 246, 247 Sand, 97 Santonin, 377 j Sapphire, 159 Sarkin, 434 | Scale compounds of iron, 334 Scamonium, 352 Scandium, 61 | Scheele’s green, 208 Schweinfurth’s green, 209, 326 Seidlitz powder, 333 | Selenium, 108 ! Serpentine, 149 Serum, 424 Sherry wine, 310 Silica, 97 Silicates, 97 Silicic acid, 98 Silicium, 97 Silicon, 97 Silver, 192. analytical reactions of, 194 antidotes to, 194 chloride, 195 chromate, 195 cyanide, 358 fulminate, 362 German, 187 iodide, 194 nitrate, 193 oxide, 194 sulphide, 194 volumetric solution, 271 Sinapine, 387 Sinigrin, 353 Skatol, 433 Slaked lime, 153 Slate, 159 Soap, 343 Soapstone, 149 Soda, 140 -ash, 141 -lime, 279 Sodium, 140 492 INDEX. Sodium acetate, 325 analytical reactions of, 144 arsenate, 206 benzoate, 373 bicarbonate, 142 bisulphite, 142 borate, 144 bromide, 144 carbonate, 141 chlorate, 144 chloride, 140 hydrate, 140 hydroxide, 140 hypophosphite, 144 hyposulphite, 143 iodide, 144 nitrate, 143 phosphate, 143 salicylate, 375 sulphate, 142 sulphite, 142 sulphocarbolate, 369 thiosulphate, 143 Solanine, 387 Solids, definition of, 18 Specific heat, 27, 50 weight, 29 Spirit of hartshorn, 87 of wine, 309 proof, 309 rectified, 309 wood-, 307 Standard solutions, 261 Stannic acid, 217 chloride, 217 sulphide, 217 Stannous chloride, 217 hydroxide, 217 sulphide, 217 Stannum, 216 Starch, 349 iodized, 350 solution, 270 Stearic acid, 322 Stearin, 342 Stearoptenes, 372 Steel, 165 Steapsin, 410 Stibium, 213 Stibnite, 213 Strontianite, 156 Strontium, 156 analytical reactions of, 156 bromide, 156 carbonate, 156 chloride, 156 nitrate, 156 sulphate, 156 Structure of flame, 97 Strychnine, 393 analytical reactions of, 394 sulphate, 393 Sublimation, 26 Sublimed sulphur, 100 Substitution, 286 Sugar, 348 cane-, 348 detection of, in urine, 459 fruit-, 347 grape-, 346 of lead, 326 milk-, 349 Sulphates, analytical reactions of, 105 Sulphides, analytical reactions of, 107 Sulphites, analytical reactions of, 103 Sulphocarbolates, 369 Sulphocyanic acid, 359 Sulphonal, 318 Sulphonic acid, 319, 369 Sulphur, 100 determination in organic compounds, 281 dioxide, 101 flowers of, 100 milk of, 101 precipitated, 101 sublimed, 100 trioxide, 103 Sulphurated antimony, 214 lime, 155 Sulphuretted hydrogen, 106, 234 j Sulphuric acid, 103 antidotes to, 105 dilute, 105 fuming, 106 Nordhausen, 106 j Sulphuric ether, 339 Sulphurous acid, 102 Superphosphate of lime, 154 j Surface-action, 32 Sweet spirit of nitre, 341 Symbols, function of, 41 Synthesis, 57 Syntonin, 406 T'ABLES of solubility, 246, 247 I Talc, 149 Tannic acid, 376 Tannin, 376 Tantalum, 61 Tartar, 434 cream of, 332 crude, 331 emetic, 333 Tartaric acid, 331 Tartrates, analytical reactions of, 332 I Taurine, 431 ! Taurocholic acid, 431 Teeth, 434 Tellurium, 108 ; Tenacity, 20 ; Tension, 20 Terebene, 371 Terpenes, 370 Tetanine, 401 | Thalline, 383 I Thallium, 61 INDEX. 493 Thebaine, 387 Theine, 398 Theobromine, 387 Thermometers, 27 Thorium, 61 Thymol, 373 Tin, 216 -amalgam, 217 analytical reactions of, 217 chlorides of, 217 -plate, 217 -stone, 216 Titanium, 61 Titration, 263 Titre, 263 Toluene, 364 Toxalbumins, 402 Toxines, 400 Trichloraldehyde, 314 Trichlormethane, 315 Trinitro cellulose, 351 -phenol, 369 Triple phosphate, 470 Trommer’s test, 460 Trypsin, 418 Tungsten, 61 Turpentine, 371 Turpeth mineral, 199 Type-metal, 214 Types, chemical, 286 Tyrotoxicon, 401 11LTI MATE analysis, 279 U Ultramarine, 162 Ultzmann’s test, 466 Uranium, 61 Urates, 468 Urea, 443 determination of, 446 nitrate, 435 Uric acid, 447 Urinary calculi, 468 sediments, 467, 469 Urine, 442 analysis, 449 color, 450 composition, 443 reaction, 451 secretion, 442 specific gravity, 452 Urinometer, 452 Urobilin, 450 Uroxanthin, 450 WALENCE, 45 » Valerianates, 327 Valerianic acid, 327 Vanadium, 61 Vaseline, 302 Veratrine, 397 analytical reactions of, 397 oleate, 328 Verdigris, 326 Vermilion, 201 Vinegar, 325 Vitellin, 406 Vitriol, blue, 188 green, 170 oil of, 103 white, 181 Volatile oils, 304 Volumetric analysis, 259 Vulcanite, 372 Vulcanized rubber, 371 WASTE products of animal life, 420 Water, 81 distilled, 82 drinking, 81 -gas, 96 lead-, 326 lime-, 153 mineral, 81 of ammonia, 87 of bitter almond, 374 Wax, 338 Weight, 29 atomic, 40 specific, 29 molecular, 41, 51 Whey, 437 Whiskey, 310 White arsenic, 205 lead, 185 precipitate, 201 vitriol, 181 Wine, 310 Witherite, 157 Wood-naphtha, 307 -spirit, 307 Wrought-iron, 165 V ANTHIN, 434 A Xylene, 364 YELLOW oxide of mercury, 197 prussiate, 360 subsulphate of mercury, 199 -wash, 197 Ytterbium, 61 Yttrium, 61 y INC, 179 Lj acetate, 325 analytical reactions of, 182 antidotes to, 181 -blende, 179 bromide, 180 carbonate, 181 chloride, 180 ferrocyanide, 182 hydroxide, 180 iodide, 180 oxide, 180 phosphide, 181 sulphate, 181 valerianate, 328 -white, 180 Zirconium, 61