QUALITATIVE CHEMICAL ANALYSIS: A GUIDE IN THE PRACTICAL STUDY OF CHEMISTRY AND IN THE WORK OF ANALYSIS. BY ALBERT B. PRESCOTT, Ph.D, AND OTIS O. JOHNSON, M.A. FOURTH FULLY REVISED EDITION, WITH DESCRIPTIVE CHEMISTRY EXTENDED THROUGHOUT NEW YORK: D. VAN NOSTRAND COMPANY, 23 MURRAY AY I) 27 WARREN STS. 1895. Entered according to Act of Congress, in the year 1888, by W. 11. FARRINGTON, In the Office of the Librarian of Congress, at Washington. PREFACE TO THE FOURTH REVISED EDITION. In this edition the text has been mostly rewritten, to bring in the later results and to agree with fuller experience. The metalloids and acids have received more adequate treatment, and throughout the work the descriptive chemistry has been enlarged, in systematic arrangement. To gain room for these additions, some verbal condensation has been made all through the book, and such matter as could best be spared, including the introductory chapter, has been omitted. With a design to keep within pievious limits of size, the edition is only 13 pages larger than the previous one. In the preface to the first edition it was said to be the chief object in the work to aid the student in gaining accurate acquaintance with the facts whereby analyses aie made, and a clear understanding of the co-ordination of these facts the principles of ana lysis.” And the authors still believe that the best study of analytical chemistiy is among the best of the methods for the study of general chemistry. In this edition, with the purpose to bring together varied resources of analysis for co-operation with each other, to a wider knowledge of chemical reactions, the chief plans for quantitative analysis have been introduced. These are in no case to be resorted to as directions foi quanti tative operations, which require a strict control of details not presented in this woik. With the same purpose, under a like limitation, the leading methods of preparation 01 of manufacture of compounds are indicated, in order following the statements of “sources in chemical description. The special treatment of reactions of oxidation and reduction was introduced by one of the present authors in the third revised edition, and is now extended through the text upon metalloids and acids. The use of negative as well as positive units of valence has been found most helpful for the statement of reactions of oxidation, and is now offered, in its proper notation, where required throughout Part 11., but this additional notation will not in the least embarrass readers who would omit it. The first edition of this work appeared in 1873, the second revised edition in 18/6, and the third in 1880. With the first edition were incorporated the Tables for Qualitative Analysis by Professor Silas H. Douglas, editions of which had appeared in 1864, 1860, and 1868. Professor Douglas having some years since retired from engagements in chemistry, and being by his long and most valuable services in chemical education well entitled to release, now withdraws his name from the authorship of this work. The Chemical Laboratory op the University op Michigan, August, 1888. CONTENTS. INTRODUCTORY: page The Notation of Metallic Compounds 9 Table of Atomic Weights. 14 Order of Study 15 PART I. THE METALS : Classification of the Metals 16 Group V.: The Alkali Metals 17 Potassium 19 Sodium 21 Ammonium , 22 Lithium 24= Rubidium and Caesium 26 Group IV.; The Alkaline Earth Metals 29 Magnesium 27 Barium 30 Strontium 83 Calcium 34 Separation of the Fourth-Group Metals 36 Separation of Magnesium from the Alkali Metals 37 Group III.: The Metals of the Earths and the more Electro-positive of the Heavy Metals 38 Aluminium 40 Chromium as a Base 43 Chromic Acid 45 Manganese 54 Iron 46 Cobalt 59 Nickel 63 Zinc 65 Comparison of Reactions of Third-Group Bases 69 Separation by Ammonium Hydroxide, etc., Table 70 Separations of the Metals of the Third Group 72 Separation of Phosphates, with Alkali Acetate, Table 75 Cerium, Beryllium 77 Uranium 78 Titanium 79 Thallium 80 Ytterbium, Terbium, Scandium, Samarium, Gallium, Decipium 81 Comparative View of the Ten Remaining Metals of Group 111. 82 5 6 Contents. Groups I. and 11. Metals whose Sulphides are Insoluble in Dilute Acids 83 Copper 84 Bismuth 90 Cadmium 94 Comparison of Bismuth, Copper, and Cadmium 95 Lead 96 Silver 101 Mercury ... 105 Mercurous Salts 107 Mercuric Salts 109 Comparison of First-Group Metals 114 Arsenic 115 Antimony 127 Tin 135 Comparison of Aisenic, Antimony, and Tin 141 Separation of the Metals of the First and Second Groups 141 Methods for separating the Bases without the use of Sulphides 151 Gold 153 Platinum 156 Palladium 158 Molybdenum 160 Germanium, Norwegium 163 Ruthenium, Iridium, Rhodium, Osmium 164 Tellurium, Selenium, Tungsten 164 PART 11. THE NON-METALS AND THE ACIDS : Valence and Negative Bonds 165 Oxidation Valence compared with Structural Valence 166 Rule for Balancing Equations 169 Hydrogen 171 Boron, Boric Acid 173 Carbon 175 Acetic Acid 176 Citric Acid 177 Tartaric Acid 178 Carbon Monoxide 179 Oxalic Acid 180 Carbonic Anhydride 183 Cyanogen, Hydrocyanic Acid 187 Ferrocyanic Acid 191 Ferricyanic Acid 193 Cyanic Acid 195 Thiocyanic Acid 196 Nitrogen 198 Nitrous and Nitric Oxides 200 Nitrous Acid 200 Nitric Acid 303 Nitrogen Peroxide 203 Oxygen 309 Ozone 310 Contents. 7 Hydrogen Peroxide 211 Fluorine 212 Hydrofluoric Acid 213 Silicon 214 Silicic Acid 215 Phosphorus 216 Phosphorus Hydride .' 218 Hypophosphorous Acid 218 Phosphorous Acid 221 Phosphoric Acid 221 Sulphur 225 Hydrosulphuric Acid 226 Thiosulphuric Acid 230 Dithionous Acid 232 Dithionic, Trithionic, and Tetrathionic Acids.... 233 Pentathionic Acid 234 Table Comparing the Thionic Series of Acids 235 Sulphurous Acid 236 Sulphuric Acid 240 Chlorine 244 Hydrochloric Acid 247 Hypochlorous Acid 251 Chlorous Acid 253 Chlorine Peroxide 253 Chloric Acid 253 Perchloric Acid 257 Bromine 257 Hydrobromic Acid 250 Hypobromous Acid 262 Bromic Acid 263 lodine 264 Hydriodic Acid 267 lodic Acid 271 Periodic Acid 273 Comparative View of Reactions of the Acids of Chlorine, Bromine, and lodine 274 part hi. systematic examinations : Separation of the Acids of Chlorine, Bromine, and lodine 275 Separation of the Acids from the Bases and from each other 280 Conversion of Solids into Liquids 381 Separation from Organic Matter 282 Preliminary Examinations of Solids 283 Systematic Examinations in the Dry Way 287 Tests in Beads of Borax and Phosphorus Salt 291 Analysis for the Metals, Tables 394 Analysis for the Acid Radicals, Tables 302 Solubilities of Salts 308 General Table of Solubilities 311 Reagents used in Analysis 31g> THE NOTATION OF METALLIC COMPOUNDS. ACIDS. An acid is a salt of hydrogen. It consists of an Acid Radical,* united with hydrogen which can be exchanged for a metal, this being the forma-, tion of a salt. The hydrogen is the base of the acid, as the metal is the base of the salt. Sulphuric acid, for instance, is sometimes written hydro- gen sulphate. Oxacids are those whose radicals contain oxygen, as HNO,. Hydracids are those whose radicals have no oxygen, as HCI; their names begin with hydr and end with ic, and the names of their salts end in ide. bhe Anhydride of an oxacid is what remains after removing from the acid its basic H, and enough Oto form H2O with the H, Thus, the anhydride °f H2SQ4 is S03, and carbonic anhydride is C02, carbonic acid being Acids whose molecules contain but one atom of basal hydrogen are termed monobasic, as HNOs and HCI; those with two atoms of hydro- gen in the molecule, dibasic, as H2S04 and H2S ; those with three hydro- gen atoms, trihasic, as h3po4 ; etc. Some of the more important acids are given in the following list : HC1, liydrocliloric acid, hydrogen chloride, or h ydric chloride. hcio3, chloric 6 6 66 chlorate. 6 6 chlorate. HBr, hydrobromic 6 6 66 bromide. 6 6 bromide. HBrO„ bromic < 6 66 bromate, 6 6 bromate. HI, hydriodic 66 6 6 iodide, 6 6 iodide. HIO„ iodic 6 6 66 iodate, 6 6 iodate. hno., nitric 6 6 6 6 nitrate, 6 6 nitrate. H.S, aAo., H.SO., H,CO., hydrosulphuric 66 6 6 sulphide, 6 6 sulphide. sulphuric 6 6 6 6 sulphate. 66 sulphate. sulphurousf 66 66 sulphite, f 6 6 sulphite, f carbonic 66 66 carbonate, 6 6 carbonate. a,Cro., BPo * 15 h.Aso„ chromic 6 6 66 chromate, 66 chromate. phosphoric 66 66 phosphate, 6 6 phosphate. arsenic cc 66 arseniate, 66 arseniate. a^e(CN)., >. aA°.. arsenious 66 66 arsenite, « arsen i te. hydroferricyanic 66 66 ferricyanide, 66 ferricyanide. hydroferrocyanic 66 66 ferrocyanide. 66 » ferrocyanide. acetic oxalic 66 66 66 66 acetate, oxalate, 66 66 acetate. oxalate. onp mnJ10 *erm Radical is applied to a group of atoms which retains its integrity while transferred from m°lecuie to another. (ren ,^XaC'^S " ose names en 00 QO fe L. Meyer and K. Seubert, 1882. Aluminium.. . . A1 27.009 27.04 Nickel. Ni 57 928 58 6 Antimony Sb 119.P55 119.6 Niobium Nb 93 812 93.7 Arsenic As 74 918 74.9 Nitrogen N 14.0210 14.01 Barium Ba 130.703 136.86 Norwegium j Beryllium. .. . 9.085 9.08 Osmium Os 198.494' 195. Bismuth Bi 207 523 207.5 Oxygen o 15 9633 15.96 Boron B 10 941 10.9 Palladium Pd 105 737 106.2 Bromine. . . . • Br 79 768 79 76 Phosphorus P 30 958 30.96 Cadmium. . Cd 111 835 111.7 Platinum Pt 194.415 194.3 Caesium Cs 132.583 132.7 Potassium K 39.019 39.03 Calcium Ca 39 990 39.91 Rhodium. . Rh 104.055 104.1 Carbon..... o 11 9736 11.97 Rubidium Rb 85.251 85.2 Cerium Ce 140 424 141 2 Ruthenium Ru 104.217 103.5 Chlorine... . Cl 35 370 35 37 Samarium Sm 149.801 Chromium. . Or 52 009 52 45 Scandium Sc 43 980 43.97 Cobalt Co 58 887 58 6 Selenium Se 78.797 78.87 Copper Ou 63.173 63.18 Silicon Si 28.195 28.0 Decipium Dp * Silver As 107.675 107.66 Didyrnium Di 142.121 145 0 Sodium 22.998 22.995 Erbium Er 165.891 166. Strontium Sr 87.374 87.3 Fluorine F 18.984 19 06 Sulphur s 31.984 31 98 Gallium Ga 68.854 69.9 Tantalum Ta 182.144 182. Germanium... Gr 72.32f Tellurium Te 127.960 127.7 Gold Au 196 155 196 2 Terbium Tr § H 1 000 1 Thallium T1 203.715 203.7 In 113 398 113 4 Thorium Th 232.020 231.96 I 126 557 126 54 Tin Sn 117 698 117 35 Iridium Ir 192 651 192.5 Titanium Ti 47.980 50 25 1 ron Fe 55 913 55 88 Tungsten W 183 610 183 6 1 .lanthanum La 138.019 138.5 Uranium u 238 482 239 8 Lead Pb 206.471 206.39 Vanadium V 51 256 51.1 Lithium Li 7 0073 7 01 Ytterbium Yb 172 761 172 6 Magnesium Ms 23 959 23 94 Y 88 900 89 6 Manganese Mn 54 855 54 8 Zinc Zn 64 9045 64.88 Mercury Hg 199.7J2 199.8 Zirconium Zr 89.367 90.4 Molybdenum Mo 95.527 95.9 The discovery of a few other elements has recently been announced. It is hoped that more accurate investigations will soon place their existence beyond doubt. For list of new elements announced see a report, by H. 0. Bolton, “ Trans. N. Y. Acad. Sci- ences,” Yol. V. * Probably 171, Delafontaine. X Probably 218.93, Dahll. t Winkler, 1886. § 148.5, Marignac. .4 Course in Qualitative Analysis. ORDER OP STUDY. The order of study of qualitative analysis, in the laboratory under the authors’ charge, varied from year *° year, is at this time about as follows : Preparatory—a drill in “writing salts,’’ to memorize quantivalence and make the notation familiar. Then, FIRST, a study of the solubilities of metallic salts and hydrates— namely . A, obtaining all the precipitates by potassium hydrate or sodium hydrate, with the metals successive groups, then the same with ammonium hydrate, the student writing equations for all pre- Staining the precipitates by potassium carbonate, with the successive bases, and formu- thi*11" i Ul_l.'lan«es- In the same way the students work out the precipitates with, O, the sulphates; I), II common88 ' e° sulpl:iurous acid; E, ammonium sulphide ;G, hydrosulphuric acid ; iodides,°” s.oclium Phosphate; I, free phosphoric acid ;J, chlorides; K, bromides ;X, the anal r i 10(^a^es 1 iV, potassium dichromate ;O, ammonium oxalate. SECOND, a study of other fit- a react,'ons °I each base, and then, for the first, practice in the separation of metals from each t^om *n t,le or(ler of their groups. THIRD, the analytical reactions of each acid, and then the P'oducts>nS -°f aC.i(*s' FOURTH, practice with synthetic operations, devised by the student, for required niake lead" ".'VCTI materials—equations of all changes being given byr the student. Thus, required to Point a LSU,!)Me’ the lead from the metallic state, and the sulphur from calcium sulphate. At this further 'ixamination is held, and qualification upon all the work passed over is required before going The c >.• ■ o|c analysis of unknown solid mixtures, each containing from two to seven compounds, mixture matlon base, in the greater number, to be determined by the action of solvents upon the SIXTH th Cl)orts received after analyses of each ten, and results of first and second reports preserved. Lastly afi G |inalySlS °f mixt"reB in solution, mostly involving the action of oxidizing and reducing agents, nal examination. There is a daily recitation, with the daily laboratory work. PAET I—THE METALS. 1. Classification of Metals or Bases.—ln chemical analysis, the metals are commonly divided into five groups, according Lo their deportment, in solution of their salts, with certain general reagents, as follows : I. Those metals forming chlorides insoluble in water are precipitated from the solutions of their salts by the first group reagent, hydrochloric acid: Pb, Ag, (Hg2)". 11. Metals which are not precipitated by the first group reagent (hydro- chloric acid), but are precipitated by the second group reagent, which is hydrosulphuric acid slightly acidulated with hydrochloric acid: Sn, Sb, As, Hg", Bi, Cu, Cd, Pb. Lead chloride is slightly soluble in water ; hence only a part of the lead is precipitated in the first group, and the remainder in the second. 111. Metals which are not precipitated by the first or second group re- agents, but by ammonium hydroxide and sulphide in presence of ammo- Mum chloride: Fe, Cr, Al, Co, Ni, Mn, Zn. IV. Those metals which are not precipitated by the first, second, or third group reagents, but by ammonium carbonate in presence of ammo- nium hydroxide and ammonium chloride : Ba, Sr, Ca. Y. Metals which are not precipitated by the reagents of the preceding groups : Mg, K, Na, NH(. Only the more common metals are here given. The position of the rarer metals in the five groups is given later. 2. The group reagents of the second group will precipitate the metals of the first group. The group reagents of the third group will precipitate the metals of the first two groups, except Sn, Sb, and As, which are soluble in excess of ammonium sulphide. The fourth group reagents will precipi- tate the metals of the first three groups, except Ag, Cu, As, Cd, Co, Ni, Zn, Mn. After filtering out a group precipitate, the reagent which produced it should be added to the filtrate to insure its entire removal, before testing for the next group. 16 Tut: Metals of the Alkalies. 17 GROUP V. 3. Magnesium, Mg" = 23.959 An Alkaline Earth Metal. The Alkali Metals. Potassium, . . K' = 39.019 Lithium, . . Li' = 7.0073 Sodium, . . . Ka' = 22.998 Rubidium, . . Rb' = = 85.251 Ammonium, . • W Caesium, . Os' = 132.583 4. The metals of the alkalies are highly combustible, oxidizing quickly in the air, displacing the hydrogen of water even more rapidly than zinc °r iron displaces the hydrogen of acids, and displacing non-alkali metals from their oxides and salts. As elements they are very strong reducing agents, while their compounds are very stable, and not liable to either re- duction or oxidation by ordinary means. The five metals, Cs, Rb, K, Na, frij present a gradation of electro-positive or basic power, caesium being strongest, and the others decreasing in the order of their atomic weights, lithium decomposing water with less violence than the others. Their speci- fic gravities decrease,* their fusing points rise, and as carbonates their sol- ubilities lessen, in the same order. In solubility of the phosphate, also, lith- -111111 approaches the character of an alkaline earth. Ammonium is the basal radical of ammonium salts, and as such has Uiany of the characteristics of an alkali metal. The water solution of the ammonia, NH3 (an anhydride), from analogy is supposed to contain ammonium hydrate, NH4OH, known as the volatile alkali. Potassium and s°dium hydrates are the fixed alkalies in common use. 5. The alkalies are very soluble in water, and all the important salts of le Alkali metals (including NHj are soluble in water, not excepting their Caibonates, phosphates (except lithium), and silicates ; while dll other Petals form hydrates or oxides, either insoluble or sparingly soluble, and Caidonates, phosphates, silicates, and certain other salts quite insoluble in Water. I heir compounds being nearly all soluble, the alkali metals are 7iot pre- printed by ordinary reagents, and, with few exceptions, their salts do not Precipitate each other. In analysis, they are mostly separated from other letals by non-precipitation. 6- In accordance with the insolubility in water of the non-alkali hy- -1 eg and oxides, the alkali hydrates precipitate all non-alkali metals, ex- c P ammonium hydrate does not precipitate barium, strontium, and ,lUnu These precipitates are hydrates, except those of mercury, silver, 1(1 antimony. ♦Except those of potassium (0.865) and sodium (0.973). 18 The Metals of the Alkalies. But certain of the non-alkali hydrates and oxides, though insoluble in water, dissolve in solutions of alkalies ; hence, when added in excess, the alkalies redissolve the precipitates they at first produce with salts of certain metals, viz.; the hydrates of Pb, Sn, Sb, Zn, Al, and Cr dissolve in the fixed alkalies ; and oxide of Ag and hydroxides of Cu, Zn, Co, and Ni dis- solve in the volatile alkali. Precipitations by Alkali Hydroxides (KOH, NaOH, and NH4OH). Note the color of the precipitates. For a full statement of the composition of the precipitates, as obtained under different conditions, see the text for the several metals. At certain stages in the addition of alkalies, basic salts of numerous metals are formed. The proportion of hydroxyl is reduced by elevation of temperature, short of that of boil- ing water, in many instances; thus Cu(OH)2 is changed to Cu3Oj(OH)2, and at 100° C. to CuO. Barium hydrate, Ba(OH)2, not caused by NH4OII, sol, in 20 parts water. Strontium 66 Sr(OH)2, “ “ “ 60 “ Calcium 66 Ca(OH)2, o o Magnesium 66 Mg(OH)2, soluble by NH4C1, sol. in 6,000 “ Aluminium 66 Al2(OH)6, soluble in excess of fixed alkali hydrates. Chromium 66 Cr2(OH)6, soluble in cold sol. of fixed alkali, precipitat- ed on boiling. Ferrous 66 Fe(OH)2, slightly soluble by JMII4C1. Oxidizes in air. Ferric 66 Fes(OH)r Manganous 66 Mn(OH)2, soluble by NH Cl. Oxidizes in air. Manganic 66 Mn,,(OH)( Cobalt 66 Co(OH)2, soluble in excess of NH4OH, and by NH401. Nickel 66 Ni(OH)a, 66 66 6( (C Zinc 66 Zn(OH)2, soluble in both fixed and volatile alkalies. Copper 66 Cu(OH)2, soluble in NH4OH (with bine color). Cadmium 66 Cd(OH)2, “ “ (colorless). Bismuth 66 Bi(OH)3. Lead 66 Pb(OH)2, soluble in excess fixed alkalies. Silver oxide, Ag„0, soluble in excess NH4OH. Mercurous oxide, Hg,2C (by fixed alkalies). Mercurous-ammonium chloride, NH2Hg2Cl, from Hg2Cl2 by NH4OH. P OTASSIUM. 19 Mercuric oxide, HgO (by fixed alkalies). Mercur-ammonium chloride, NHaHgCl, from HgCI2 by NH4OH. Antirnonious oxide, Sb4Oe, soluble in excess fixed alkalies. Stannous hydrate, Sn(OH)a, “ “ “ Stannic “ Sn(OH) , “ “ u rt- Solutions of the alkalies are caustic to the taste and touch, and turn le(l litmus blue ; also, the carbonates, acid carbonates, normal phosphates, ai)d some other salts of the alkali metals, give the “alkaline reaction” test papers. Sodium nitroferricyanide, with hydrogen sulphide, £['es a delicate reaction. S- The hydrates and normal carbonates of the alkali metals are not de- composed by heat alone (as are those of other metals), and these metals 01111 the only acid carbonates obtained in the solid state. 9; -the fixed alkalies, likewise many of their salts, melt on platinum 11 111 the flame, and slowly vaporize at a bright red heat (distinction from magnesia). All salts of ammonium, by a careful evaporation of their solu- -118 011 platinum foil, may be obtained in a solid residue, which rapidly Vtlpoiizes, wholly or partly, below a red heat (distinction from fixed alkali metah and magnesium). hydrates of the fixed alkali metals, and those of their salts most vo.atile at a red heat, preferably their chlorides, impart strongly character- ic colors to a non-luminous flame, and give well-defined spectra with the spectroscope. POTASSIUM. K = 39.019. H. Specific gravity 0.865 (Gay-Lussac). Fuses at 63.5° C. (Bunsen). 01 s at 719°-731° C. (Carnelley). 12. Preparation.—From the carbonate by fusion with carbon. Also t^a 6 % electrolysis of the hydroxide. A white metal, brittle at 0° 0., above wlii *'emPera^ure s°fl like wax. It forms two oxides—the peroxide, K204, UIS n° conesPonc^ng salts, and potassium oxide, KnO. The so- xoY) *SU^°X^e a mixture tlie oxide and metallic potassium (Lup- -13. The hydrate, carbonate, dimetallic phosphate, sulphite, nitrite, ace- e, and normal tartrate are deliquescent. tu 14 7ne Potassium salts are quite insoluble in water ; the pla- Per ° l(^C"’ tartrate, silico-fluoride, picrate, phosphomolybdate, and eohol °mte’ belng only slig]ltly soluble in water, and quite insoluble in ab 0 • Ihe chlorate is but sparingly soluble in cold water, and nearly * Jour. Chem. Soc., 1876; S, 565. 20 Potassium. insoluble in alcohol. Also, the carbonate and sulphate are insoluble in al- cohol. 15. In ordinary qualitative analysis, potassium compounds are identified by their flame-color, using blue glass to exclude the color of the sodium- flame ; also by precipitation of potassium acid tartrate in alcoholic acidified solution, or of potassium platinic chloride. Both these precipitations are used in quantitative analysis. 16. Platinic Chloride (PtCl4), added to neutral or acid solutions not too dilute, with hydrochloric acid if the compound be not a chloride, preci- pitates potassium platinic chloride, (KCI)„PtGI4, crystalline, yellow. Non- alkali bases also precipitate this reagent, and if present must be removed before this test. The precipitate is soluble in 19 parts of boiling water, or 111 parts of water at 10° 0. Minute proportions are detected by evaporat- ing the solution with the reagent nearly to dryness, on the water-bath, and then dissolving in alcohol; the yellow crystalline precipitate, octahedral, re- mains undissolved, and may be identified under the microscope. 17. Tartaric acid (H2C4H406), or more readily sodium hydrogen tartrate (NaH C 4H406), precipitates, from solutions sufficiently concentrated, potas- sium hydrogen tartrate, KHC4H4Oe, granular-crystalline. If the solution be alkaline, tartaric acid should be added to strong acid reaction. The test must be made in absence of non-alkalin bases. The precipitate is increased by agitation, and by addition of alcohol. It is dissolved by fifteen parts of boiling water or eighty-nine parts water at 25° C., by mineral acids, by solution of borax, and by alkalies, which form the more soluble normal tartrate, K204H4O6, but not by acetic acid, or at all by alcohol of fifty per cent. 18. Picric acid, HC6H2(N02)30, precipitates, from solutions not very dilute, the yellow, crystalline potassium pier ate, KC6H2(N02)sO, insoluble in alcohol, by help of which it is formed in dilute solutions. The dried pre- cipitate detonates strongly when heated. 19. It will be observed that ammonium salts form precipitates with plati- nic chloride and with tartaric acid, closely resembling those formed with salts of potassium, but the latter is the only fixed alkali which is precipi- tated by these reagents. 20. Potassium compounds color the flame violet. A little of the solid substance, or residue by evaporation, moistened with hydrochloric acid, is brought on a platinum wire into a non-luminous flame. The wire should be previously moistened with hydrochloric acid, and held in the flame until it does not color. The presence of very small quantities of sodium enables its yellow flame completely to obscure the violet of potassium ; but owing to the greater volatility of the latter metal, flashes of violet are sometimes seen on the first introduction of the wire, or at the border of the flame, or in its base, even when enough sodium is present to conceal the violet at full beat. Sodium. Silicates may be fused with pure gypsum, giving vapor of potassium sul- phate. The interposition of a blue glass, or prism filled with indigo solu- tion, sufficiently thick, entirely cuts off the yellow light of sodium, and en- ables the potassium flame to be seen. The red rays of the lithium flame are also intercepted by tiie blue glass or indigo prism, a thicker stratum being Required than for sodium. If organic substances are present, giving lumi- nosity to the flame, they must be removed by ignition. Certain non-alkali bases interfere with the examination. 21. The volatile potassium compounds, when placed in the flame, give a Widely-extended continuous spectrum, containing two characteristic lines ; °ae line, K a, situated in the outermost red, and a second line, K (S, far iu the violet rays at the other end of the spectrum. 22. Oxidation.—Potassium does not oxidize in perfectly dry air. Ox- idizes rapidly in moist air, and when thrown upon water bursts into flame. % the aid of heat it reduces the oxides of carbon, phosphorus, and nitro- gen to the free elements, and similarly reduces to the metallic state all the °xides of lead, silver, mercury, tin, antimony, arsenic, bismuth, copper, cadmium, zinc, manganese, cobalt, nickel, iron, and the oxides of many of the rarer metals. SODIUM. Na = 32.998. 23. Specific gravity .973 (Gay-Lussac). Puses at 95.6° 0. (Bu^tseiv). Preparation.—(l) By igniting the carbonate with carbon ; (3) by ignit- lng 'he hydroxide with metallic iron ; (3) by electrolysis. 24. A white, soft metal, resembling potassium in its properties, and, being cheaper, is in greater demand. Sodium forms two oxides, Ha20 and a Peroxide, Na„02. 25. Of ordinary compounds of sodium, only the hydrate, nitrate, and chlorate are deliquescent. The carbonate (10 aq.), sulphate (10 aq.), sul- phite (8 aq.), phosphate (13 aq.), and acetate (3 aq.), are efflorescent. 26. The salts of sodium are freely soluble in water, except the metanti- ni°niate and the silico-fluoride, the latter being sparingly soluble. . Sodium is identified chiefly by its flame-color (30), and by non-pre- Clpitation with various reagents. Its soluble salts are weighed in gravi- nietric operations. 28. Solution of potassium metantimoniate (KSbOs) produces, in neutral °! ‘hhaline solutions, a slow-forming, white, crystalline precipitate, NaSbOa, a “jest insoluble in cold water. The reagent must be carefully prepared ancl dissolved when required, as it is not permanent in solution. (See Under Antimonic Acid.) Ce j29’ platinic chloride, (NaCl)2PtCl4, crystallizes from its con- -1 1 aled solutions in red prisms, or prismatic needles (distinction from P° assmm or ammonium). A drop of the solution to be tested is slightly Ammonium. acidified with hydrochloric acid from the point of a glass rod on a slip of glass, treated with two drops of solution of platinic chloride, left a short time for spontaneous evaporation and crystallization, and observed under the microscope. 30. Sodium compounds color the flame intensely yellow—the color being scarcely affected by potassium (at full heat), but modified to orange-red by much lithium, and readily intercepted by blue glass. Infusible compounds may be ignited with calcium sulphate. The test is interfered with by some non-alkali bases. 31. The spectrum of sodium consists of a single band, Na a, at Fraun- hofer’s line D, in the yellow of the solar spectrum. 32. The amount of sodium in the atmosphere, and in the larger num- ber of substances designed to be “ chemically pure,” is sufficient to give a distinct but evanescent yellow color to the flame and spectrum. 33. Oxidation.—Sodium, in its reducing power, is similar to potassium, but a little less rapid in its action. AMMONIUM. 34. Specific gravity of NH3 gas, 0.5901 (Davy); of the liquid, 0.6234 (Jolly). The liquid melts at —7s° 0. (Fakaday); boils at —38.5° 0. (Eegnault). 35. The anhydride, ammonia (NH3), gaseous at common temperatures, dissolves in twice its weight of cold water, forming a volatile solution lighter than water. 36. The “ sesquicarbonate,” (NH4)4H2(C03)3 (1 aq.), or tetra-ammo- nium dihydrogen tricarbonate, and the phosphate (2 aq.), are efflorescent ; the nitrate is deliquescent, and the sulphate slightly deliquescent. The normal carbonate is very instable, and used only in solution. 37. The solubilities of the salts of ammonium correspond very nearly tvitli those of potassium salts. Ammonium is found by obtaining the anhydride, ammonia, in vapor. Precipitation as mercur-ammonium iodide is also used. Ammonium pla- tinic chloride is weighed in quantitative work. For the nitroferricyanide test, see under Hydrosulphuric Acid. 38. Ammonia gas (NHJ escapes from its solutions (having alkaline re- action) at ordinary temperatures, more rapidly when heated ; and from its combinations, in any mixture (alkaline, neutral, or acid), by heating with an alkali or alkaline earth (potassium or calcium hydrate). I*ll,ol + KOH = KCI + NH3 + H2O 39. Ammonia gas is recognized, Ist, by its odor; 3d, by turning moistened red litmus-paper to blue ; 3d, by changing red logwood paper blue ; 4th, by rendering paper wet with solution of cupric sulphate blue j A MMONIUM. 23 sth, by blackening paper wet with solution of mercurous nitrate ; 6th, by forming white fumes with the vapors of volatile acids, vapor of HCI form- ing solid WH4CI; vapor of HC2H302 forming solid NH4 G2H302, etc. 40. A solution of potassium mercuric iodide, (KI)2Hgl2, containing also potassium hydrate—Nessler’s test *—produces a brown precipitate of nitrogen dimercuric iodide, NHgJ (dimercur-ammonium iodide—see under Mercury), soluble by excess of KI and by HCI; not soluble by KBr (dis- tinction from HgO): NH3 + 2HgI2 = NHg2I + 3HI NH4OH + 2(KI)2Hgl2 + 3KOH = NHgsI + 7KI + 4H2Q This very delicate test is applicable to ammonium hydrate or salts ; traces forming only a yellow to brown coloration. The potassium mercuric iodide, alone, precipitates the alkaloids from neutral or acid solutions, but does not precipitate ammonium salts from neutral or acid solutions. 41. Mercuric chloride (HgCIJ forms, in solutions of ammonium hy- drate or ammonium carbonate, the “white precipitate ” of nitrogen dihy- drogen mercuric chloride, WH2HgCI, or mercur-amrnonium chloride. If ammonium is in a salt, not carbonate, it is changed to the carbonate aftd precipitated, by addition of mercuric chloride and potassium carbonate previously mixed in solutions (with pure water), so dilute as not to preci- pitate each other (yellow). This test (Bohlig’s) is intensely delicate, re- vealing the presence of ammonia derived from the air by water and many substances. 42. Add a small quantity of recently precipitated and well-washed sil- ver chloride, and, if it does not dissolve after agitation, then add a little Potassium hydrate solution. The solution of the AgCl, before the addition the fixed alkali, indicates free ammonia; after the addition of the fixed alkali, ammonium salt. (Applicable in absence of thiosulphates; iodides, bromides, and sulphocyanides.) 43. Platinic chloride and tartaric acid form precipitates with am- monium, which, in conditions of production, form and color of crystals, and 111 solubility, closely resemble the potassium precipitates with the same re- cent. They may be distinguished by the effect of ignition, which, in case .ammonium platinic chloride, leaves pure spongy platinum (without and, in case of ammonium hydrogen tartrate, leaves pure carbon (without K„CO„). Also, NH,H CHO, is more soluble in water than A A / 7 4 446 . c4H406. Picric acid precipitates ammonium, in solutions not very dilute. dA. Phosphomolybdate of sodium precipitates ammonium from neu- Id,l or acid solutions ; also precipitates the alkaloids, even from very dilute si * Ttl'S rea-ent maY ho prepared as follows :To a solution of mercuric chloride add solution of potas- the U lod'de rile precipitate is nearly all redissolved ; then add solution of potassium hydrate ; leave until quid becomes clear, and decant from any remaining sediment. Lithium. solutions, and, from concentrated solutions, likewise precipitates K, Rb, and Cs (all the fixed alkalies except Na and Li). 45. Ammonium salts in solution, treated with chlorine gas, generate the instable and violently explosive “ nitrogen chloride” (NOI3?) (a). The same product is liable to arise from solid ammonium salts treated with chlorine. Gaseous ammonia, and aqueous am- monium hydrate, with chlorine gas, generate free nitrogen (b), a little ammonium chlo- rate being formed if the ammonia is in excess. Hypochlorites or hypohvomites (or chlo- rine or bromine dissolved in aqueous alkali, so as to leave an alkaline reaction) liberate, from dissolved ammonium salts, all of their nitrogen (as shown in the second equation of h)\ the measure of the nitrogen gas being a means of quantitative estimation of ammo- nium.—With iodine, ammonium iodide and the explosive iodamides (as in equation c) are produced ; also, in proportion governed by conditions, iodate (d) and hypoiodite may be formed.—Ammonia is liable to atmospheric oxidation to ammonium nitrite and ni- trate.—Permanganates oxidize to nitrate (e).—Ammonia is somewhat readily produced from nitric acid by strong reducing agents. It is formed with carbonic anhydride, in a water solution of Cyanic acid, and, more slowly, in a water solution of Hydrocyanic acid. It is generated, by fixed alkalies, in boiling solution of Cyanides!/); also, in boiling solutions of albuminoids and other nitrogenous organic compounds, this formation being hastened and increased by addition of permanganate (Wanklyn’s process). Fusion with fixed alkalies transforms all the nitrogen of organic bodies into ammonia. a. NH4CI + 8CI2 = NC13 + 4HCI b. 8NH3 + 3012 = CNH4CI + N2 3NH4CI + BCI2 = BHCI + N, c. 2NH3 -f I 2 = NH4I + NH2I d. fiNH4OH -f 3I„ = SNH4I + NH4I03 -J- 8H20 e. GNH4OH + 4H2Mns08 = BNH4NO3 + 8MnO(OH)2 -f 5H20 f. HCN -f KOH + H2O = NHS + ECHO, (formate) 46. Heat vaporizes the carbonate, and the haloid salts of ammonium, undecomposed; decomposes the nitrate with formation of nitrous oxide and water, the phosphate and borate with evolution of ammonia, and other salts with various products. 47. Ammonium compounds impart to the flame a faint and evanescent violet color. LITHIUM = 7.0073. Specific gravity, 0.5936. Melting point, 180° C. (356° F.) (Bunsen). 48. Occurrence.—lt is a sparingly but widely distributed metal. Usu- ally prepared from lepidolite, tripbylene, or petalite. Traces are found in a great many minerals, in mineral springs, and in the leaves and asbes of many plants—e.g., coffee, tobacco, and sugar-cane. 49. Preparation.—Salts of the metal were prepared in 1817, and in 1855 Bunsen and Matthiessen isolated it in considerable quantities by electrolysis. It is the lightest metal known ; harder than potassium and sodium, but .softer than lead. It is volatile at a red heat in an atmosphere of hydrogen ; Lithium. 25 oxidizes in moist but not in dr}7, cold air. When thrown upon water, rapid- ly forms the hydroxide, but does not burst into a flame like sodium and po- tassium. 50. Oxide and Hydroxide.—lt forms one oxide (Li2o) by heating the ttietal in oxygen or dry air; cheaper by the action of heat upon the nitrate. ilhe corresponding hydroxide (LiOH) is made by the action of water upon the metal or its oxide ; cheaper by heating the carbonate with calcium hydroxide. 51. Solubilities.—The chloride, chlorate, and many other salts, are very deliquescent. The carbonate, phosphate, and silico-fluoride are only spar- 11]gly soluble in water ; the other salts of lithium are freely soluble in water, and nearly all soluble in alcohol. 52. Reactions of Lithium Salts.—Lithium salts are more fusible and hiore easily decomposed by fusion than the corresponding potassium and sodium salts. 53. Sodium phosphate, Ha2HP04, precipitates trimetallic lithium phos- phate, Li3P04, soluble in 2,530 parts water ; more soluble in solutions of amrnonium salts (distinction from magnesium), but much less soluble in strong solution of ammonia. In dilute solutions the precipitate forms only atter boiling; and addition of sodium hydroxide to alkaline reaction in- creases the delicacy of the test, forming a double phosphate of sodium and hthium (0. Rammelsburg : Ann. Phys. Chem. [2], 7, 157). Its solution 111 hydrochloric acid is not at once precipitated by ammonium hydrate in the cold (distinction from alkaline earth metals); and the blow-pipe bead of lithi mn phosphate, with soda, is transparent (that of alkaline earth Petals being opaque). 54. Nitrophenic acid forms a yellow precipitate, not easily soluble in 'Water. 55. Compounds of lithium impart to the flame a carmine-mZ color, ob- scured by sodium, but not by small quantities of potassium compounds. Ue ghiss, just thick enough to cut off the yellow light of sodium, trans- mits the red light of lithium ; but the latter is intercepted by a thicker l)ait of the blue prism, or by several plates of blue glass. _ -I'he spectrum of lithium consists of a bright red band, Li a, and a aint orange line, Li fd. The color tests have an intensity intermediate be- een those of sodium and potassium. ,5e- Estimation.—After separation from other elements it may be Weighed as a sulphate, carbonate, or phosphate (Li3PO4). It may also be , "Bated by the comparative intensity of the lines in the spectroscope Cell : Amer. Chem. Jour., 7, 35). 57- Oxidation.—When heated it burns in Cl, Br, I, S, and in carbon hlQxide. 26 R ÜBIDI UM— CAESIUM. RUBIDIUM. Rb' = 85.251. 58. Specific gravity, 1.52 (Bunsen). Melting point, 38.5° 0. (101.3° F.) (Bunsen). 59. Occurrence.—Widely distributed in minute quantities ; found in lepidolite ; in the ashes of some plants; in certain mineral springs, from which source it was first obtained. 60. Preparation.—By electrolysis of the chloride, also by heating the carbonate with carbon. 61. Properties.—-A soft, wax-like metal, with a yellowish tint; inflames when exposed to the air, but less readily than cassium. It bursts into a flame when thrown upon water; burns when brought in contact with gaseous Cl, Br, I, S, and As. 62. Oxide and Hydroxide.—Rb2Q is supposed to be formed when ru- bidium burns in oxygen, but it is not yet proven. The corresponding hydroxide (RbOH) is formed when rubidium is oxidized by water, or more cheaply by treating its carbonate with calcium hydroxide. 63. Reactions of Rubidium Salts.—lts salts are nearly all soluble. Pla- tinic chloride (PtCl4) precipitates PtRb2Cl6, soluble in 157 parts of water at 100° 0., or 649 parts at 10° C. Tartaric acid precipitates rubidium hydrogen tartrate (RbHO4H4O6), soluble in 84 parts of water at 25° 0. The spec- trum shows two characteristic lines in the violet, Rb a and Rb (3, also two others in the red, Rb d and Rb y. The spectrum reaction is so delicate that 0.002 rngrm. may be detected (Bunsen). CAESIUM. Cs' = 132.583. 64. Specific gravity, 1.88 (Setteeberg). Melting point, 26° to 27° C. (Setterberg). 65. Occurrence.—Found in the mineral pollux, and in the water of some springs. 66. Preparation.—By electrolysis of the caesium and barium cyan- ides. Its reduction by carbon has not yet been accomplished (Setter- berg). 67. Properties.—The metal is silver-white, and soft at ordinary tem- peratures ; oxidizes quickly in the air, and takes fire when thrown on water (Setteeberg). 68. Oxide and Hydroxide.—The oxide formed by burning caesium in oxygen is supposed to be Cs20. The corresponding hydroxide is formed when caesium is oxidized by water, but prepared cheaper by treating the carbonate with calcium hydroxide or the sulphate with barium hydrox- ide. Magnesium. 27 69. Reactions of Caesium Salts.—Nearly all caesium salts are soluble. Platinic chloride precipitates Cs2PtCl4 in octahedral crystals, soluble in 265 parts of water at 100° C., and in 2,000 parts at 10° C. Stannic chloride precipitates Cs2SnCl6, insoluble in strong hydrochloric acid, but soluble in Water. Antimonous chloride precipitates CsSbCl4, completely separating ]t from all other alkalies (Crooke’s Select Methods, page 26). Its spectrum consists of two bright blue lines almost coincident with the strontium blue line. MAGNESIUM. Mg" = 23.959. 70. Specific gravity, 1.75 (Deville). Melting point, about 500° 0. (932° F.) (Ditte). Volatilizes at about 1100° C. (2507° F.) (Ditte). 71. Occurrence.—Magnesite (MgCOj, dolomite (CaMg(C03)2), brucite 2), epsom salts (MgS04), and combined with other metals in a great Variety of minerals. 72. Preparation.—(l) By electrolysis of the chloride. (2) By heating the chloride with potassium or sodium. 73. Properties.—A white, hard, malleable, and ductile metal; not acted Upon by water or fixed alkali hydroxides at ordinary temperatures, and only ®hghtly at 100° C. Soluble in acids and in ammonium chloride. (4NH4CI + Mg = (NH4)2MgCl4 + 2NH3 + Ha) 74. Oxide and Hydroxide.—Only one oxide of magnesium (MgO) 18 known with certainty. Formed by burning tbe metal in the air, and by action of beat upon the hydroxide, carbonate, nitrate, sulphate, oxalate, and °*-her magnesium salts decomposable by heat. Tbe corresponding hydrox- ]de (Mg(OH)2) is formed by precipitating magnesium salts with the fixed alkalies. 75. Solubilities.—The chloride, bromide, iodide, chlorate, nitrate, and a°et;ite (4 aq.) are deliquescent; the sulphate (7 aq.), slightly efflorescent. The hydroxide, carbonate, phosphate, and arseniate are insoluble in AVa,er ; the sulphite, oxalate, and tartrate, sparingly soluble ; the chromate, s°kible. The hydrate and carbonate are soluble in ammonium salts—ex- Cel)t ammonium phosphate. Reactions of Magnesium Salts.—The fixed alkali hydrates and 10 hydrates of barium, strontium, and calcium, precipitate, from magne- Slllm salts in solution, magnesium hydroxide, Mg(OH),,, nearly insoluble in "‘her, but soluble in ammonium chloride or sulphate (equation a). . hydroxide precipitates half the magnesium as a hydrox- e> leaving the other half in solution as a double salt of magnesium and ammonium {h): a. Mg(OH)2 + 4NH4CI = (NH4CI)2MgCl2 + 2NH4OH i. 2MgS04 + 2NH4OH = Mg(OH)2 + (NH^SO.MgSO^ 28 Magnesium. 77. Ammonium sulphide forms no precipitate. Tim normal carbon- ates of the fixed alkali metals—as K,C03—precipitate magnesium basic car- bonate, Mg4(C03)3(OH)2, variable to MgS(C03)4(OH)2. Carbonic acid is liberated in the formation of this basic salt: But in the cold the free C02 combines with another portion of MgCOc to form a soluble supercarbonate : 4MgS04 + 4Na2COs + H2O = Mg4(C03)3(OH)2 + C02 + 4Na2S04 On boiling, the supercarbonate is precipitated as MgC03 with escape of co2. 78. Ammonium carbonate scarcely precipitates magnesium salts, ex- cept in concentrated solutions, owing to the formation of a soluble double carbonate of magnesium and ammonium: SMgSOi + 5Na2C03 + HaO =r Mg4(C03)3(OH)2 + MgC03.C02 + SNa2S04 79. Alkaline phosphates—as Na2HP04—precipitate magnesium phos- phate, MgHP04, if the solution be not very dilute. But even in very di- lute solutions, by the further addition of ammonium hydroxide (and NH4CI), a crystalline precipitate is slowly formed, magnesium ammonium phosphate—Mgl4H4P04. Stirring with a glass rod against the side of the test-tube promotes the precipitation. The addition of ammonium chloride, in this test, prevents formation of any precipitate of magnesium hydrate (76 a). The precipitate dissolves in 15,000 parts pure water, or in 44,000 parts of water containing ammonium hydrate. 80. Alkaline arseniates—as Na2HAs04—act with magnesium salts in all respects like the phosphates, giving corresponding precipitates. 81. In the dry way, the only characteristic test for magnesium is the pale rose color, obtained by igniting, then moistening the compound with solution of cobalt nitrate, and again igniting strongly on charcoal. The color is more apparent on cooling, is not intense, and is jirevented by pres- ence of many other bases. The spectrum of magnesium, as well as the spectra of most of the metals yet to be described, cannot be obtained by means of the flame, in which their compounds are not volatile. To obtain them, recourse must be had to the electric spark. 82. Estimation.—After removal of other metals magnesium is preci- pitated as MgNH4PG4, then changed by ignition to Mg2P2G, (magnesium pyrophosphate) and weighed as such. 83. Oxidation.—Burns brilliantly in the air, giving a light of high ac- tinic power useful in photography. It precipitates the free metals from solutions of Sb, Mn, Bi, Pe, Zn, Cd, Tl, Sn, Pb, Te, Co, Ni, Cu, Ag, Pt, Pd, An. MgS04 + 2(NH4)a003 = MgC03(NH4)2C03 + (NH4),SO4 Alkaline Earth Metals. 29 GROUP IV. 84. Alkaline Earth Metals.—Barium, Ba" = 136.763. Strontium, Sr" = 87.374. Calcium, Ca" = 39.99. 85. Comparative View.—Magnesium belongs to the fifth group, and is not usually classified as an alkaline earth ; hut on account of its close re- semblance it is discussed in the following paragraphs. 86. Like the alkali metals, Ba, Sr, and Ca oxidize rapidly in the air at ordinary temperatures—forming alkaline earths—and decompose ivater with- out the aid of an acid, forming hydroxides; also these hydroxides are formed, with evolution of heat, when the oxides are brought in contact with water. Mg oxidizes rapidly in the air when ignited, decomposes water at 100° C., and its oxide—in physical properties unlike tfie alkaline earths slowly unites with water without sensible production of heat. As com- pounds, these metals are not easily oxidized beyond their quantivalence as (lyads, and they require very strong reducing agents to restore them to the elemental state. 87. In basic power, Ba is the strongest of the four, Sr somewhat stronger than Ca, and Mg much weaker than the other three. It will be observed that the solubility of their hydroxides varies in the same decreasing grada- tion, which is also that of their atomic weights ; while the solubility of their sulphates varies in a reverse order, as follows ; 88. The hydroxide of Ba dissolves in about 20 parts of water ; that of r> in 60 parts ; of Ca, in 700 parts; and of Mg, in 6,000 parts. The sul- phate ol Ba is not appreciably soluble in water ; that of Sr dissolves in 7.000 parts ; of Ca, in 400 parts ; of Mg, in 3 parts. To the extent in which they dissolve in water, alkaline earths render their solutions caustic t° tlm taste and touch, and alkaline to test-papers. 89. The carbonates, normal phosphates, silicates, and some other salts °f these four metals, are insoluble in water (as are those of the bases of 16 first three groups). Magnesium carbonate is soluble iu ammonium Salts, whereby its precipitation with the other three is prevented. Cal- c}H7n oxalate and barium chromate are insoluble (see table for Group •): the oxalates of barium, strontium, and magnesium, and the chro- niate of strontium, are sparingly soluble ; chromate of calcium freely sol- uble. . I'l qualitative analysis, the group-separation of the fourth-group metals 8 effected, after removal of the first three groups of bases, by precipitation 'Vlth carbonate in presence of ammonium chloride, after which magnesium 18 precipitated from the filtrate, as phosphate. The hydroxides of Ca, Sr, and Ba, in their saturated solutions, jjecessarily dilute, throw down, from solutions of salts of the metals of the lst three groups and of Mg, thin precipitates of hydrates of the latter 30 Barium. which precipitates are not soluble in excess of the precipitants. In turn, the fixed alkalies precipitate, from solutions of Ba, Sr, Ca, and Mg, so much of the hydroxides of these metals as does not dissolve in the water present; but ammonium hydroxide precipitates only Mg, and this but in part, owing to the solubility of Mg(OH)2 in ammonium salts. 91. Solutions containing Ba, Sr, Ca, and Mg, with phosphoric, oxalic, boracic, or arsenic acid, necessarily have the acid reaction, as occurs in dis- solving phosphates, oxalates, etc., with acids; such solutions are precipi- tated by ammonium hydroxide or by any agent which neutralizes the solu- tion, and, consequently, we have precipitates of this kind in the third group. CaCl2 + H3P04 + 2NH4OH = CaHP04 + 2WH4CI + 3H20 CaH4(P04)2 + 3WH4OH = CaHP04 + (NH4)2HPOt + 2H20 If excess of the ammonium hydroxide be added the precipitate is Ca3(p°4)2- In the case of a magnesium salt the precipitai e is MgNH4P04, 92. The cart)o7iates of the alkaline earth metals are dissociated by heat leaving metallic oxides and carbonic anhydride. This occurs with difficulty in the case of Ba, Sr, and Ca; witli readiness in the case of Mg; hence ignition of the carbonates of Ba, Sr, mid Ca causes them to present the alkaline reaction to a slip of moistened litmus-paper. 93. Compounds of Ba, Sr, and Ca (preferably with HCI) impart charac- teristic colors to the non-luminous flame, and readily present well-defined spectra. BARIUM. Ba" = 136.763. 94. Specific gravity, 3.75 (Keeist). Melting-point, 475° C. (887° F.) (Van dee Weyde). Not volatile at a red heat. 95. Occurrence—Found chiefly in heavy spar (BaSOj and witherite (BaCOj. 96. Preparation.—(l) By electrolysis. (2) Reduction by potassium or sodium. 97. Properties.—A silver-ivhite (Dayy) or yellowish-white (BuisrsEJsr) metal; ductile, malleable ; oxidizes rapidly in the air and in water. 98. Oxides and Hydroxide.—The oxide. BaO, is formed by the action of heat upon the hydroxide, carbonate, nitrate, oxalate, and all its organic salts. The corresponding hydroxide, Ba(OH)2, is made by treating the oxide with water ; is soluble in 20 parts of water at 15° C., and in 2 parts at 100° 0. The peroxide (BaOJ is made by heating the oxide almost to redness in oxygen, or air which has been freed from carbon dioxide : by heat- ing the oxide with potassium chlorate (Liebig) or cupric oxide (Wank- i,yx). Barium. 31 It is used as a source of oxygen, which it gives off at a white heat, BaO remaining ; also in the manufacture of hydrogen peroxide, H202, which is formed by treating it with dilute acids. (Bao2 + 2HCI = BaCl2 + H202) 99. Solubilities.—Most of the soluble salts of barium are permanent; the acetate is efflorescent. The chloride, bromide, iodide, sulphides, ferrocyanides, nitrate, chlorate, acetate, and phenylsutphate, are freely soluble in water; the carbonate, sulphate, sulphite, chromate, phosphate, oxalate, iodate, and silico-fluoride, Jtre insoluble in water. The chloride is almost insoluble in strong hydro- chloric acid ; likewise the nitrate in strong hydrochloric and nitric acids. tlhe chloride and nitrate are insoluble in alcohol. 100. Reactions of Barium Salts.—Barium may be separated from other alkaline earth metals by precipitation as chromate (105), and by its closer precipitation as sulphate (104). The latter precipitation is a sharp distinction from all other metals except lead, strontium, and calcium, and 18 the operation most used in quantitative analysis of barium and of sul- phates. 101. The fixed alkali hydroxides precipitate only concentrated solutions °1 barium salts, as explained by the statement in 88. 102. The alkali carbonates—as K2C03 and (NH4)„CO3—precipitate, h°m barium salts in solution, barium carbonate (BaCOj, white. The pre- cipitation is promoted by heat and by ammonium hydroxide, but is made lightly incomplete by the presence of ammonium chloride and nitrate. 103. Barium Carbonate- BaCOa—is a valuable reagent for special purposes, chiefly 0r separation of third-group metals. It is used in the form of the moist precipitate, rbich must be thoroughly washed. It is best precipitated from boiling solutions of chlo- -I]Je of barium and carbonate of sodium or ammonium, washed once or twice by decanta- n, then by filti’ation, till the washings no longer precipitate solution of nitrate of sil- Ver- Mixed with water to consistence of cream, it may be preserved for some time in Coppered bottles, being shaken whenever required for use. When dissolved in hydro- c acid, and fully precipitated by sulphuric acid, the filtrate must yield no fixed residue. . 1 his reagent removes sulphuric acid (radical) from all sulphates in solution to which is added (104). Na2SO, + BaCOg = BaSC4 + Na,C03 When salts of non alkali metals are so decomposed, of course, they are left insoluble, as carbonates or hydrates, nothing remaining in solution: FeS04 + BaCOs = BaSO., + FeCG3 Fe2(S04)a + 38aC03 + 3H.0 = 38a504 + Fe2(OH)c + 8C02 I’he chlorides of the double triads of the third group, namely, aluminic, chromic, and Jerric chlorides, are decomposed by barium carbonate; while the other metals of the u’q group, zinc, manganese, cobalt, nickel, and iron in ferrous combination, are not Precipitated from their chlorides by this reagent. But tartaric acid, citric acid, sugar, dll(l other organic substances, prevent the decompositions by carbonate of barium. 32 Barium. 104. Sulphuric acid (H2SQ4), and all soluble sulphates, precipitate ba- rium sulphate (BaSOj, white [BB], slightly soluble in hot concentrated sul- phuric acid. Immediate precipitation by the (dilute) saturated solution of calcium sulphate distinguishes Ba from Sr (and of course from Ca); but precipitation by the (very dilute) solution of strontium sulphate is a more certain test between Ba and Sr. 105. Normal chromates, as K2Cr04, precipitate barium salts (also, strontium salts in solutions not very dilute); the yellow precipitate, BaCr04, being almost insoluble in water, slightly soluble in acetic acid, but soluble in hydrochloric and nitric acids, and moderately soluble in chromic acid. (SrCro4, also yellow, is a little more soluble in water than the barium salt.) Dichromates (as K„Cr20.) precipitate barium, as normal chromate, from the acetate, in solution not dilute (but do not precipitate strontium). 106. Soluble phosphates, full metallic, or two-thirds metallic, as Na2HP04, precipitate barium phosphate, white, consisting of BaHPQ4 when the reagent is two-thirds metallic, and Ba3(P04)2 when the reagent is full metallic. 107. Oxalates, as (NH4)20204, precipitate barium from solutions not very dilute ; as BaC„04, somewhat soluble in oxalic and acetic acids. 108. Hydro-fluosilicic acid, H„SiF6, precipitates white, crystalline Ba SiFc, slightly soluble in water, not soluble in alcohol (distinction from strontium and calcium). 109. Solutions of iodates, as Nal03, precipitate, from barium solutions not very dilute, barium iodate, Ba(I03)2, white, soluble in 600 parts of hot or 1,746 parts of cold water (distinction from the other alkaline earth metals). 110. Barium compounds impart to the flame a yellowish green color, which appears blue-green when viewed through green glass. 111. The spectrum of barium is at once distinguished from all others by the green bands, Ba a, Ba ft; Bay is less distinct but more char acteristic. 112. Estimation.—Barium is weighed as a sulphate, carbonate, or a sili- co-fluoride (BaSiFj. It is separated from strontium and calcium : (1) By digesting the mixed sulphates at ordinary temperatures for 12 hours with ammonium carbonate. The calcium and strontium are thus converted into carbonates, which are separated from the barium sulphate by dissolving in hydrochloric acid. (2) By hydro-fluosilicic acid. The hydrate and carbon- ate are also determined by alkalimetry. One part of a volumetric process recommended depends upon the separation by potassium dichromate in ex- cess of ammonium hydroxide (Crooke’s Select Methods, page 45). Strontium. 33 STRONTIUM. Sr" = 87.374. 113. Specific gravity, 2.4 (Franz). Melts at a red heat (Franz}. 114. Occurrence.—Strontium occurs chiefly in strontianite (SrUOJ and in celestine (SrSOj. 115. Preparation.—First isolated in 1808 by Davy. Is made by elec- trolysis ; also by fusion of the chloride with sodium amalgam. 116. Properties.—A white (Davy) or faintly yeilow metal (Franz). ft is malleable and ductile; somewhat harder than lead; oxidizes rapidly 111 air and water. 117. Oxides and Hydroxide.—The oxide (SrO) is formed by burning tile metal in the air ; also by igniting the hydroxide, carbonate, nitrate, °Xalate, and all organic strontium salts. Tne hydroxide is formed by ac- ti°n of water on the oxide. The peroxide, Sr02, is made by treating the hydroxide with hydrogen peroxide (Sr(OH)2 -(- H202 = Sr02 + 2H20). It °annot be made, like the corresponding barium peroxide, by heating the hy- droxide with oxygen or potassium chlorate (Brodie). 118. Solubilities.—The chloride is slightly deliquescent; crystals of the nitrate and acetate effloresce. j fn solubility most compounds of strontium closely resemble those of drium (99)—the hydrate being a little less soluble, and the sulphate and chrornate more soluble, in water than the corresponding barium compounds, atld the silico-fluoride quite soluble (see 88). The chloride is soluble, the 11Trate insoluble, in alcohol absolute. 119. Reactions of Strontium Salts.—The fixed alkalies precipitate, from c°n centra ted solutions, strontium hydroxide (Sr(OH)2), soluble in 60 parts of 'Water. Strontium is identified, in the fourth group, after removal of barium, by Precipitation with calcium sulphate solution (121) ; also, quite clearly, by le flame-color and spectrum (122, 123). °ntiuni is not to be distinguished from barium ; the differing reactions Or fi o z o aji lle two metals with sulphates, chromates, and hydro-fluosilieic acid under the head of Barium, Strontium sulphate is soluble in *° sbo parts of concentrated nitric or hydrochloric acid. „ . Saturated solution of calcium sulphate (CaSO4) slowly produces a chi1 .h1‘ecd)itate of SrS04, prevented or dissolved by presence of hydro- aol lc anCentlated So^u^'oll ammonium sulphate, which separates it from 4* Strontium compounds color the flame crimson. In presence of IPO crimson color appears at the moment when the substance, moist- With hydrochloric acid, is first brought into the flame. The paler. 34 Calcium. yellowish-red flame of calcium is liable to be mistaken for the strontium- flame . 123. The spectrum of strontium is characterized by eight bright bands —namely, six red, one orange, and one blue. The orange line, Sr a, at the red end of the spectrum ; the two red lines, Sr ft and Sr y, and the blue line, Sr d, are the most important. 124. Estimation.—Strontium is weighed as a sulphate or a carbonate. The carbonate may be determined by alkalimetry. It is separated from calcium by the insolubility of its sulphate in ammonium sulphate. It is separated from barium as stated in 112. CALCIUM. Ca" == 39.99. -125. Specific gravity, 1.57 (Bunsen). Melts at a bright red heat (Matthiessen). 126. Occurrence.—Found in the mineral kingdom as a carbonate in marble, limestone, chalk, and arragonite ; as a sulphate in gypsum, selen- ite, alabaster, etc.; as a fluoride in fluor-spar; as a phosphate in apatite, phosphorite, etc. It is found as a phosphate in bones ; in egg-shells and oyster-Shells as a carbonate. It is found in nearly all spring and river waters. 127. Preparation.—By electrolysis ; also by fusion with sodium amal- gam or with an alloy of sodium and zinc. 128. Properties-—Calcium is a pale yellow metal, softer than zinc, but harder than tin and lead. It is malleable and ductile ; may be kept for several days in dry air without oxidation, but rapidly oxidizes in moist air and in water. 129. Oxides and Hydroxide.—The oxide, CaO, is formed by oxidation of the metal in air ; by ignition of the hydroxide, the carbonate (limestone), nitrate, oxalate, and all organic acids. Its usefulness when combined with sand, making mortar, is too well known to need any description here. The corresponding hydroxide, Ca(OH)2 (slaked lime), is made by treating the oxide with water. The peroxide, CaO„, is made by adding hydrogen perox- ide to the hydroxide (Ca(OH)2 + H202 = Ca02 + 2H20). It cannot be made by heating the oxide in oxygen or with potassium chlorate (Brodie). 130. Solubilities,—The chloride, bromide, iodide, nitrate, and chlorate are deliquescent; the acetate is efflorescent. 131. The carbonate, oxalate, and phosphate are insoluble in water ; ihe hydrate, sulphate, sulphite, and iodate are slightly soluble in water (88), but are insoluble in alcohol. The chloride, iodide, and nitrate are soluble in alcohol. The ferrocyanide is soluble, the potassio-ferrocyanide insolu- ble, in water. 132. Reactions of Calcium Salts.—The fixed alkali hydrates precipi- Calcium. tate calcium hydrate, Ca(OH)2, from solutions of calcium salts not very dilute. The precipitate is less soluble in solution of potassium or sodium hydrate, and more soluble in solution of ammonium hydrate than in pure Water. 133. In their deportment with soluble carbonates (precipitation of CaC03) and with alkaline phosphates (precipitation of CaHP04 or Ca,(Po4)s), solutions of calcium cannot be distinguished from solutions of strontium and barium (102, 106). 134. Sulphuric acid and soluble sulphates (not calcium sulphate) pre- cipitate CaS04 from calcium salts, in moderately concentrated solutions, h’he precipitate is distinguished from barium and strontium sulphates by dissolving in concentrated solution of ammonium sulphate. 135. Alkaline oxalates, as (NH4)2C„O4, precipitate calcium oxalate, CaC204, from even dilute solutions of calcium salts. The precipitate is scarcely at all soluble in acetic or oxalic acids (separation of oxalic from phosphoric acid), but is soluble in hydrochloric and nitric acids. The pre- cipitation is hastened by presence of ammonium hydrate. Formed slowly, Ifoni very dilute solutions, the precipitate is crystalline, octahedral. If Sr ol'Ba are possibly present in the solution tested, an alkaline sulphate must fir8t be added, and after digesting a few minutes, if a precipitate ap- pears, SrS04, BaS04, or, if the solution was concentrated, perhaps CaS04, it 18 filtered out, and the oxalate then added to the filtrate. Observe the pre- cipitate formed by ammonium oxalate in the reagent solution of calci- Urn sulphate. Ignition of Ca02O4 changes it first to CaC03, then to CaO, giving alkaline reaction to test-paper, 136. Neutral alkaline sulphites, as Na2S03, precipitate CaS03, nearly Hisoluble in water, soluble in hydrochloric or nitric acid, and in sulphurous acid. This reaction is common to the alkaline earths, 137. Alkaline arsenites precipitate, from neutral calcium solutions, calcium arsenite, CaHAs03, soluble in acids and in ammonium hydrate. Ihe precipitate forms slowly. Other alkaline earth metals are not precipi- tated by arsenites, unless in concentrated solutions, 138. An ammoniacal solution of arsenious acid gives a precipitate of calcium arsenite in neutral calcium salts. Under similar circumstances ba- llllrn and strontium give no precipitate. t-39. Compounds of calcium, preferably the chloride, render the flame yellowish red. The presence of strontium or barium obscures this reaction, . a mixture containing calcium and barium, moistened with hydrochloric dcid, gives the calcium color on its first introduction to the flame. 140. The spectrum of calcium is distinguished by the bright green line, a P, and the intensely bright orange line, Ca a, near the red end of the spectrum. l4=l. Estimation.—Calcium is weighed as an oxide, carbonate, or sul- Separation of the Fourth Group Metals. pluite. The carbonate is obtained by precipitating as oxalate, and gently igniting the dried precipitate. The sulphate is precipitated in a mixture of two parts of alcohol to one of the solution. The best method of separation from strontium, is to treat the nitrates, with a mixture of equal volumes of alcohol and ether. The calcium nitrate dissolves, but not more than one part in 60,000 of the strontium is found in the solution. For other methods of separation from barium and strontium see 113 and 134. SEPARATION OP THE FOURTH-GROUP METALS, 142. Barium, strontium, calcium, and magnesium may be completely precipitated together, either as carbonates or as phosphates ; but a precipitate of phosphates would be intractable in further operations, owing to the difficulty of removing the non-volatile phosphoric acid. Hence, they are precipitated as carbonates, and this could be done by any alkaline carbonate; but the necessity for subsequent examination for fixed alkali metals restricts us to ammonium carbonate. Now, this reagent but imperfectly preci- piiates magnesium, and from this difficulty, and also because magnesium is more easily separated from alkali metals than from metals of the fourth group, the ordinary scheme of separation provides for the precipitation of Ba, Sr, and Ca, by ammonium carbonate in presence of ammonium chloride, so as to leave Mg either with the fifth group, or as a distinct division of the fourth group. 143. The precipitation of barium, strontium, and calcium by ammonium carbonate in the presence of chloride, is not as complete as would be desirable in very delicate analyses. For the carbonates of barium, strontium, and calcium are all slightly soluble in ammonium chloride solution; and while the prescribed addition of ammonium hydrate, and excess of ammonium carbonate, greatly reduces the solubility of the precipitated carbonates, yet even with these the precipitation is not absolute, though more nearly so with strontium than with barium and calcium. Thus, in quantitative analyses, if barium and calcium are precipitated as carbonates, it must be done in the absence of ammonium chloride or sulphate, and the precipitate washed with water containing ammonium hydrate. 144. But a more accurate precipitation of barium is effected by sulphates, and of calcium by oxalates, and these tests may be applied to portions of the filtrate from the precipitation by carbonates, or of the liquid that has given no precipitate by carbonates. Also, the complete removal of barium and calcium is not only a test for traces of these two metals, but it enables us to accept a slight precipitation of phosphate afterwards as conclusive evidence of the presence of magnesium (unless lithium be present). This pre- cautionary work, done after the ordinary work for barium, strontium, and calcium, may be tabulated as follows : Divide the filtrate from the fourth group into three portions. Test in I. for Ba with a drop of H2S04, leav- | Test in II. for Ca with (NH,)2C204, leaving ing some time. I some time. If both Ba and Ca appear, mix I. and II.; let the mixture stand ; filter and test the filtrate for Mg' by Na.,HPO4 and NH4OH. If either Ba or Ca appears, filter it and test the filtrate for Mg. If neither Ba nor Ca appears, test portion 111. for Mg. 145. The solution of calcium sulphate can be used to distinguish between barium, strontium, and calcium, provided that but one metal of the group is present, and that the solution be at least moderately concentrated, and not notably acid. Separation of Magnesium from the Alkali Metals. 37 146. The unlike solubilities in alcohol, of the chlorides and nitrates of barium, strontium, and calcium enable us to separate them quite closely by absolute alcohol, and approximately by “ strong alcohol,” as follows: Dissolve the carbonate precipitate in HCI, evaporate to dryness on the water-bath, rub the residue to a fine powder in the evaporating dish, and digest it with alcohol. Fil- ter through a small filter, and wash with alcohol. Residue, BaCl2. Dissolve in water, test with CaS04, etc. Filtrate SrCl2 and CaCl2. Evaporate to dryness, dissolve in water, change to nitrates,by pre- cipitating with (NHjiaCOj, washing, and dissolving in HNOj. Evaporate the nitrates to dryness, powder, digest with alcohol, filter and wash with ulcohi,). Residue, Sr(N03)2. Precipitation by CaS04 in wa- ter solution ; flame test, etc. Filtrate, Ca(N03)2, Precipitation by H2S04 in alco- hol solution;, by (NH4laC204,etc. 147. Or, the alcoholic filtrate of SrCl2 and OaCl2 may be precipitated with (a drop °f) sulphuric acid, the precipitate filtered out and digested with solution of (NH4)2SO4 aild a little NH4OH. Residue, SrSO4. Solution contains CaSO4, precipitable by Oxalates. SEPARATION of magnesium from the alkali metals. 148. By ignition on platinum foil, magnesium compounds do not vaporize, as do those ammonium, nor melt, as do many compounds of fixed alkalies. Magnesium is the only one of these metals precipitated by ordinary salts—viz., by phosphates, carbonates, and hydroxide. 149. The presence of magnesium slightly impairs the delicacy of the flame test for the fixed alkali metals, and entirely prevents their recognition or separation by precipita- tions. Phosphate of ammonium will remove magnesium from solution ; but, after eva- porating the filtrate and igniting its residue, the phosphoric acid remains—combined Wlth the fixed alkali metals, if they are present. Thus: (NH4).,HPO4 (excessof reagent), ignited = HP03 + 2NH3 + H2O, and 2(NH4)2HPO4 + 4KCI, on ignition = K,P207 + 4NH4CI + H2O ■the residual phosphates of the alkali metals, when moistened with hydrochloric acid, Slye the flame-tests, but the residue of phosphoric acid obstructs the analysis. The phos- phoric acid may be removed by acetate of lead, and the excess of lead by hydrosulphuric acid. 150. A more convenient method of removing magnesium is to precipitate it with solution of barium hydrate, and filter, and remove the excess of barium hydrate from le filtrate by addition of sulphuric acid, filtering again. 38 The Metals of the Third Group. GROUP 111. 151. The Metals of the Earths, and the more Electro-Positive of the Heavy Metals. Aluminium A1 = 27.009 Didymium D = 142.121 Chromium Cr = 52.009 Titanium Ti - 47.980 Iron Pe = 55.913 Tantalum Ta = 182.144 Manganese Mn 54.855 Niobium Nb — Cobalt Co = 58.887 Yttrium Y = 88.900 Nickel Ni 57.928 Erbium E . 165.871 Zinc Zn = 64.9045 Vanadium V = 51.256 Uranium Ur = 238.482 Decipium = 171. ? Indium ... In = 113.398 Gallium .Ga - 68.584 Beryllium = 9.085 Samarium ....... Sm - 149.801 Thorium. Th 232.020 Scandium Sc - 43.980 Zirconium Zr = 89.367 Terbium Tr = 148.5 Cerium Ce = 140.424 Thallium Tl = 203.715 Lanthanum La = 138.019 Ytterbium Yb = 172.761 -152. The metals above named gradually oxidize at their surfaces in the air, and their oxides are not decomposed by heat alone. Zinc, iron, cobalt, nickel, and, with more difficulty, manganese, chromium, and most of the other metals of the group, are reduced from their oxides by ignition at white heat with charcoal. They are all reduced from oxides by the metals of the alkalies. Iron is gradually changed from ferrous to ferric combina- tions by contact with the air. Chromium and manganese are oxidized from bases to acid radicals by ignition with an active supply of oxygen in presence of alsalies; these acid radicals acting as pretty strong oxidizing agents. 153. The oxides and liydroxides of third-group metals are insoluble in water, hence they are precipitated from all their salts by alkalies. In the case of zinc, the precipitate redissolves in all the alkalies ; the aluminium hydroxide redissolves in the fixed alkalies, but very slightly in ammonium hydroxide ; the precipitate of chromium redissolves in cold solution of fixed alkalies, precipitating again on boiling ; the hydroxides of cobalt and nickel dissolve in ammonium hydroxide. The oxides of Al, Cr, and Pe, after ig- nition, are difficultly soluble by acids. The presence of tartaric acid, citric acid, sugar, and some other organic substances, prevents the precipitation of base* of this group by alkalies. 154. Salts of ammonium (as NH4CI) dissolve moderate quantities of the hydroxides of manganese, zinc, cobalt, nickel, and ferrous hydroxide; but, so far from dissolving the hydroxide of aluminium, they lessen its slight solubility in ammonium hydroxide. 155. It thus appears that ammonium hydroxide, with ammonium chloride, the latter necessary on account of magnesium, manganese, and alu- The Metals of the Third Group. 39 minium, Avill fully precipitate only aluminium, chromium, and ferricum °f the important metals named in third group. In many plans of separa- tion these three metals constitute a separate group, and we shall refer to them as Division First of the group. 156. Ammonium sulphide precipitates all the metals of the third group from neutral or ammoniacal solutions, as follows : The sidphides of the group—those of Fe, Mn, Co, Hi, and Zn—are soluble in dilute acids, which acids keep them in solution during the second group precipitation ; but are insoluble in icater, Avhich enables thorn to be precipitated by alkaline sul- phides, and separated from the fourth and fifth groups. The other two metals, A 1 and Cr, do not form sulphides, in the Avet Avay, but are precipi- tated as hydroxides by alkaline sulphides. 157. Hydrosulphuric acid scarcely precipitates the metals of this group, unless it be from some of their acetates—owing to the solubility of the sulphides in the acids, Avhich Avould be set free in their formation, bhus, this change cannot occur— because the two products would decompose each other. Therefore, neu- tralized hydrosulphuric acid—a soluble sulphide—is employed for this group, and in a neutral or amraoniacal solution. As most of the chemically normal salts of heavy metals have an acid reaction to test-paper, we can ollly assure ourselves of the requisite neutrality by adding sufficient ammo- -11 mm hydroxide, which itself precipitates the larger number of the bases, as we have just seen (153). But the resulting precipitate of hydroxide, as is immediately changed to sulphide, FeS, by subsequent addition °f ammonium sulphide ; as the student may observe, by the alteration in color of the precipitate. Ferric and manganic salts are reduced to ferrous and manganous salts, % hydrosulphuric acid, in solution, with a precipitation of sulphur, and tbe corresponding reaction occurs Avith chromates. 158. Soluble carbonates precipitate all the metals of this group, in ac- c°rdance with the general statement for bases not alkali. With aluminium and chromium, the precipitates dissolve sparingly in excess of potassium or s°dium carbonate; with zinc, the precipitate dissohres in excess of ammo- lHum carbonate. In the case of ferrous and manganous salts, the precipi- tates are normal carbonates ; Avith zinc, cobalt, and nickel salts, they are basic carbonates ; while with ferric, aluminic, and chromic salts, the pre- Clpitates are almost or quite Avholly hydroxides. Barium carbonate precipi- tates the pseudo-triads, which, in the cold and from salts not sulphates, is a SBparation from the other bases of this group. 159. Soluble phosphates precipitate these as they do other non-alkali 'ases. The acid solutions of phosphates of the metals of the third group are precipitated by neutralization. The recently-precipitated phosphates. FeCl, + H2S = FeS + 2HCI 40 A L UMINIUM. of all the metals of this group which form sulphides, are transformed to sul- phides by ammonium sulphide : FeHP04 + (NH4)2S = FeS + (NH4)2HPO4 Hence, the only phosphates which may occur in a sulphide precipitate are those of Al, Or, Ba, Sr, Ca, and Mg. 160. The metals of the third group are not easily reduced from their compounds to the metallic state by ignition before the blow-pipe, even on charcoal, except zinc, which then vaporizes. Three of them, however— iron, cobalt, and nickel—are reducible to magnetic oxides. The larger num- ber of them give characteristic colors to beads of borax and of microcosmic salt, fused on a loop of platinum wire before the blow-pipe. None of them color the flame or give spectra, unless vaporized by a higher temperature than that of Bunsen’s burner. ALUMINIUM. Al = 27.009. 161. Specific gravity, 2.583 (Mallet). Melting point about 700° 0. (1292° F.) (Hebrew). 162. Occurrence.—Is not found in nature. Is found in corundum, ruby, and sapphire, as nearly pure A1203 ;in diaspore (A1202(0H)2); in bauxite (A120(0H)J ; in felspar (K2AI2SiOe); in cryolite (Na6AI2FI2). As a silicate in all clays and in very many minerals. It is widely distributed, constituting about one-twelfth of the earth’s crust. 163. Preparation.—(l) By electrolysis of the fused Wa2Al2Cl6. (2) By fusion of cryolite or the chloride with FT aor K. (3) By heating Ha2Al2Cl6 with zinc, with which it forms an alloy from which the zinc is driven off by a white heat. (4) By fusion of the chloride with potassium cyanide. (5) By fusing A12S3 witli iron. A great many new methods have been patented. Its alloys, especially that with copper, promise extensive use- fulness. 164. Properties.—A tin-white metal; after fusion about as hard as silver, after hammering about- as hard as soft iron ; very malleable, very ductile, and very sonorous. Its tenacity is nearly equal to that of copper. It conducts electricity eight times better than iron (Deyille). The pure metal remains untarnished in the air. Impure specimens become coated With a film of oxide. 165. Oxide and Hydroxide.—A1203 is formed by heating the hydrox- ide, nitrate, acetate, or other organic salt. It is insoluble in acids after ignition, but may be dissolved after fusion with KHSO, or Na„C03. Al2(OH)6 is formed when aluminium salts are precipitated with cold am- monium hydroxide. A120(0H)4 is formed if the precipitation is made at 100° C. 166. Solubilities.—The chloride and bromide are deliquescent and in- Aluminium. 41 stable. The iodide is known only in solution, the cyanide is not known, the acetate is delicpiescent. Aluminium is the most representative con- stituent of that large class of isomorphous double salts, called alums, per- nianent or slightly efflorescent, as KA1(504)2.12 aq., or K2A12(S04)4.34 aq. The oxide, hydroxide, and phosphate are the principal insoluble com- binations. Existence of the carbonate is doubtful. The sulphide, made only by heat in the dry way, is decomposed by water. Most insoluble salts of aluminium are changed to soluble compounds by action of fixed alkali hydroxides. In analysis, aluminium is obtained in the First Division of O-roup 111., by precipitation, by excess of ammonium hydroxide, with am- monium chloride ; then separated from the other members of the First Division by solution with excess of potassium or sodium hydrate (167). Excess of fixed alkali hydrate in boiling solution leaves only aluminium and zinc, of the third-group metals, dissolved, and it is separated from zinc, by non-precipitation with sulphides, and by precipitation with excess of am- monium hydrate, 167. Reactions of Aluminium Salts.—The alkali hydroxides precipi- tate aluminium hydroxide, grayish-white, gelatinous, A12(OH)0, soluble in fixed alkali hydroxides, slightly soluble in ammonium hydroxide, though not so if ammonium chloride be present : A12CIo + GKOH - A12(0H)6 + 6KCI A12(OH)o + 2KOH = K2A1204* + 4H20 This alkaline solution of aluminium differs from that of zinc, both in not being at all precipitated by boiliyg, and in being precipitated by excess °f ammonium chloride, morn readily when heated : K2A1204 + 2NH4CI + 4H20 '= Al2(OH)6 + 2KCI + 2NH,OH Sufficient ammonium chloride must be added, first to salify the free po- tassium or sodium hydroxide. Hydrosulphuric acid does not precipitate aluminium from any com- bination ; but ammonium sulphide precipitates the aluminium hydroxide, with evolution of hydrosulphuric acid (15G): Al2(50.,)3 + B(NH4)2S + OHnO = A12(0H)6 + 3(NH4)2504 + 3H2S 168. Alkali carbonates also precipitate the hydroxide with evolution of carbonic anhydride—the precipitate being sparingly soluble in excess of so- dium or potassium carbonate, scarcely at all soluble in excess of ammonium carbonate : AMS04)3 + 3K2CO;, + 3HoO = Alo(OH)o + , 3K2504 + 3C02 Barium carbonate, on digestion in the cold, precipitates the whole °f aluminium from its chloride, as hydroxide mixed with a little basic salt. * Or A1202(OK'2. A series of volumetric determinations, made by Mr, J. N. Ayres and the author ■Tovr. Am. Chem. Soc., Feb., 1880), give results according closely with this formula for potassium alu- Inina(e, and Na2Al204 for sodium aluminate—as fixed by the constituents of the solutions when the preci- Pdates are held dissolved by least excess of alkali. A L UMINIUM. Basic acetate of aluminium is precipitated ah follows : To the solution of aluminium salt add a little sodium or ammonium carbonate, as much as can be added without leaving a precipitate on stirring, then add excels of sodium or ammonium acetate, and boil for some time, when the precipi- tation at, length becomes very nearly complete. 169. Alkali phosphates precipitate aluminium phosphate, white, Al2- (P04)2, soluble in the fixed alkali hydroxides, not in acetic acid. To separate Al from PO4, fuse the precipitate or powdered substance with parts finely divided silica and 6 parts dried sodium carbonate in a platinum crucible, for half an hour. Digest the mass for some time in water; add ammonium carbonate in excess, filter and wash. The residue consists of aluminium sodium silicate; the solution con- tains the PO4, as sodium phosphate. The Al can be obtained from the residue by dis- solving it in hydrochloric acid, evaporating to dryness to render the silica insoluble. Treat with hydrochloric acid, and filter; the filtrate containing aluminium chloride. Also, Al (and ferricum) may be separated from PO4 by dissolving in hydrochloric acid, adding tartaric acid and then ammonia, and digesting some time with the mixture of magnesium sulphate, ammonium chloride, and ammonium hydroxide. The fil- trate contains most of the aluminium. 170. Sodium thiosulphate precipitates, from aluminium salts, in neutral solutions, aluminium hydroxide, with free sulphur, and liberation of sulphurous anhydride (a) The liquid should be dilute, and boiled till it no longer gives the odor of sulphur dioxide. This precipitation (Chancel’s) is a separation from iron. If phosphates are present, and sodium acetate with acetic acid to acidify slightly, the aluminium is precipitated as phosphate. Potassium ferrocyanide very slowly precipitates a white mixture of aluminium hy- droxide and ferrous cyanide with formation of hydrocyanic acid. Ferricyanides do not precipitate aluminium; neither do oxalates. Solution of borax precipitates an acid alu- minium borate, quickly changed to aluminium hydroxide. In very concentrated solu- tions, addition of potassium sulphate causes the crystallization of alum, potassium alu- minium sulphate, in regular octahedrons or cubes. a. 2AI2(S04)3 + 6Na.S.03 + (iH.O = 2Al2(OH)8 + BS2 + OSO2 + 6Na2SO4 171. Compounds of aluminium are not reduced to the metal, but most of them are reduced to the oxide, by ignition on charcoal. If now this residue is moistened with solution of cobaltous nitrate, and again strongly ignited, it assumes a blue color. This test is conclusive only with infusible compounds, and applies only in absence of colored oxides. 172. Estimation.—Aluminium is invariably weighed as the oxide, after ignition. It is separated from zinc as a basic acetate ; from chromium by oxidizing the latter to chromic acid, by boiling with potassium chlorate and nitric acid, or by fusing with KNOs and Na2C03, or by action of Cl or Br in presence of KOH, and after acidulating with HCI precipitating the alu- minium with ammonium hydroxide. It may be separated from iron by the thiosulphate process ; also by precipitation with ammonium sulphide after adding tartrate of potassium and excess of ammonium hydroxide (Carnot). Chromium. 43 173. Oxidation.—Pure aluminium does not dissolve in nitric acid. It dissolves very slowly in cold sulphuric acid, evolving hydrogen, and if hot evolving sulphur dioxide. It dissolves in HCI and in solution of KOH, but molten KOH has no action on it. It reduces solutions of Pb, Ag, Hg, Sn (Bi incompletely), Cu, Cd, Fe, Co, Hi (Zn in alkaline solution only), Te, Se, Au, and Pt lo the metallic state. As becomes AsH3 with alkalies and acids, Sb becomes SbH3 with acids, but with alkalies becomes metallic antimony. Compounds of Mn having more than two bonds are reduced to the dyad. Chromic acid is changed to a chromic compound. It also re- duces very many acids. CHROMIUM. Cr = 52.009. 174. Specific gravity varies, according to method of preparation, from 5.9 (Richter) to 7.3 (Bunsen). Melting point above that of platinum (Deyille). 175. Occurrence.—Not found native. It is found in several minerals. Ch rome-ironstone or chromite (PeOor2O3) is the chief ore of chromium, and is usually employed in the manufacture of chromium compounds. 176. Preparation.—(l) By electrolysis of the chloride. (2) By fusing the chloride with potassium or sodium. (3) By ignition of the oxide with carbon. (4) By fusing Cr2Cl6 with zinc, using KCI and T4aCl as a flux, and Removing the excess of zinc by dissolving it in nitric acid, which does not dissolve metallic chromium. 177. Properties.—A light-green, crystalline, and almost infusible pow- der, .5 to .75 per cent, renders steel harder and improves its quality. 178. Oxide and Hydroxide.—Chromous oxide (CrO) has not been delated. The corresponding hydroxide Cr(OH)2 is made by treating CrCl2 with KOH. Chromic oxide (Cr„o 3) is made by a great variety of methods, among which are fusing the nitrate, or higher or lower oxides and hydrox- ]des in the air; heating mercurous chromate, or the dichromates of the alkalies. 4Hg„Cr04 = 2Cr„03 + BHg + 502 . (HH4)2Cr2O7 = Cr203 + + 4H20 4K2Cr207 = 2Cr203 + 4K20rO4 + 302 the last the K2Cr04 may be separated bv water. After heating to red- "ess 0r2O3 is insoluble in acids. Chromic hydroxide is precipitated by add- lng NH4OH to chromic solutions. That formed by precipitating with ?"OH or NaOH retains traces of the alkali, not easily removed by washing, " addition to chromic acid (H,,0r04) a per-chroroic acid, supposed to be Cr04, is formed when H Cr04 is treated with peroxide of hydrogen (aA). 179. Reactions of Chromous Salts.—Chromons chloride (CrCl.,) is 44 Chromium. formed by dissolving the metal in HCI, or by gently heating Cr2Cl6 in hy- drogen gas. The blue solution formed by treating a solution of Cr2Cl6 with zinc contains CrCl2 (J. Pr. Chem , 90, 12), It is a strong reducing agent, oxidizing rapidly when exposed to the air. From its solution KOH and HaOH precipitate Cr(OH)2. Precipitates are formed by NH(OH, by solu- ble carbonates, sulphides, sulphites, etc. 180. Solubilities of Chromic Salts.—There are two modifications of chromic salts, one having a green color and the other violet to red. There are many double salts. The chloride is deliquescent. Chromic oxide, hydroxide, and 'phosphate are insoluble in water. The carbonate and sulphide are not formed in the wet way. There are modifi- cations of the chloride and sulphate insoluble in water. In analysis, chro- mium is precipitated in the third group as a hydroxide, and identified by the oxidation of this hydroxide to a salt of chromic acid, known by its col- ored precipitates with lead and barium salts. It is separated in the First Division of the group. 181. Reactions of Chromic Salts.—The fixed alkali hydroxides, as, KOH, precipitate tiie bluish-green chromic hydroxide, Cr2(OH)6. Other hydroxides are formed in certain conditions. The precipitate redissolves readily in excess of the alkalies while cold, the green solution being K2Cr204. Long boiling reprecipitates the whole of the chromium, as hydroxide ; the same result is effected on heating by addition of ammonium chloride. (Compare 167, last equation.) Ammonium hydroxide precipitates chromic hydroxide, which but slightly redissolves with excess of the alkali in the cold and all reprecipi- tates readily on heating. The precipitate from solutions of green chromic salts is grayish green, dissolving with acids to form a green solution again ; from solutions of violet chromic salts the precipitate is grayish-blue, dis- solving with acids to reproduce the violet solution. The tints are, how- ever, modified by the degree of concentration of solution, and by other con- ditions. 182. Hydrosulphuric acid does not affect solutions of chromic salts, whether acid, neutral, or alkaline, and ammonium sulphide precipitates the hydroxide with evolution of hydrosulphuric acid. The equation corre- sponds to that for aluminium (167). Both hydrosulphuric acid and ammonium sulphide, acting on Chromic Acid or chromates, abstract oxygen and form the chromic base. In the neutral solution for the third-group precipitation this deoxidation leaves the chromium in the precipitate as a hydrate ; whence it is that the occur- rence of chromium in the third group of bases, as frequently as otherwise, must be referred to the existence of combinations of chromic acid in the material examined. (See 194 a.) 183. Alkali carbonates precipitate chromium hydroxide, nearly free Chromic Acid. 45 from carbonate, somewhat soluble in excess of potassium or sodium car- bonate : Barium carbonate precipitates chromium from its solutions (better from the chloride) as a hydrate with some basic salt, the precipitate being complete after long digestion in the cold. For removal of excess of reagent, add H2SOi and the filtrate will contain the chromium as a sulphate. 184. Soluble phosphates—as NaoHPO*—precipitate chromic phosphate, Cr2(POj)2, insoluble in acetic acid. Cyanide of Potassium precipitates the hydroxide. Ferrocy- anides, and oxalates, cause no precipitates. Potassium chromate colors an acid solu- tion of chromic salt brown-yellow ; on addition of ammonium hydroxide, a precipitate of the same color is obtained, chromic chromate. (Maus) Pogg. Ann. 9, 127. Cr2(S04)3 + SK2C03 + SH20 = Cr2(OH)6 + SK2S04 + SCO* Cr2(SO4)3 + SK2CrOi = Cr202CrO4 + 3K2Cr.20, + 8K2S04 185. Chromic oxide and chromic salts dissolve in beads of mieroeos- ttdc salt, and of borax, before the blow-pipe, in both reducing and oxidiz- hig flames, with a yellowish-green tint while hot, becoming emerald-green when cold. 186. Estimation.—Chromium is weighed as an oxide. It is brought hito this form either by precipitation as a hydroxide and ignition, or by simple ignition. It may, however, be changed to a chromate, and .esti- mated as such. 187. Oxidation.—Metallic chromium does not dissolve in HN03, but quickly dissolves in HCI, forming CrCl2, which is a strong reducing agent, changing HgCl2 to Hg2Cl2, CuS04 to Cu, SnCl2 to Sn, etc. Chromic compounds in alkaline mixture are oxidized to chromates by reducing PbO„ to PbO, Ag20 to Ag, Hg„o and HgO to Hg, CuO to CuzO, KMn,OB to Mn02; also by Cl, Br, and I forming a chloride, bromide, or lQdide. A chromate is formed when, in the presence of Na2C03, any chro- mic compound is fused with KN03, KCI03, KBr03, or KIOs; NO and a chloride, bromide, or iodide being produced. CHROMIC ACID. 188. Chromic anhydride, Cr03, commonly called “chromic acid,” is a Bcarlet-red solid, usually in acicular crystals, very deliquescent in the air, aild soluble in a small proportion of water. It is a very powerful oxidizing agent, acting explosively with combustible substances, and as a caustic to living tissues. Its soluble salts are poisonous, and have a bitter metallic taste. 189. The alkali metals form yellow normal chromates and reddish di- chromates ; most other metals form normal chromates, yellow or red ; a, e'v form only basic or instable chromates. Most soluble salts of chromic acid crystallize in permanent forms ; sodic normal chromate is efflorescent. 190. All the chromates of the alkali metals, and those of magnesium. Iron. calcium, zinc, and copper, are soluble in water; strontium and mercuric chromates sparingly soluble ; barium, manganous, bismuth, mercurous, silver, and lead chromates insoluble in water. Nitric acid transposes chro- mates. 191. Lead salts precipitate, from normal and from superchromates, the yellow, lead chromate, PbCr04, slowly soluble in nitric acid, not soluble in acetic acid, difficultly soluble in potassium hydrate. 192. Barium salts precipitate from solutions of normal chromates, also from concentrated solutions of superchromates, the normal barium chro- mate, yellow, soluble in hydrochloric and nitric acids, slightly soluble in chromic acid (105). 193. Silver salts precipitate silver chromate, Ag.CrCq, dark red, soluble in nitric acid and in ammonia. Mercurous nitrate precipitates mercurous chromate, Hg2Cr04, dark red, decomposed by ignition into chromic oxide, oxygen, and vapor of mercury (178). 194. Chromic anhydride and chromates are DEOXIDIZED TO CHROMIC com- pounds by various reducing agents The following instances occur frequently in quali- tative analysis : other examples are given, in the study of Chromate reductions. Hydrosulphuric acid, in acid solutions, quickly causes reduction to a green chromic salt solution (a). At first the sulphur is all precipitated, white in the green liquid ; but on warming, it slowly dissolves by oxidation to sulphurous acid, with precipitation of brown basic chromate (b), the action being continued, with slow oxidation of the sul- phurous acid (c) [IT. B. Parsons]. Ammonium sulphide, in solutions neutral or alkaline, precipitates chromic hydrate, green, with oxidation of the sulphide. The precipitate is liable to contain sulphur. In case of yellow or supersulphide of ammonium, it is stated that thiosulphate is obtained in the solution (d). a. THoCrOi + 12HC1 + 6H,S = 2Cr2CI3 + 352 + 16H.0 I). 12H2Cr04 + 352 = 40r20,0r04 + 6H2503 + 6H20 c. 2H2CrO4 + 3H2S03 = Cr2(S04)3 + 5H20 d. K2Or20, + (NH4)2S2 + 4H20 = Or2(OHj3 + K2S203 + 2NH4OH 195. By ignition on charcoal the carbon deoxidizes chromic anhydride free or combined, and a green mas*, Cr„Oa, is left. Chromates give, in the beads, the results described for chromic base in 185. For a more complete statement of the oxidizing action of this acid see under chromic acid. Part 11. IRON. Re ■= 55.913. 196. Specific gravity variable; reduced by electricity 8.1393 (Smith) ; by hydrogen 8.007 (Schiff); by carbon 7.130 (Playfair). Melting-point, gray cast, 1275° 0. (2327° F.) (Ledebur); cast steel, 1375° C. (2507° F.) (Bloxam). 197. Occurrence.—Native iron is rarely found except in meteorites. The chief ores of iron are red hematite or specular iron ore (Fe„o3), brown hematite (2Fe203,3H„0), magnetic iron ore (Fe3o4), iron pyrites (FeSj, Iron. 47 spathic iron ore (FeCOj, clay iron-stone (FeCO3 with clay), black band (FeoO3 mixed with bituminous matter). 198. Preparation.—Pure iron is not usually found in the market. It is made : (1) by electrolysis ; (2) by heating its purified salts with hydrogen; (3) by heating the purified salts with some form of carbon ; (4) in metal- lurgy it is made from the ores, and the reducing agent is coal, charcoal, and, more recently, natural gas. 199. Properties.— Pure iron is the most tenacious of all the metals ex- cept cobalt and nickel. It softens at a red heat, may be welded at a white heat, but above the melting point is brittle under the hammer. Finely di- vided iron burns in the air when ignited. When made by (2) it is a very fine powder, and if made at as low a temperature as practicable, it takes fire spontaneously when exposed to the air. Steel contains from 0.2 to 1,5 per cent, of carbon, while cast iron contains from 2to 5 per cent. Pure iron is attracted by the magnet, but does not retain its magnetism. Permanent magnets are made of steel. All the ordinary properties of iron are too well known to need any description. 200. Oxides and Hydroxides.—Ferrous oxide (PeO) is made from Fe203 by heating it to 300° C. in an atmosphere of hydrogen ; also by heat- ing Fe2C204 to IGO° C., air being excluded. It takes fire spontaneously in the air, oxidizing to Fe203. Ferrous hydroxide, Fe(OH)2 is formed by pre- cipitating ferrous salts with KOH or HaOH, perfectly white when pure, but frequently green from partial oxidation. Ferric oxide, Fe„03, is formed ky heating FeO, Fe(OH)2, or any ferrous salt consisting of a volatile or organic acid in the air; more rapidly by heating Fe2(OH)6, Fe;(N03)6, or I'e!(SO<).. Ferric hydroxide is formed by precipitating cold dilute ferric salts with alkalies or alkali carbonates, and drying at 100° C. If KOH or HaOH is used, the precipitate requires longer washing than when NH(OH is employed. By increasing the temperature and concentration of the solu- tions the following definite compounds may be formed : Fe202(OH)2, ;0(OH)4, Fe4Or(OH)2, Fe403(OH)6, Fe0O4(OH)10 (hexaferric-tetroxide- dekahydroxide). Fe;i04 is slowly formed by heating FeO or Fe203 to a "’kite heat. Its corresponding hydroxide may be made by precipitation, Feci2 + Fe.Cl, + BHH.OH = Fe,(QH)B + 8HH4CI Fe3(OH)8 when keated to 90° C. forms Fe304. The black color and magnetic properties skow that it is a chemical salt and not a mechanical mixture of FeO and Fe203. (Fe2)vi acts as an acid towards the Fe" ; this oxide, Fe304, or FeFe204, '"ay be called ferrous ferrite. Other ferrites have been formed—e.g., cal- c'"m ferrite, CaFe.O, (Percy) ; MgFe204 and BaFe204 (List, Ber. 1878, 1;>12); zinc ferrite, ZnFe204 (Ebelmen). Compare potassium aluminate (K2A1204), and potassium chromite (K„Cr„O4) (167 and 181). 201. Ferric Acid, H„FeOt, and its anhydride (Feo3) have not been iso- kited. Potassium ferrate (K2FeO4) is made (1) by electrolysis ; (2) by heating 48 Iron. iron filings, FeO or Fe203, to a red heat with 003 ; (3) by heating Fe2(OH)6 with potassium peroxide K204 ; (4) by passing Cl or Br into a so- lution of 5 parts of KOH in 8 parts of water in which Fe2(OH)c is suspend- ed ; the temperature should be not above 50° C. (122 F.) It has a purple color; is a strong oxidizing agent. It slowly decomposes on standing (4K2FeQ4 + 10H2O = BKOH + 2Fe2(OH)6 + 30J. With barium salts ic precipitates a stable barium ferrate (BaFeOj. 202. Solubility of Iron.—Iron dissolves, in hydrochloric acid, and in dilute sulphuric acid, to ferrous salts, with liberation of hydrogen (a); con- centrated cold H.,SOt has no action, but if hot SO„ is evolved and a ferric salt formed (b) ;in moderately dilute nitric acid, with heat, to ferric ni- trate, liberating chiefly nitric oxide (c); in cold dilute nitric acid, forming ferrous nitrate with production of ammonium nitrate (d), of nitrous oxide (e), or of hydrogen (/): a. Fe + HjSOi = FeSO* + H2 b. 2Fe + 6H2504 = Fe.hSO.O3 + BSO2 + 6H20 c. 3Fe + BHNO3 = Fe2(N03)6 + 3NO + 4H,0 d. 4Fe + 10HNO3 = 4Fe(NO3)2 + NH4NO3 + 3H.0 e. 4Fe + 10HNO3 = 4Fe(NO3)2 + N2O + 5H..0 /. Fe + 2HNO.J = Fe(NO3)2 + H2 In dissolving the iron of commerce in hydrochloric acid the carbon which it always contains, so far as combined in the carbide of iron, will pass off in gaseous hydrocarbons, and so far as uncombined will remain un- dissolved, as graphitoid carbon. Iron acts as a base in two kinds of salts: the ferrous and the ferric ; botli are stable, in considerable variations of temperature, when undis- turbed by other substances ; but the ferrous compounds are changed to fer- ric by contact with the air, and by oxidizing agents generally ; while the ferric compounds are permanent in the air, but are changed to ferrous com- binations by reducing agents. In the systematic course of analysis, by the treatment necessary in separation from other metals, the ferric compounds are reduced to ferrous compounds, and then, by air and by reagents, par- tially or wholly changed to ferric compounds again, and the original sub- stance must always be tested for determination, whether ferrous.or fer- ric. The metal oxidizes in moist air to tetraferric-trioxide-hexahydrox- ide, Fe403(OH)6. T3y ignition in the air, chiefly ferrous oxide is formed, but by a long-continued white heat Fe3Q4 is formed, see (200). Scale ox- ide is (Feo)6Fe„o3. 203. Solubilities of Ferrous Salts.—FERROUS salts, in crystals and in solution, have a light green color. The oxide is black ; the salts slightly redden litmus. The sulphate (7 aq.) is efflorescent; the chloride, bromide, iodide, and citrate are deliquescent; the hydrate, chlorate, and sulphite are especially instable. The hydroxide, oxide, carbonate, sulphite, phosphate, borate, oxalate, Iron, 49 cyanide, ferrocyanide, ferricyanide, tartrate, and tannate are insoluble in Pater. In analysis, ferrous compounds are identified as ferrous by their blue precipitate with ferricyanide (208); and, as iron, by the red solution which, after oxidation, they form with sulphocyanate. 204. Reactions of Ferrous Salts.—-The alkali hydrates precipitate ferrous hydroxide, Fe(OH)2, white if pure, but seldom obtained sufficiently ftee from ferric hydrate to be clear white, and quickly changing, in the air, to ferroso-ferric hydroxide, of a dirty-green to black color, then to fer- ric hydroxide (212), of a reddish-brown color. The fixed alkalies adhere to this precipitate. Ammonium chloride or sulphate, sugar, and many or- ganic acids, to a slight extent, dissolve the ferrous hydrate or prevent its formation (compare 154). 205. The soluble carbonates precipitate, from purely ferrous solutions, ferrous carbonate, PeC03, white if pure, but soon changing, in the air, to the reddish-brown ferric-hydrate (212). 206. Hydrosulphuric acid does not disturb ferrous salts—the acetate being only slightly precipitated, as explained in 157. Ammonium sul- phide precipitates ferrous sulphide, FeS, black. The moist precipitate is slowly converted, in the air, to ferrous sulphate ; and afterward to basic ferric sulphate, Fe„O(SO4)2. 207. Alkali phosphates—as Na„HPO-precipitate two-thirds metallic ferrous phosphate, FeHP04, mixed with the full-metallic salt, Fe3(POj2, 'vhite to bluish -white. By the addition of an alkali acetate, the precipitate 18 obtained of full-metallic phosphate exclusively : 3PeSO, + 2Na,HPO, + 2Na0.H302 = Fe3(P04)2 + 8Na2S04 + 2HC2H302 208. Cyanides—as KCN—give a yellowish-red precipitate, chiefly fer- r°ns cyanide, soluble in excess of the reagent; the solution constituting Potassium ferrocyanide. iPerrocyanides—as K,Fe(CN)r—precipitate potassium-ferrous ferrocy- yride, K2PeFe(CN)6 (Everitt’s salt), bluish-white, insoluble in acids. This 18 converted to Prussian blue (217), gradually by exposure to the air, imme- diately by oxidizing agents : 4K„FeFe(CN)6 +O2 + 4HCI = Fe4(Fe(ON)6)3 + K.,Fe(CN)6 + 4KCI + 2H20 Perricyanides—as K6Fe2(CN)12—precipitate (even from dilute solutions) ferrous ferricyanide, Fe3Fe2(CN)12, dark blue, insoluble in acids. This ln\portant test reveals the presence of traces of ferrous salt, in ferric solu- tions. For this purpose the solution must be dilute, as stated at 217, and original solution always employed, because the oxidation of iron is al- ei’ed by chemical operations. ■Alkali hydroxides decompose the precipitates above named : with po- tassium ferrous ferrocyanide, forming alkali ferrocyanide and ferrous hy- droxide ; with ferrous ferricyanide, forming alkali ferricyanide and ferrous hydroxide : 50 Iron. K2FeFe(CN)6 + 2KOH = K4Fe(CN)6 + Fe(OH)2 Fe3Fe2(CN)i2 + 6KOH = KOFe2(ON)12 + 3Fe(OH)2 Sulphocyanates give no reaction with ferrous salts. 209. Oxalic acid and oxalates precipitate ferrous oxalate, PeC204, yellowish-white, crystalline, sparingly soluble in boiling water, decomposed by mineral acids not too dilute. 210. Tannic acid, and tincture of galls, with concentrated solutions of purely ferrous salts, give a white gelatinous precipitate of ferrous tannate, which is quickly oxidized by exposure to the air to blue-black ferric tan- nate—long used for writing ink. 211. By ignition and in beads before the oxidizing flame of the blow- pipe ferrous salts give the same reactions as ferric (220). 212. Solubilities of Ferric Salts.—FERRIC salts form solutions hav- ing a brownish yellow color, and reddening litmus Most soluble ferric salts are deliquescent. Ferric oxide, in powder, is reddish-brown ;in native crystal, steel-gray. It is soluble in hydrochloric acid, not very readily, but much quicker than in other acids. The hydroxide, oxalate, phosphate, ferrocyanide, tannate, gallate, bo- rate, and sulphite are insoluble in water. The chloride is soluble in alcohol and in ether ; the sulphate is soluble in alcohol, a separation from ferrous sulphate. Ferric carbonate is not formed, and ferric sulphide is not formed in ordinary conditions of wet analysis. In analysis, ferric compounds are identified by the red solution they form with sulphocyanate, and distin- guished from ferrous forms by not causing a blue precipitate with ferricy- anide (217). Ferricum is separated in the First Division of Group 111., with the other pseudo-triads. 213. Reactions of Ferric Salts.—The alkali hydroxides precipitate ferric hydroxide, Fe2(OH)6, variable to Fe„0„(0H)2, reddish-brown, insol- uble in alkalies or ammonium salts. Salts of fixed alkalies adhere to this precipitate with great tenacity. Alkali carbonates—as K.,00^—also precipitate the hydroxide, contain- ing traces of carbonate. Regarding barium carbonate, see 219. Fe2CIO + 3K2COs + 3H20 = Fe2(OH)6 + 6KCI + 3CO* 214. Hydrosulphuric acid does not precipitate iron from ferric solu- tions ; but reduces them to the ferrous combination, with precipitation of sulphur. 215. Ammonium sulphide precipitates the ferrous sulphide with free sulphur, FeS with S, a reduction of the metal to the condition of a dyad. Hence, the ammonium sulphide precipitate contains iron in ferrous condi- tion only. 216. Phosphates—asNa2HP04—precipitate ferric phosphate,TPefPO,)2, scarcely at all soluble in acetic acid, but readily soluble in hydrochloric, Iron. 51 nitric, and sulphuric acids. Hence, ferric salts which are not acetates are precipitated by phosphoric acid with co-operation of alkali acetates (207): In this way phosphoric acid is removed from alkaline earth bases—in so- lutions of alkaline earth phosphates, in hydrochloric or nitric acid. 217. Soluble cyanides—as KCN—precipitate, from ferric salts, the hy- drate, with evolution of hydrocyanic acid (a). Ferrocyanides—as K4Fe(CH‘)6—precipitate ferric ferrocyanide, Fe4(Fe- Prussian blue, insoluble in acids, decomposed by alkalies fb). Strong acids color the reagent blue, and render the test fallacious ; acetic acid is free from this objection, and addition of potassium acetate enables the test to be made in acid solutions. By excess of the reagent the precipitate is somewhat soluble to a blue liquid. Sulphoeyanates—as KCNS—form, in solution, ferric sulphocyanate, Fe2(CNS)6, of a blood-red color so intense that this is an exceedingly deli- cate test for iron when in the ferric condition (c). Fe2Cl5 + 2H3POi + 6NaC2H302 = Fe2(P04)2 + GNaCI + CHC2H302 The red salt is freely soluble in water, alcohol, and ether, and extracted by ether from aqueous solutions; is decomposed by alkalies, but not by acids. Traces of ferric salts are revealed by adding the reagent, slightly over-saturating the mixture witli ether; the excess of which will rise to the surface, colored by any ferric sulphocyanate, concen- trated from the mixture.* The color of the liquid is destroyed by mercuric chloride (d); also by phosphates, borates, acetates, oxalates, tartrates, racemates, malates, citrates, succinates, and the acids of these salts. Molybdenum dioxide, also nitric and chloric acids, give red color with the sulphocyanate, removed by heat. To determine the condi- tion of iron the original solution only can be used (302). Ferrieyanides—as K.Fe,(CN), 2—form no precipitate in ferric solutions, but give a green, or, in some proportions, brown color to the liquid (e), which should he diluted until transparent enough to reveal minute portions of bine precipitate if ferrous salt is present (208). The addition of dannous chloride. SnCl,„ or some other strong deoxidizing agent to the mixture of ferricyanide, wherein no precipitate is found, constitutes a delicate test for ferric salts. Some of the above-named reactions of ferric salts with cyanogen com- pounds are defined in the following equations : а. Fe2C16 + 6KON + 6H.0 = Fe2(OH)G + GHCN + OKCI б. Fe4(Fe(CN)f,)3 + 13KOH = BK4Fe(CN)6 + 2Fe2(OH)6 (Decomposition of K4Fs(CN)6 by acids, see Part TI.) e. Fe.jCls + 6KCNS = Fe,(CNS)s 4- GKCI d. Fe2(CNS)6 + 3HgCI. = 3Hg(CNS)2 + Fe2Cl6 e. Fe2Cl6 -+- K6Fe2(CN)u = 6KCI + Fe2Fe2(CN)]2 218. The acetates—as NaC2H30,2—form, in the cold, a dull red liquid, ferric acetate, Fe2(CoH30„)6, not decolorized by mercuric chloride. On boil- * Natanson, Zeitsch. analyt. Chem., iii. 370. 52 Iron. ing the solution basic ferric acetate is precipitated, finally becoming hy- drate. Sulphites give, likewise, a red solution of ferric sulphite, decom- posed by boiling.* 219. Tannic acid—and tincture of galls—precipitate ferric salts blue- black, as ferric tannate, the basis of common ink. Ammonium Succinate precipitates reddish-brown ferric succinate. Carbonates of Ba, Ca, Mg, Mn, Zn, and Cu precipitate ferric hydroxide from the chloride even in the cold (leaving barium chloride in solution) : Fe2Cl6 + SBaCOo + 3H20 = Fe2(OH)6 + CEaCl2 + 3C02 The excess of the barium carbonate is filtered out with the ferric hydroxide, and may be separated by addition of sulphuric acid,which changes it to insoluble barium sulphate, and leaves ferric sulphate in solution. If ferrous chloride were in the original solution, the barium chloride formed in the reaction may be separated from it, likewise, by addi- tion of sulphuric acid. 220. The larger number of iron salts are decomposed, as solids, by heat; ferric chloride vaporizes, undecomposed, at a very little above 100° C. (212° F.) Ignition in the air changes ferrous compounds, and ignition on char- coal or by the reducing flame changes ferric compounds to the magnetic oxide, which is attracted to the magnet. In the outer flame, the borax bead, when moderately saturated with any compound of iron, acquires a reddish color while hot, fading and becoming light yellow when cold, or colorless, if feebly saturated. The same bead, held persistently in the reducing flame, becomes colorless unless strongly saturated, when it shows the pale green color of ferrous compounds. The reactions with microcosmic salt are less distinct, but similar. Cobalt, nickel, chromium, and copper conceal the reaction of iron in the bead. Ferric compounds, heated briefly in a blue borax bead holding a very little cupric oxide, leave the bead blue ; ferrous compounds so treated change the blue bead to red—the color of cuprous oxide. 221. Recapitulation of Distinctions between Ferric and Ferrous Com- pounds : Ferric Compounds. Ferrous Compounds. (1) Ferricyanides. No pre., green color, 217. Deep bine pre. Fe,Fe0(C]Sr)12. (2) Sulphocyauates. Red sol. Fe„(CNS)6. No cliange. (3) Ferrocyanides. Blue pre. Fe4(Fe(CN)6)3. Pale bine pre. K„FeFe(CN),. (4) Carbonates. Effervescence (213). No effervescence (205). (5) Cyanides. 217. 208. * Meconic acid and Formic acid form red solutions with ferric salts. Benzoic acid gives a flesh-colored precipitate; salicylic acid a deep violet color; phenol and creosote, each a blue color; saligenin a blue color; and various compounds of the “ aromatic group,” hydroxyl substitutions in benzine derivatives, give blue to violet colors. Morphine, pseudomorphine, and daphnin give the blue color. Iron. 53 (6) Reducing agents. Hydrosulph. acid. (Deodorized and S. precipitated.) Sulphurous acid. (Deodorized.) (7) Oxidizing Agents, Nitric acid, .... 202 [e) (Brown gas, by heat.) Bromine, (Deodorized.) Chlorine water, . . . (Deodorized.) 222. Estimation of Fe".—(1) By conversion into Fe„03 and weighing as such ; (2) by converting it into FeS and weighing as such ; (3) by treat- ing with AuC13 and weighing the reduced gold ; (4) volumetrically (a) by a solution of K„Mn208 of known strength, (IOFeSCU + SjMnaOs + BH2S04 = SFe2(S04)3 + K2SQ4 + 2MnSQ4 + 8HsO); (6) by a solution of K2Cr207 of known strength, (6FeSO4 + K2Cr207 + 7H2S04 = BFe2(S04)3 + K2SQ4 + Cr2(S04)3 + 7H„0), the end of the reaction is determined by KcFe2(HC)12. 223. Estimation of (Fe2)VI.—(f) It is converted into Fe203 or FeS, and weighed as such ; also (2) volumetrically (a) by a solution of sodium thiosulphate of known strength (Fe2016 + 2Na2S203 = 2FeCI2 + Na2S,Os + 2NaCI) A few drops of a solution of CuSC4 are added. The CuSC4 in some manner not well understood hastens the reaction and gives more accurate results. The end of the reaction is shown by KCHS ; (d) by a solution of SnCl2 of known strength The excess of SnCl2 used is determined by a solution of iodine in KI, and deducted from the whole amount used ; (e) by a solution of Cu2Cl2 of known strength (Fendr, + 0u2C12 = 2FeClo + 2CuCI2) A drop of KCNS is first added, and the redaction is complete when the color disappears. 224. Oxidation.—Metallic iron precipitates the free metals from solu- tions of An, Pt, Ag, Hg, Bi, and Cu. Solutions of Fe" are changed to (Fe2)VI solutions by treating with solu- tions of An'", Ag', CrVI, Mnvn, MnVI, and H202. In presence of some di- kite acid, such as H2S04 or H3P04 by Pb02, Pb304, Mn304, Mn02, Mn20„ Co203, Ni203. The following acids also oxidize Fe" to (Fe2)VI, HNO,, HN03, HCIO, HCI02, HCI03, H2S04 (if concentrated and hot), HBrO, HBrOj, HI03, Br, Cl. Br and Cl in presence of KOH changes Fe" and (Pe2)vi to K2Fe04. 225. (Fe2)VI is changed to Fe" by solutions of Sn", (Cu2)", H3P02, S3P03, H2S, H2S02, ]\fa2S203, and HI. Also by nascent hydrogen, or by (Fe„C]s + SnCl2 = 2FeCl2 + SnCl4) 54 Manganese. any of the metals which produce hydrogen when treated with acids, includ- ing Zn, Sn, Cd, Al, As, Sb, Bi, Pb, Cu, and Be. MANGANESE. Mn == 54.855. 226. Specific gravity 7.14 to 7.20 (Brunner). Melts at the highest heat of the blast furnace (Clarke). 227. Occurrence.—Not found native. Its chief ore is pyrolusite (MnOj. It is also found as braunite (Mn2o3), hausmannite (Mn3o4), manganiie (Mn2Q2(OH)2), manganese spar (MnCOa), manganese blende (MnS), and as a constituent of many other minerals. 228. Preparation.—(l) By electrolysis of the chloride. (2) By reduc- tion with metallic sodium. (3) By reduction with some form of carbon. It has not been reduced by hydrogen. 229. Properties.—A brittle metal, having the appearance of cast iron; harder than steel, and when mixed with it improves its quality. It oxidizes in the air, and must be kept under naphtha. It decomposes water at ordi- nary temperatures, rapidly when heated. 230. Oxides and Hydroxides.—Manganous oxide (MnO) represents the only base capable of forming stable manganese salts. It is formed (1) by simple ignition of Mn(OH)2, MnCOa, or Mn02O4, air being excluded, (2) By ignition of any of the higher oxides of manganese with hydrogen in a closed tube. If prepared at as low a temperature as practicable, it is a dark gray or greenish gray powder and oxidizes quickly, in the air, to Mn304. If prepared at a higher heat it is more stable. Manganous hydrox- ide (Mn(OH)s) is formed from manganous salts by precipitation with alka- lies. It quickly oxidizes in the air, forming Mn202(OH)2, thus changing from white to brown. 231. Manganic oxide (Mn2o3) is formed, (1) by heating MnO, Mn304, or MnO, to a red heat in oxygen gas (Schneider ; Pogg. Ann. 107, 605); (2) by heating Mn(OH)„ Mn(HOs)2 or any of the higher or lower oxides in the air (Fehling’s Handbuch dev Chemie, 1886, 4, 253). Mn,o,(oH)s (di- manganic dioxide dibydroxide) is formed (1) by oxidation of Mn(OH), in the air ; (2) by treating MnO, with concentrated H2S04 at a temperature of about 130° C., forming Mn,(SO4)3 and then adding water, (CARIUS, 1856, Ann. Oh. Pharrn 98, 63). Mn2(SOP3 + 4H20 = Mn202(0H)2 + 3H2504 232. Trimanganese tetroxide (Mn3o4) is formed when any of the higher or lower oxides of manganese or any manganese salts with a volatile acid are heated in the air to a white heat. The corresponding hydroxide would be Mn,(OH)8; this has not been isolated. A corresponding oxide hydrox- ide is formed by adding freshly formed and moist MnOa to an excess of MnCl2, containing NH4CI (Otto, 1855). .l/.l -A (J .1 .VAASA'. 55 233. Manganese peroxide (MnOj is formed (1) by heating Mn(NOs)2 to 200° 0. (Kuhlmanh, 1874) ; (2) by heating MnCOs with KCI03 to 300° C.; (3) by boiling any manganous salt with concentrated HN03 and KCI03. A corresponding hydroxide, Mn(OH)4, has not been isolated. Several oxide hydroxides—e.g., MnO(OH)s, Mn203(OH)2, Mn304(OH)4, etc., have been produced. 234. Manganites.—The tetrad manganese (Mnrvq sometimes plays the part of an acid toward the stronger bases. Potassium manganite, K2MnsOn, is formed by treat- ing potassium manganate, K2MnO4, with carbon dioxide (Post, 1879; Bar. 1459). Calcium manganite, (CaMasOu), lias been isolated. K2MnO4 with HNOs or H2S04, or other strong non-reducing acids, forms Mn02, or one of its hydroxides. 235. Manganese Dioxide, Mn""o2, constitutes the commercial source of manganese and an important oxidizing agent, as Pyrolusite. Salts of this type are not formed, and the action of hydrochloric acid with heat slowly dissolves (reduces) the dioxide to man- ganous chloride, with evolution of chlorine: 15K2Mn04 + 9C02 = K2Mn6Ou + SK2Mn2Os + 9K2COs Free chlorine, bromine, and iodine are obtained in accordance with this reaction (or some modification of it), and immense quantities of native manganese dioxide are re- quired for the liberation of these elements in manufacturing operations. The production of chlorine is frequently effected by using sulphuric acid and common salt, instead of hydrochloric acid: MnO, + 4HCI = MnCl2 + Cl2 + 2H20 MnO, + 2NaCI + 2H.SO = MnSQ4 + Na2S04 + Cl2 + 2HaO Oxygen also can be obtained from binoxide of manganese, by action of sulphuric acid: Further, regarding Mn"" as an oxidizing agent, see Part 11. 3MnO, + SHoSO., = 2MnSO4 + 3H.0 + 02 236. Manganates.—Manganic acid, H2Mnvi04, is not known in a free state. The corresponding potassium salt, K2MnO4, is formed when any form of manganese is fused ■with KOH or K2003 (1) in the air, see (6) and (c) (249), oxygen being absorbed; or (2) with KNO3 or KClOs, NO or KCI being formed. A manganate of the alkali metals is solu- ble in water, with gradual decomposition into manganese dioxide and permanganates : 3K2Mn04 + 3H.0 = K2Mn,OB + MnQ2 + 4KOH Free alkali retards, and free acids and boiling promote, this change. Manganates have a green color, which turns to the red of permanganates during the decomposition inevita- ble ir solution. This is the usual method of manufacturing K2Mn208. 237. Permanganic Acid is not in use as an acid, but is represented by the perman- ganates, as K..(Mno)xivO«" The permanganic acid radical is at once decomposed by addition of strong acids to a solid permanganate, but in water solution this decomposi- tion does not at once take place, except by contact with oxidizable substances. The ox-1 idizing power of permanganates extends to a great number of substances, possesses dif- ferent characteristics in acid and in alkaline solutions, acts in many cases so rapidly as to be violently explosive, and is of such quantity that four parts of the absolute potassium salt furnish over one part of oxygen (equation a). For the study of Permanganates in oxidation, see Part 11. The reactions with ferrous salts (6) and with oxalic acid (c) are much used in volumetric analysis : * In permanganates, manganese may he considered as an octad, in the compounds of which two of its atoms are held to each other by one bond of each ; the pair having twice seven bonds for other elements, »nd having always an even number of atoms in correctly written formulae. 56 Manganesk. a. 3X„Mn208 + 6H2SQ4 = 4MnSO, + 2X,S04 + 50* + 6H.0 b. X2Mn208 + IOFeCI. + 16HC1 = 3MnCl2 + 3XOI + 5Fe2O16 + 8H2Q c. K.Mn.Og + 5H.0204 + 6HCI = 3MnCI2 + BH2Q + 10CO2 + 2KCI 238. Permanganates are all soluble in water, silver permanganate being only sparingly soluble. The most of them are deliquescent. Their solutions have a deep red color. Slight deoxidation may give the green color of manganate. 239. Manganese is reduced to the manganous condition, from all its other degrees of combination, by boiling with hydrochloric acid. In this, its only stable form, it is most perfectly identified as manganese, and the various reactions of (a) the manganous base in the wet way obtained—34o and after. b. For reactions characteristic of the manganic base, see 251 and after. c. If the substance be a black powder, insoluble in water, but dissolving to manganous chloride in hydrochloric acid, with evolution of chlorine even in the cold (335), it is MnOj, Mn203, Mn304, or a mixture of these or their hydroxides. d. IE having a green color (336). and being soluble in water with decomposition, etc., lea ving MnO.., it is a manganate, representing manganic acid. e. If soluble in water to a red color, and, by deoxidation, losing color (and leaving manganous base ov Mno2), it is indicated as a permanganate (237), representing perman- ganic acid. 240. Solubilities of Manganous Salts,—The chloride (4 aq.), bro- mide, iodide, and nitrate, are deliquescent; the sulphate (7 aq.) is efflor- escent. Manganous oxide, hydroxide, sulphide, carhoyiate, phosphate, oxalate, borate, and sulphite, are insoluble in water. The hydroxide is insoluble in alkalies, but soluble in solution of ammonium salts. 241. Reactions of Manganous Salts.—ln analysis, manganese is iden- tified by the oxidation of manganous hydroxide or oxide, to manganate (248) or permanganate (248), each recognized by its bright color. As to determination of the oxidation of manganese, see 239 242. The alkali hydrates precipitate, from soluble manganous salts, manganous hydroxide, Mn(OH)„, white, soon turning brown in the air by oxidation to dimanganic dioxide dihydroxide, Mn20„(GH)2. The precipitate is insoluble in excess of alkali, but—before oxidation—is soluble in solution of ammonium salts, by formation of soluble double salts of ammonium and man- ganese—ammonio-manganous salts—(corresponding to those of ammonium and magne- sium ; compare 76). And hence, ammonium hydroxide precipitates but part of the manganese in solution, forming in the reaction a salt of ammonium, which holds the rest of the manganese.from precipitation. The manganic hydroxide is not only insoluble in ammonium salts, but it is formed and precipitated from the ammoniacal solut ion of manganous hydroxide in salts of ammo- nium by action of the air. After standing, all the manganese is so precipitated, dark brown ; this precipitation by action of the air upon solution in ammonium salt being pe- culiar to manganese. As free ammonia facilitates the oxidation of metallic copper and of eobaltous salts, it may here promote the oxidation of the manganous compounds; also, it neutralizes the acid which would otherwise be set free. Manganese. 57 243. Hydrosulphurie acid, precipitates manganous acetate but imper- fectly, and nut in presence of acetic acid, and does not precipitate other salts, as manganous sulphide is soluble in very dilute acids, even acetic acid. Ammonium sulphide precipitates from neutral solutions, and forms from the recent hydrate of mixtures made alkaline the fiesh-colored man- ganous sulphide, MnS. Acetic acid, acting on the precipitated sulphides, separates manganese from cobalt and nickel, and from the greater part of zinc. 244. Alkali carbonates precipitate won/tmtms carbonate, MnC03, white, oxidized by the atmosphere to the brown manganic hydroxide (243), and, before oxidation, somewhat soluble in solution of ammonium chloride. 245. Alkali phosphates—as Wa2HP04—precipitate, from neutral solu- tions of manganous salts, normal manganous phosphate, Mn3(P04)2, white, slightly soluble in water, and soluble in dilute acids. It turns brown in the air. The manganous hydrogen phosphate—MnHPOj— is more soluble in water, and is obtained by crystallization from a mixture of manganous sulphate acidulated with acetic acid, and disodium hydrogen phosphate, Na .HP04, added till a precipitate begins to form. From the ammonio-manganese solution (343) phosphates precipitate all the man- ganese as ammonium-manganous phosphate. 246. Alkaline oxalates precipitate manganous oxalate, soluble in acids not very dilute, and formed with difficulty by addition of oxalic acid. 247. Soluble cyanides—asKCN—precipitate manganous cyanide, Mn(CN)2, white, but darkening in the air, soluble in excess of the precipitant by formation of double cyanides—as K2Mn(CN)4. This solution, exposed to the air, prod uces mangan icy an ides ■—analogous to ferricyanides—with oxidation of a portion of the manganese : 12K.,Mn(CN)4 + 302 + 2H_,O - 4K6Mn2(CN)I2 + 3Mn,02(0H)2 Perrocyanid.es precipitate white manganous ferrocyanide, Mn2Fe(CN)6, soluble in hydrochloric acid. Perricyanides precipitate brown manganous ferriryanide, Mn3- Pe2(CN)12, insoluble in acids. 248. Manganese is most easily and certainly identified through oxida- tion, by several methods, each method giving a color-product. A small portion of manganous solution, when boiled with nitric acid or dilute sulphuric acid and lead dioxide, or Pb304, is oxidized to per- manganic acid, H„MnO0B, giving a red color to the solution when the sedi- ment subsides. The oxidation is derived from the lead dioxide, reduced to a lead salt. If other reducing agents are present, they also must be oxi- dized. The lead dioxide should be used in such excess as to leave a black sediment. An excess of MnO„ must not be mistaken for an excess of *bO, 249. Ignition with alkali and oxidizing agents, forming a bright green mass of alkaline manganate, constitutes a delicate and convenient test for Manganese, in any combination. A small portion of precipitate or fine pow- der is taken. If the manganese forms but a small part of a mixture to be 58 MANG ANESE. tested, it is better to submit the substance to the systematic course of anal- ysis, and apply this test to the precipitate by alkali, in the third group. A convenient form of the test is by ignition*on platinum foil with potas- sium or sodium nitrate and sodium carbonate (a). Ignition, by an oxi- dizing flame, on platinum foil, with potassium hydroxide, effects the same result, less quickly and perfectly (b). Ignition by the oxidizing flame of the blow-pipe, in a bead of sodium carbonate, on the loop of platinum wire, also gives the green color (c). a ?Mn(OH)2 + 4KN03 -f Na2C03 = 2K2MnO4 + Na2Mno4 + 4NO + CO2 + 3H20 h. Mn(OH)2 + 2KOH +O2 = K2MnO, + 2H20 c. Mn(OH)2 + Na2C03+02 = Na„MnQ4 + H.O + C02 250. With beads of borax and microcosmic salt, before the outer blow- pipe flame, manganese colors the bead violet while hot, and amethyst-red when cold. The color is due to the formation of manganic oxide, the color- ing material of the amethyst and other minerals, and is slowly destroyed by application of the inner flame, which reduces the manganic to manganous oxide. 251. MANGANIC SALTS are somewhat instable compounds, of a reddish-brown or purple red color, becoming paler and of lighter tint in reduction to the manganous combination. The chloride and sulphate are deliquescent. Manganic chloride, Mn2Cl6, exists only in solution, which is reduced to MnCl2 by boiling, also by evaporation to a solid. Manganic sulphate -Mn«(SO4)3—is soluble in dilute sulphuric acid, but is reduced to MnSC4 by the attempt to dissolve it in water alone, see 231 ; potassium manganic sul- phate and other manganic alums are also decomposed by water. 252. Hydrosulphuric acid reduces manganic salts to the manganous combination, with precipitation of sulphur. Ammonium sulphide reduces manganic chloride, and precipitates manganous sulphide—MnS—with free sulphur. Alkali hydroxides, car- bonates, and barium carbonate, all precipitate from solution of manganic chloride, man- ganic hydroxide, Mn202(OH)2. Ferrocyanides precipitate gray-green manganic ferro- cyanide, Mn4(Fe(CN)6)3. Ferricyanides precipitate manganic ferricyanide—MnsPes- (CN)i2—brown. When a manganic compound is mixed with aqueous phosphoric acid, the solution evaporated to dryness and gently ignited, a violet or deep blue mass is ob- tained, from which water dissolves a purple-red manganic hydrogen phosphate, a distinc- tion from manganous compounds Simple ignition changes manganic compounds to Mn3o4. In the tests in the dry way, manganic compounds give the same reactions as manganous oxide (249). 253. Estimation.—(l) By converting into Mn3G4 (see 232), and weigh- ing as such. (2) By precipitating as a sulphide and weighing after ignition in hydrogen. (3) By precipitating as MnWH4P04, and after ignition weighing as Mn2P2Q7. (4) By addition of K6Fe2(CN)]2 in presence of KOH (and a trace of Fe2(OH)6), which converts the manganese into MnO„, and the resulting K4Fe(CN)6 is estimated by K2Mn2Og. (5) By treating the manganous salt with a solution of K2Mn208 of known strength. If some Cobalt. 59 ZnS04 is added the action is more satisfactory (Wright and Menee : Jour. Chem. Soc., 1880, 37, 42). 3MnS04 + K2Mn208 + 2HoO = 5MnOa + KaS04 + 3H2S04 254. Peroxide of manganese is estimated (1) by treating with H2C204, and measuring or weighing the CO,2 which is produced (246); (2) by boil- ing with HCI and estimating the evolved chlorine ; (3) by boiling with Pb02 and HNOs, and comparing the color with a permanganate solution of known strength (Peters : Chem. Neivs, 33, 35). A remarkable number of other methods have been recommended. 255. Oxidation.-—Mn" is oxidized to Mnlv in alkaline mixture by Cl, Br, I, K6Fe2(ON)12, KGIO, KBrO, etc. In presence of HNOs by Pb02, Pb304, HCI03, H2Mn2Og, etc. 256. All compounds of manganese having more than two bonds are re- duced to Mn" by H2C204, H3P02, H2S, H2S03, HCI, HBr, HI, HONS, (Hg2)", Sn", As'", Sb'" (Cu )", Fe", Cr", etc. For oxidizing action of K2Mn04 and K2Mn208 see Part 11. COBALT. Co = 58.887. 257. Specific gravity of powder reduced by hydrogen varies from 8.132 to 9.495; mean of five samples, 8.957 (Rammelsburg, 1849). Melting point above that of qold and below that of iron (Clarke): 1500° 0. (2712° F.), Pictet, 1879. 258. Occurrence.—Cobalt does not occur in a free state, except in me- teoric iron. It. is found in linnaeite (Co3S4); skutterudite (CoAs3); speiss cobalt (CoNiFeAsJ; glance cobalt (CoFeAsSj; wad (CoMnO,3MnOs + 4H„0), etc. 259. Preparation.—(l) By electrolysis of the chloride. (2) By heating with potassium or sodium (too expensive for commercial production). (3) By heating any of the oxides hydroxides or the chloride in hydrogen gas. (4) By fusion of the oxalate under powdered glass. (5) Also reduced by carbon in various ways. 260. Properties.—Cobalt in fine powder oxidizes rapidly in the air, and when made according to (3) sometimes takes fire spontaneously, like iron when reduced by hydrogen. When in a compact mass is permanent in dry air, and is only slightly tarnished in moist air. At a white heat it burns- tapidly. Can be made magnetic, but, unlike steel, retains its magnetism at a white heat. It is harder than iron, malleable and ductile, cobalt wire being about twice as strong as iron wire. It is soluble in most dilute acids. 261. Oxides and Hydroxides.—Cobaltous oxide (CoO) is made (1) by beating any of its oxides or hydroxides in hydrogen to (not above) 350° C.; (2) by ignition of Co(OH)2 or CoC03, air being excluded ; (3) by heating Cobalt. Co304 to redness in a stream of C02 (Russell, 1863, Jr. Client. Soc., [3], 1, 51) ; (4) by beating any of tlie higher oxides to a white heat (0. D. Braux, 1867, Zeit. Anal. Ghent., 6, 76). Cobaltous hydroxide is made from cobalt- ous salts by precipitation with fixed alkalies ; oxidizes if exposed to the air (264). The most stable oxide is the cobaltoso-cobaltic or tri-cobaltic tetrox- ide (Oo.O,); is made by heating any of the oxides or hydroxides, the car- bonate, oxalate, or nitrate to a dull red heat in the air or in oxygen gas. Several oxide hydroxides are known—e.y., Co302(OH)4, C030(0H)6, Co303 (OH),, Cobaltio oxide, Co2Os, is made by heating the nitrate just hot enough for decomposition, but not hot enough to form Co304. Cobaltic hydroxide, Co2(OH)6, is made by treating any cobaltous salt with Cl, HCIO, Br, or lin presence of a fixed alkali or alkaline carbonate. It dis- solves in HCI, with evolution of Cl, in H2S04, with evolution of oxygen, forming a cobaltous salt. 262. Solubilities.—Cobaltous nitrate and acetate are deliquescent; chloride, hy- groscopic ; sulphate (7 aq ), efflorescent. The chloride vaporizes, undecomposed, at a high temperature. The hydroxide, basic carbonate, sulphide, phosphate, borate, oxalate, cyanide, ferro- cyanide, and ferricyanide are insoluble in water ; the potassio-cobaltous oxide is inso- luble ; the ammonio-cobaltous oxide, soluble ; the double cyanides of cobalt and the alkali metals are soluble in water. Alcohol dissolves the chloride and nitrate ; ether dis- solves the chloride, sparingly. Most of the salts insoluble in water form soluble com- pounds with ammonia. In analysis, cobalt is pretty clearly identified in the dry way, by the bead test (2(19). 263. Reactions of Cobaltous Salts.—Cobaltous oxide is gray-green, the hydroxide is rose-red ; they are easily soluble in acids forming COBALTOUS SALTS, which ex- hibit bright colors, varied by different physical states, and by different chemical combi- nations. In crystals, they are red ; anhydrous, mostly lilac. Their solutions are mostly blue when concentrated, but pink when diluted. At a certain stage of dilution, these solutions are red when cold, and blue when hot. The pink dilute solution of the chloride spreads colorless on white paper when cold, becomes blue on heating, and colorless -when cold again, used as “ sympathetic ink.'’ Cobaltous oxide dissolves in meltedglass, coloring it blue—used to cut off the light of yellow flames ; also, with the same color, in fused borax—the most delicate test for cobalt (269), and in other vitreous substances. The black, cobaltoso-cobaltic oxide, Co304—as left by ignition of cobaltous oxide or nitrate —combines or mixes, by ignition, with zinc oxide from zinc compounds to form a green mass, with aluminium compounds to a blue, and with magnesium compounds to a pink mass. Cobalt forms many double salts, and compounds with alkalies, noted for their various bright colors. 264. The fixed alkalies precipitate, from solutions of cobaltous salts, blue basic salts, which absorb oxygen from the air and turn olive-green, as cobaltoso-cobaltic hydroxide; or if boiled before oxidation in the air, become rose-red, as cobaltous hydroxide, Co(OH)2. This last result is favored by excess of the reagent, which does not redissolve the preci- pitate. But ammonia and ammonium salts dissolve the precipitate. Ammonium hydrate causes the same precipitate as feed alkalies ; incomplete, even at first, because of its solubility in the ammonium salt formed in the reaction, and solu- ble in excess of the ammonia to a solution which turns brown in the air by combination Cobalt. 61 with oxygen, and is not precipitated by potassium hydrate. The reaction of the preci- pitate with ammonium salts forms a soluble double chloride (as with magnesium) ; the reaction of the precipitate with ammonia produces, in different conditions, different soluble color compounds, ammonio-cobaltous and ammonio-cobaltic, as (NH3).,C0C12, (NH3)6CoCI2, (NH3)BCo2CI6, etc. 265. Alkali carbonates precipitate cobaltous basic-carbonate, peach red, which when boiled loses carbonic anhydride and acquires a violet, or, if the reagent be in excess, a blue color. The precipitate is soluble in ammonium carbonate (or in excess of that precipitant), and very slightly soluble in fixed alkali carbonates. Carbonates of Ba, Sr, Ca, and Mg do not precipitate cobaltous salts in the cold (except the sulphate), but by prolonged boiling they precipitate cobaltous chloride com- pletely. 266. Hydrosulphuric acid, with normal cobaltous s Its, gradually and imperfectly precipitates the black cobalt sulphide, CoS; from cobalt acetate, the precipitation is more prompt, and is complete; but in presence of mineral acids, as in the second-group precipitation, no precipitate is made. When formed, the precipitate is scarcely at all soluble in dilute hydrochloric acid or in acetic acid ; slowly soluble in moderately con- centrated hydrochloric acid, as in dissolving the third-group precipitate ; readily soluble in nitric, and most easily in nitro-hydrochloric acids. By exposure to the air, the recent cobaltous sulphide is gradually or slowly oxidized to cobalt sulphate, soluble, as occurs with iron sulphide. Ammonium sulphide precipitates immediately and perfectly the black CoS, described above. 267. Phosphates—as Na2HP04—precipitate the reddish cobaltous phosphate, CoHPOi, soluble in acids and in ammonia. Oxalic acid and oxalates precipitate the reddish white, cobaltous oxalate, C0C204, soluble in mineral acids and in ammonia. 268. Alkali cyanides—as KON—precipitate the brownish-white cobaltous cyanide, Co(ON)2, soluble in hydrochloric, not in acetic or in hydrocyanic acid, soluble in excess of the reagent, as double cyanides of cobalt and alkali metals—(KCN)2Co(CN)2—potas- sium cobaltous cyanide, etc., the solution having a brown color: CoCl2 + 2KCN = CofCN)2 + 2KCI Co(CN)2 + 2KCN = (KCN)2Co(CN)2 Dilute acids, without digestion, reprecipitate cobaltous cyanide from this solution (the same as with nickel, 280): But if the solution, with excess of the alkali cyanide and with a drop or two of hy- drochloric acid, insuring free HCN, be now digested hot for some time, the cobaltous cy- anide is oxidized and converted into alkali cobalticyanide—as K6Co2(CN)i2—correspond- ing to ferricyanides, but having no corresponding nickel compound: (KON)2Co(CN)2 + 2HCI = Co(CN)2 + 3HCN + 3KOI 4Co(CN)2 + 4HCN + 02 = 3002(CN)6 (cobaltie cyanide) + 2H.0 Co2(CN)6 + 6KCN =. KOCo2(CN)i2 potassium cobalticyanide. In the latter solution acids cause no precipitate {important distinction from nickel, whose solution remains (KON)2Ni(CN)2, and after digestion as above is precipitated with acids). Sulphocyanate, in highly concentrated solution, gives a blue color, Co(CNS)2, crys- tallizable in blue needles, soluble in alcohol, not in carbon disulphide. In less concen- trated solutions, the color appears on warming. In neutral solutions, -nickel, iron, man- ganese, and zinc do not interfere (Schoenn, 1870). Ferrocyanides—as K,Fe(CN)6—precipitate cobaltous ferrocyanide, Co2Fe(CN)6, gray-green, insoluble in acids. Ferricyanides—as K6Fe2(CN) i2—precipitate cobaltous ferricyanide, Co3Fe2(CN)i2, brownish-red, insoluble in acids. But a more distinctive Cobalt. test is made by adding ammonium chloride and hydroxide, with the ferricyanide, when a blood-red color is obtained, in evidence of cobalt. If, in this test, manganese be pre- sent, a white precipitate is obtained at once, becoming brown with more ferricyanide; if nickel be present, a copper-red precipitate forms on boiling; zinc gives no precipitate, hot or cold, but on addition of ferrocyanide to the same solution, gives a white precipi- tate (Allkn, 1871). 269. In the bead of borax, and in that of microcosrnic salt, with oxidizing and with reducing flames, cobalt gives an intense blue color. The blue bead of copper changes to brown in the reducing flame. If strongly saturated, the bead may appear black from in- tensity of color, but will give a blue powder. This important test is most delicate with the borax bead. If sulphur or arsenic is present, it must be previously expelled by roast- ing. If manganese, copper, nickel, or iron is present, the continued application of the reducing flame will destroy the interfering color, and bring out the blue of cobalt. By ignition, with sodium carbonate on charcoal or with the reducing flame, com- pounds of cobalt are reduced to a magnetic mass. 270. Cobaltous compounds are oxidized to cobaltic combinations, in' the following tests, all of which distinguish cobalt from nickel, which is scarcely capable of higher oxi- dation. Potassium nitrite, KNOj, added to a somewhat concentrated solution of cobaltous salt, with addition of sufficient acetic acid, after warm digestion, on standing some time, better for twenty-four hours, causes a yellow crystalline precipitate of potassium cobaltic nitrite (a separation from nickel): 2C0C12 + 12KNO, + 3H02H302 + H2Q = (KNOj),, Co2O(NOn)4, (H20)2 + 4KCI + 2KC2H300 4- 2NO Chlorine gas, passed into dilute cobaltous solutions, changes them to cobaltic com- binations, which are then precipitated by digestion with barium carbonate in the cold (compare 265). Lead dioxide, with warm digestion, precipitates from neutral solutions all the cobalt, as cobaltic oxide hydroxide. Zn(OH)2, Pb(OH)2, and HgO precipitate Co(OH)2 from CoCl2 at 100° C. 271. Estimation.—(1) As metallic cobalt, all compounds that may be reduced by ignition in hydrogen gas—e.g., CoCl2, C0(N03)2, CoCOs, and all oxides and hydroxides. (2) As CoO, all soluble cobalt salts, all salts whose acids aye expelled or destroyed by ignition, all oxides and hydroxides. The salt is converted into Co(OH)2 by precipitation with a fixed alkali, and ig- nited in a stream of C02. The carbonate and nitrate may be ignited di- rectly in C 0.,, and organic salts are first ignited in the air until the carbon is oxidized, and then again ignited in C02. (3) After converting into a sul- phate it is ignited at a dull red heat and weighed as a sulphate. (4) After converting into the oxalate titrating with K.Mn.O^. 272. Oxidation.—Co" is oxidized to (Co„)VI in presence of a fixed alkali by PbO,2, Cl, KCIO, Br, KBrO, and I; in presence of acetic acid by KNO,, (270). (Co 2)VI is reduced to Co" by H2C204, H3P02, H2S, H2SOs, HCI, HBr' and HI. Metallic cobalt is precipitated from solution of CoCl2 by Zn, Cd, and Mg. Nickel. 63 NICKEL. Ni = 57.928. 273. Specific gravity, reduced by carbon, 8.900 (Schroder, 1859). Melting point, 1450° 0. (2642° F.) (Pictet, 1879). 274. Occurrence.—Nickel almost always occurs in nature together with cobalt. It is found ns kupfernickel, NiAs, millerite or nickel blende, NiS, etc. 275. Preparation.—(l) By electrolysis. (2) By heating in a stream of hydrogen. The oxide is reduced in this manner at 270° C. (W. Muller, 18G9 : Ann. Ohim. Phyl36, 51). (3) By fusing the oxalate under pow- dered glass (CO2 being given off). (4) Reduction by igniting in CO. (5) Re- duction by fusing with carbon in a variety of methods. 276. Properties.—A hard white metal; malleable and ductile, making a stronger wire than iron ; does not oxidize in the air at ordinary tempera- tures. It is much used in plating other metals. It is magnetic, but loses its magnetism like steel by heating to redness. Soluble in dilute HCI and H.SOt, hydro gen being given off ; in concentrated H2S04, SO., being formed ; in HNOs, NO being evolved ; in concentrated HNOs it becomes passive like iron. 277. Oxides and Hydroxides.—Nickelous oxide is formed when the carbonate, nitrate, or any of its oxides or hydroxides is strongly ignited. Nickelous hydroxide is formed by precipitation of nickelous salts with fixed alkalies. Nickelic oxide, Ni203, is made from NiCOs, Ni(NO3)2, or NiO by heating in the air not quite to redness, with constant stirring. It is changed to NiO at a red heat. Nickelic hydroxide, Ni2(OH)6, is formed by treating nickelous salts first with a fixed alkali and then with Cl, NaCIO, Br. or NaßrO (not formed by iodine). A black powder, forming no correspond- ing salts. Soluble in HCI, HBr, and HI, with separation of Cl, Br, and I. Soluble in HNOs H„S04, and in most non-reducing acids with evolution of oxygen ; in H,C204, C02 being evolved, and in each case a nickelous salt being produced. A trinickelic tetroxide, Ni304 (corresponding to Co304, Pe304, Mn3C4, and Pb3C4), is formed, according to A. Baubigny (1878, CJom.pt. Read., 87, 1082), by heating NiCl2 in oxygen gas at from 350° to 440° 0. 278. Solubilities.—The salts of nickel have a delicate green color in crystals and in solution; when anhydrous, they are yellow. The nitrate and chloride are deliquescent or efflorescent, according to the hygroraetric state of the atmosphere; the acetate is ef- florescent. The chloride vaporizes at high temperatures. The hydroxide, carbonate, sulphide, phosphate, borate, oxalate, cyanide, ferroeyan- -I{le, ferrieyanide, insoluble in water. The compounds of the oxide with potassium oxide and sodium oxide are insoluble; that with ammonia is soluble; and the double cyanides °f nickel and alkali metals are soluble in water. The chloride is soluble in alcohol, and the nitrate in dilute alcohol. Most salts of nickel form soluble compounds by action of ammonium hydrate. In analysis nickel is separated, with cobalt, by the sparing solu- 64 Nickel. bility of the sulphide in dilute acids. Its separation from cobalt is more difficult (380), In absence of cobalt, it is easily identified in the bead (283). 279. Reactions of Nickelous Salts.—The fixed alkali hydroxides precipitate niekelous hydroxide (Ni(OH)2), pale green, insoluble in excess of the reagent and not oxi- dizable in the air, but soluble in ammonium hydroxide or ammonium salts to a greenish blue liquid (158). Ammonium hydroxide, also, precipitates nickel hydroxide, soluble in excess, and in ammonium salts, with formation of compounds similar to those of cobalt (364), giving a violet-blue color to the solution. Sufficient potassium or sodium hydroxide will slowly re- precipitate nickel hydroxide from its ammoniacal solution, a distinction from cobalt. In dilute ammoniacal solutions, the blue color appears only after exposure to the air. The alkaline carbonates precipitate basic carbonate of variable composition, green color, and soluble in ammonium carbonate, or excess of that precipitant—with blue or greenish-blue color. Carbonates of Ba, Sr, Ca, and Mg precipitate on boiling the whole of the nickel from NiCl2. With hydrosulpliuric acid, and with sulphide of ammonium, nickel has the same deportment as cobalt (366): the precipitate being niciel sulphide, slightly soluble in ex- cess of ammonium sulphide. Phosphates—as Na2HPO4—throw down nickel phosphate, greenish-white, mostly full metallic. 280. Alkali cyanides—as KCN—precipitate nickel cyanide, Ni(CN)2, yellowish- green, insoluble in hydrocyanic acid, and in cold dilute hydrochloric acid; dissolving in excess of the cyanide, by formation of soluble double cyanides—as potassium nickel cy- anide (KCN)2Ni(CN)2. The equation of the change corresponds exactly to that for co- balt (368); and tiie solution of double cyanide is reprecipitated as Ni(CN)2 by a careful addition of acids (like cobalt); but hot digestion, with the liberated hydrocyanic acid, forms no compound corresponding to cobalticyanides, and does not prevent precipitation by acids (unlike cobalt). It will be observed that excess of hydrochloric or sulphuric acid will dissolve the precipitate of Ni(CN)2. Ferrocyanides—as K4Fe CN|,—precipi- tate a greenish-white nickel ferrocyanide, Ni .Fe(CN),;, insoluble in acids, soluble in am- monium hydroxide, decomposed by fixed alkalies. Ferricyanides precipitate greenish- yellow nickel ferricyanide. For the test by ferricyanide, with ammonium chloride and hydroxide, in distinction from cobalt, see 368. Oxalic acid and oxalates precipitate, very slowly, but almost completely, after twenty-four hours, nickel oxalate, green. 281. Chlorine, or hypochlorite, in neutral solution, or, better, with fixed alkali hy- droxide, forms a black precipitate of nickelic hydroxide. Ni2(OH)6, reduced by heat or by solution in acids or in ammonium hydroxide. The separation of nickel from cobalt (370), by this test, is more accurate if potassium cyanide in excess be added previously to the chlorine or hypochlorite. Nitriles, with acetic acid, do not oxidize nickel as they do cobalt. 282. Nickel compounds dissolve clear in the borax bead, giving with the oxidizing flame a purple-red or violet color while hot, becoming yellowish-brown when cold; with the reducing flame, fading to a turbid gray, from reduced metallic nickel, and finally becoming colorless. The addition of any potassium salt, as potassium nitrate, causes the borax bead to take a dark purple or blue color, clearest in the oxidizing flame. With microcosmic salt, nickel gives a reddish-brown bead, cooling to a pale reddish-yellow, the colors being alike in both flames. Hence, with this reagent, in the reducing flame, the color of nickel may be recognized in presence of iron and mnngane-e. which are color- less in the reducing flame; but cohalt effectually obscures the bead-test for nickel. The BING. 65 yellow-red of copper in the reducing flame, persisting in beads of microcosmic salt, also masks the bead-test for nickel. By ignition with soda on charcoal, compounds of nickel are reduced to a powder at- tracted by the magnet. 283. Estimation.—ln gravimetric determinations nickel is converted into NiO, and after intense ignition weighed as such. 284. Oxidation.—Ni" is changed to (Ni2)VI in presence of fixed alkalies by Cl, NaCIO, Br, and NaßrO (not by I). (Ni2)VI is reduced to Ni" by all non-reducing acids with evolution of oxygen ; by reducing acids, H2C204 is oxidized to C02, HNOs to HNOa, H3P02 to H3PC4, H2S to S, H2S03 to H2S04, HCI to Cl, HBr to Br, HI to I, HCyS to HCy and H2S04, H4Fe- (CN)6 to H6Fe2(CN)12. A solution of NiCl2 is reduced to the metallic state by zinc-dust and by finely divided cadmium or tin. ZINC. Zn= 64.9045. 285. Specific gravity, 7.14 (Rammelsbukg, 1880). Melting point, 450° C. (842° F.) (Pictet, 1879). Boiling point, 930° 0. (1706° F.) (Yiolle, 1882). 286. Occurrence.—lt is found as calamine (ZnCO3), as zinc-blende (ZnS); also associated with other metals in numerous ores. 287. Preparation.—The process usually employed consists of two opera- tions : (1) Roasting; in case of the carbonate the action is ZnC03 = Zno + CO>2; if it is a sulphide, 2ZnS + 3C2 = 2ZnO -\- 2S02. (2) Reduction with distillation ; after mixing the ZnO with one-half its weight of pow- dered coal it is distilled at a white heat. Its usual impurities are As, Cd, Pb, Cu, Pe, and Sn. It is purified by repeated distillation, each time re- jecting the first portion, which contains the more volatile As and Cd, and the last, which contains the less volatile Pb, Cu, Pe, and Sn, Strictly chemi- cally pure zinc is best prepared from the carbonate which has been purified by precipitation. Other methods of preparing zinc not financially profitable are : (1) Elec- trolysis ; (2) reduction by K, Na, or Mg; (3) reduction by hydrogen; (4) reduction by carbon monoxide. 288. Properties —A bluish-white metal ; retains its lustre in dry air, and is only slightly tarnished in moist air or in water. It is more malleable between 100° and 150° C. than at other temperatures, and may then be drawn into wire or rolled into sheets. At 205° 0. it is so brittle that it may be easily powdered in a mortar. For use in the laboratory it is usually granulated by pouring it when nielted into cold water. The water is then poured off and the zinc thrown into a solution containing about one grain of platinic chloride to a gallon of Zinc. water. Metallic platinum is precipitated upon the zinc (3Zn + PtCl4 = Pt + 3ZnCl2), and increases its solubility in H2S04. Otherwise (that is, if strictly chemically pure), it, would be unfit for use in Marsh’s test for arsenic. 289. Oxide and Hydroxide.—Zinc oxide (ZnO) is made by igniting in the air either metallic zinc, its hydroxide, carbonate, nitrate, oxalate, or any of its organic oxysalts. Zinc hydroxide, Zn(OH)2, is made from solu- tions of zinc salts by precipitation with fixed alkalies (393). 290. Solubilities.—Pare zinc dissolves very slowly in acids or alkalies, unless in contact with copper, platinum, or some less positive metal. The metallic impurities in ordinary zinc enable it to dissolve easily with acids or alkali hydroxides. In contact with iron, it is quite rapidly oxidized in water containing air, but not dissolved by water, unless by aid of certain salts. All the agents which dissolve the metal, dissolve also its oxide and hy- droxides. The metal dissolves in hydrochloric, sulphuric, and acetic acids (a), and in the aqueous alkalies (/;)—with evolution of hydrogen ; in very dilute ni- tric acid, without evolution of gas (c); in moderately dilute cold nitric acid, mostly with evolution of nitrous oxide (d); and, in somewhat less dilute nitric acid, chiefly with evolution of nitric oxide (e). Concentrated nitric acid dissolves zinc but slightly—the nitrate being very sparingly solu- ble in nitric acid: a. Zn + H2S04 = ZnSOi + H2 h. Zn + 3KOH = K2OZnO +H„ c. 4Zn + IOHNOs - 4Zn(NO3)2 + NH.NO, + 3H20 d. 4Zn + IGHNO., = 4Zn(NO3)2 + N.O + 5H.0 o)a!, dark brown, changing in the air to titanic acid, H2Ti03 ; ammonium sulphide throws down the same precipitate, hydrosulphuric acid producing no change; calcium carbonate separates the hydroxide.—Ferric and cupric salts are re- duced to ferrous and cuprous compounds, and from salts of mercury, silver, and gold 80 Thallium. the metals are separated, by titanous salts, which are thereby changed to titanic com- pounds. b. Titanic salts are mostly insoluble in water, or decomposed by it with precipitation of titanic acid, H2O.Ti02 or H2TiO3. Of this compound there are two modifications, one soluble and one insoluble in hydrochloric and nitric acids; strong sulphuric acid dissolves both modifications; but the titanic sulphate is decomposed and precipitated on dilution, and the chloride on long boiling (distinctive). Titanic chloride, TiCl4, and nitrate, Ti(NO3)4, are permanently soluble in water.—From these. Alkalies and their carbonates and sulphides throw down the white voluminous titanic hydrate or titanic acid, insoluble in excess of the precipitants, and in ammonium salts; the same precipitate is produced by barium carbonate. Ferrocyanide of potassium gives a dark-brown precipitate of ti- tanic ferrocyanide; tannic acid, an orange precipitate. c. Titanates, as shown above, are not formed by treating titanic acid, even wdien re- cent, with aqueous alkalies, but are produced by fusion of titanic acid with alkalies or their carbonates So prepared, the neutral alkali titanates have a yellow color, and are decomposed by hot water with separation of insoluble acid titanates of the same bases, but soluble in acids as titanic salts. d. Compounds of titanium acids with microcosmic salt dissolve in the outer flame to a clear bead, pale yellow when hot, and colorless when cold. The strong reducing flame now turns the bead yellow while hot (reddish when cooling), and violet when cold (titan- ous oxide). If sulphate of iron be added, the bead by the inner flame is blood-red. In the borax bead the same reactions are obtained, less intense.—lgnition on charcoal with soda does not reduce titanium to the metallic state (distinction from tin). THALLIUM. T1 == 203.715. 314. Specific gravity of the wire, 11.91(Crookes, 1864). Melting point, 293.9° C. (561° F.) (Crookes, 1864). Is found in crookesite and in many varieties of iron and copper pyrites. Preparation.—(l) By electrolysis. (2) By reduction from its solutions by Zn or Al. (3) By fusion with KCH. (4) By fusion of the oxalate or with some other form of carbon. (5) Fu- sion in hydrogen gas reduces it with difficulty. It is a bluish-white metal, softer than lead, malleable and ductile ; tarnishes rapidly in the air ; may be preserved under water, which it does not decompose below a red heat ; soluble in H2S04 and HNO,, in HCI with great difficulty ; combines di- rectly with Cl, Br, I, P, S, Se, and precipitates from their solutions Cu, Ag, Hg, Au, and Pb in the metallic state. Reactions.—As a monad its compounds are stable, and not easily oxidized ; as a triad it is easily reduced to the univalent condition a. Thallious oxide, Tl2O, is black; on contact with water it forms an hydroxide, TIOH, freely soluble in water and in alcohol, to colorless solutions. The carbonate is soluble in about 20 parts of water ; the sulphate and phosphate are soluble; the chloride very sparingly soluble; the iodide insoluble in water. Hydrochloric acid precipitates, from solutions not very dilute, thallious chloride, TICI, white, and unalterable in the air. As a first-group precipitate, thallious chloride dissolves enough in hot water to give the light yellow precipitate of iodide, Til, on adding a drop of potassium iodide solution— the precipitate being slightly soluble in excess of the reagent. H2S precipitates the ace- Various Metals. 81 tate, but not the acidified solutions of its other salts. (NH4)2S precipitates TI2S, which, on exposing to the air, soon oxidizes to sulphate. Perrocyanides give a yellow precipi- tate, Tl4Pe; phosphomolybdic acid a yellow precipitate; and potassium permangan- ate, a red-brown precipitate, consisting in part of Tlj03. Chromates precipitate yellow, normal chromate; and platinic chloride, pale orange, thallium platinic chloric e, (TICI)2- PtCl4. Thallium compounds readily impart an intense green color to the flame, and one emerald-green line to the spectrum (the most delicate test). The flame-color and spectrum, from small quantities, are somewhat evanescent, owing to rapid vapori- zation. b. Thallic oxide, T1203, dark violet, is insoluble in water; the hydroxide, an oxy- hydroxide, TIO(OH), is brown and gelatinous. This hydroxide is precipitated from thal- lic salts by the caustic alkalies, and not dissolved by excess. Chlorides and bromides do not precipitate thallic solutions; iodides precipitate Til with I. Sulphides, and H2S, precipitate thallious sulphide, with sulphur. Thallic oxide, suspended in solution of potassium hydroxide, and treated with chlorine, develops an intense violet-red color. Thallic chloride and sulphate are reduced to thallious salts by boiling their water solu- tions. 315. YTTERBIUM. Yb = 172.761.—1n 1878 Marignac prepared the sulphate, Yb2(S04)3, the oxalate, Yb2(C204)3, the nitrate, and some other ytterbium salts. The metal has not been isolated. 316. TERBIUM. Tr = 148.5?—Terbium oxide, Tr2Q3, occurs in sa- marskite ; the sulphate, Tr2(S(D4)3, and a few other salts, have been pre- pared, but as the metal has not been isolated the composition of its salts is not certain. 317. SCANDIUM. Sc = 43.98.—Scandium oxide, Sc2Os, is found in gadolinite; the sulphate, Sc2(SC4)3, the oxalate, Sc2(C204)3, and a few other salts, have been prepared ; the metal lias not been isolated. 318. SAMARIUM. Sm = 148.801.—Discovered by Boiseaudran in samarskite by its peculiar spectrum. The oxide, Sm203, the chloride, SmCl3 or Sm2Cl6, the sulphate, the nitrate, and a few other salts, have been prepared. The metal has not been isolated. 319. GALLIUM. Ga = 68.584.—Specific gravity, 5.9. Melting point, 30.15° C. (86.37° F.) Found in zinc-blende. Prepared by electrolysis. It is a soft metal, may be cut with a knife, flexible and malleable, soluble in HCI and KOH, and in hot HNOs, The oxide, Ga203, the chloride, Ga2Cl6, and sulphate are best known. 320. DECIPIUM. Dp = 171.?—Discovered by Delafojsttinb in sa- marskite by means of the spectroscope. The oxide, Dp2Q3, Dp2(S04)8, Dp,(IO,)„ bp,(0.H.0,)„ and Dp2(C204)3 are best known. The metal has not been isolated. Rare Metals of the Third Group. 321. The leading reactions of the remaining rare metals of the Third Group are given in the following Comparative Table. The six metals first named form earthy oxides ; tantalum and niobium, like titanium, form acids. Indium and vanadium, also forming acids, can be placed in the Second Group. Zr Th Y E La D Ta Nb In V NH4OH with NH4C1. (NH4)5S KOH or NaOH, excess. K2C03, excess. (NH4)2COs, excess. BaC03, cold. k2so4 h2c2o4 K4Fe(CN)6 K6Fe2(CN)12 N a2S203 Zr(OH)4 Zr(OHi4 Pre. Sol. * t Sol. t Pre * Pre. Pre. ThO(OH)2 ThO(OHi2 Pre. Sol. t Sol. + Pre. Pre. t Pre. Wt. pre. No pre. Pre. * Y2(OHl6§ Y2(OHi. Pre. Pre. * Sol. * Pre. No pre. Pre. Wt. pre. No pre. No pre. Basic salt. Basic salt. Eg (OH), Pre. * Sol. * No pre- No pre. 1! Pink pre. La2(OH>6 La2(OH)8 Pre. If Pre. Pre. Pre. Pre. * Pre. Di2(OH)6* Pre. * Pre. Pre. Pre. Pre. Pre. Pre. Wt pre. HTa03 HTa03 tt Pre. ** Pre. ** tt ‘ In2(OH)6 In2Ss §§ Sol. t Pre. Sol. * Pre. ITT ill Green pre. No pre. Red pre. Yel. pre. Pre. * Wt. pre. Pre. No pre. Pre. * * In part. t Eeprecipittited on boiling. t In boiling, concentrated solutions, slowly but completely. § Slightly soluble in ammonium chloride. [ Partial pre. in the cold, dissolving when warm. ** Contains NH4OH, t When the acetate is precipitated by excess of ammonia, and the washed precipitate treated with iodine, a blue color appears. tt Fusion with KOH gives KTaOg, soluble in water; with NaOH, NaTa03. insoluble in solution of sodium hydroxide. The sol. of KTaOa is pre. by C02 ; also by HC1 (as Ta203) soluble in excess. tt Solutions of niobates, as KNb03, are pre. by mineral acids, as HNb03. In acidulated solutions, zinc forms a blue color (characteristic), lower oxide being formed. §§ Formed yellow, in slightly acid solutions, by H2S, soluble ir excess of (NH4)2S, and in HC1, 111 In neutral solutions, a brown color ; on acidulating, a brown pre., V2S5. H2S reduces vanadates to V204, with blue color. Acid solutions of V203 are green: of V205, red. For the characteristics of the several conditions of this element, see Watts’ Dictionary, Yols. Y. and VI., and Grnelinkraut's Handbuch, 2, 22T Groups I. and 11. 83 GROUPS I. AND 11. 322. The first group includes metals whose chlorides are insoluble in water (Pb, Ag, and (Hg2)"). The second group includes those metals whose chlorides are soluble in water, but whose sulphides are insoluble in dilute acids. Copper Cu = 63.173 Gold = 196.155 Platinum Pt - 194.415 Bismuth Bi = 207.523 Palladium Pd = 105.737 Cadmium Cd = 111.835 Ruthenium Ru = 104.217 Iridium Ir = 192.651 Lead Pb = 206.471 Rhodium Rh ; 104.055 Silver Ag = 107.675 Osmium Os = 198.494 Tellurium Te - 127.960 Mercury Hg - 199.712 Selenium Se = 78.797 Arsenic = 74.918 Tungsten W = 183.610 Molybdenum Mo = 95.527 Antimony Sb = 119.955 Norvvegium Ng : . 218.93? Tin Sn = 111698 Germanium. Gr = 72.32 323. Compakatiye View.—Owing to the partial solubility of lead chloride in water, it is never completely precipitated in the first group ; hence it must also be tested for in the second group. (Hg2)" belongs to the first group and Hg" to the second. Silver, then, is the only exclusively first-group metal. 324. The metals included in these groups are less strongly electro-posi- tive than those of the other groups. Only bismuth, antimony, tin, and mo- lybdenum decompose water, and these only slowly and at high tempera- tures. The oxides of silver, mercury, gold, platinum, and palladium are decomposed below a red heat. Copper, lead, and tin tarnish by oxidation in the air. In general, these metals either do not dissolve in acids with evolution of hydrogen, or do so with difficulty. Nitric acid is the best sol- vent for all, except antimony and tin, which are rapidly oxidized by it. 325. Mercury, arsenic, antimony, and tin form, each, two stable classes of salts. Therefore, the lower oxides, chlorides, etc., of these metals act as reducing agents ; and their higher oxides, chlorides, etc., as oxidizing agents, each to the extent of its chemical force. Arsenic, antimony, tin, molybdenum, and several of the rare metals of these groups enter into acid- ulous radicals, which form stable salts. Arsenic and selenium are metal- loids rather than metals. Arsenic, antimony, and bismuth belong to the Nitrogen Series of Elements. 326. A largo proportion of the compounds of these metals are insoluble in water. Of the oxides or hydroxides, only the acids of arsenic are soluble in water. The only insoluble chlorides, bromides, and iodides are in these groups. The sulphides, carbonates, oxalates, phosphates, borates, and cy- anogen compounds are insoluble. Most of the so-called soluble compounds 84 Coffer. of bismuth, antimony, and tin, and some of those of arsenic and mercury, dissolve only in acidulated water, being decomposed by pure water, with formation of insoluble basic salts. 327. The oxides of arsenic, antimony, and tin—in general terms—dis- solve in alkali hydroxides. Oxides of silver, copper, and cadmium dis- solve in ammonium hydroxide ; oxide of lead, in fixed alkali hydroxide. Metallic lead, like zinc, dissolves in fixed alkali hydroxide, with evolution of hydrogen, though it scarcely decomposes any acid by displacing hy- drogen. 328. Many double salts are formed with the metals of this group. Those whose sulphides dissolve in alkali sulphides, owe this property to the forma- tion of soluble sulpho-salts or double sulphides. Platinum forms a large number of stable double chlorides, soluble and insoluble ; and gold forms double chlorides, cyanides, etc. 329. Mercury, antimony, silver, and gold do not form hydroxides. The oxides of gold are very instable. 330. The metals of this group are all easily reduced to the metallic state by ignition on charcoal. Except mercury and arsenic, which vaporize, and certain rarer metals difficultly fusible, the reduced metals melt to metallic grains on the charcoal. Mercury and antimony vaporize from the liquid, arsenic from the solid state. COPPER. Cu = 63.173. 331. Specific gravity, variable : hammered, 8.9587 (Berzelius); crys- tallized, 8.94 ; electrolytic, 8,914 (Marc hand). Melting point, variable ; Yiolle finds it to be 1054° C. (1929° F.) Valence, a dyad in cupric salts, as Cu"Cl2, and a pseudo-monad in cuprous salts, as (Cu2)"Cl2. 332. Occurrence.—Copper is found native in various parts of the world, and especially in the region of Lake Superior. The must important English ore is copper pyrites, CuFeS2 ; copper glance is Cu2S ; green mala- chite is Cu2(OH)2C03 ; blue malachite is Cu3(OH)2(C03)2 ; red copper ore is Cu20 ; tenorite is CuO ;it is found in many other ores, and is utilized as a by-product in their reduction. 333. Preparation.—For the details of the various methods of copper- smelting the works on metallurgy should be consulted. In the laboratory pure copper may be produced (1) by electrolysis ; (2) reduction by igni- tion in hydrogen gas ; (3) reduction of the oxide by ignition with carbon, carbon monoxide, illuminating gas, or other forms of carbon ; (4) reduc- tion of the oxide by K or Na at a temperature a little above the melting point of these metals; (5) reduction by fusion with potassium cyanide (2CuO -|- 2KCN= 2Cu + 2KCbN'O). For its reduction in the humid way see 351. Copper. 85 334. Properties.—A red metal, but thin sheets transmit a greenish-blue light, and it also shows the same greenish-blue tint when in a molten con- dition. Of the metals in ordinary use only gold and silver exceed it in malleability. In ductility it is inferior to iron and cannot be so readily drawn into exceedingly fine wire. Although it ranks next to iron in tenacity, its wire bears about half the weight which an iron wire of the same size would support. As a conductor of heat it is surpassed only by gold. Next to silver it is the best conductor of electricity. Dry air has no action upon it ; in moist air it becomes coated with a film of oxide which protects it from further action of air or of water. 335. Oxides and Hydroxides.—Cuprous oxide (Cu2o) is found na- tive ; it is prepared (1) by reducing CuO by means of grape-sugar in alka- line mixture ; (2) by igniting CuO with metallic copper; (3) by treating an ammoniacal cupric solution with metallic copper, then adding KOH and drying. A red powder, soluble in HCI, forming Cu2Cl2. Cuprous hydroxide is formed by precipitating cuprous salts with KOH or NaOH. Cupric ox- ide is formed by igniting the hydroxide, carbonate, sulphate, nitrate, and some other cupric salts in the air. Cupric hydroxide is formed by precipi- tating cupric salts with KOH or NaOH. It is stated by H. Kose that te- tracupric monoxide (Cu4o) is formed by treating a cupric salt with KOH and a quantity of K„Sn02 insufficient to reduce it to the metallic state. A peroxide of copper (Cuo2) is supposed to be formed by treating Cu(OH)2 with H202 at 0° C. 336. Solubilities.—Copper does not dissolve in acids with evolution of hydrogen ; it dissolves most readily in nitric acid, chiefly with the evolu- tion of nitric oxide (a); also, in hot concentrated sulphuric acid, with evo- lution of sulphurous anhydride {b): a. BCu + BHNO3 = 3Cu(NOs)2 + 4H.,Q + 2NO b. Cu + 2H0504 = CIISO4 + 2HoO + SO. The atmosphere oxidizes copper very rapidly when in contact with sol- vents of the oxide of copper ; and in this manner the metal becomes oxi- dized and dissolved in hydrochloric acid and nearly all acids, in ammo- nium hydroxide, in solutions of many salts, in fats, sugars, and other or- ganic substances. Copper forms two oxides, and corresponding series of salts ; cuprous salts being infrequent and instable compounds, nearly all insoluble in water, and easily resolved into metallic copper and cupric salts, the stable and rep- resentative compounds of the metal. Cupric salts are readily reduced to cuprous combinations by most strong reducing agents acting with alkalies, as, by sulphites (a) with free alkali (difficultly, without alkali); by arsenious acid, with excess of alkali ; by glucose, and certain other sugars and organic materials, with excess of alkali. Also, by ferrous salts, in presence of iodides (345 b). Metallic iron 86 Coffer. and zinc separate, from solutions of cupric salts, metallic copper, without formation of cuprous salt. a. 2CuSO4 + 4KOH + SO2 = Cu2S04 + 2K2S04 + 2HaO 337. CUPROUS oxide—Cu20—is of a brownish-red color ; cuprous hy- droxide—Cu2(OH)2—brownish-yellow. Cuprous salts are insoluble in water. The chloride, Cu2Cl2, dissolves in strong hydrochloric acid to a colorless solution, which turns green in the air. From this solution water throws down the cuprous chloride, white; fixed alkalies, in small quantity, neutralize the free acid, and precipitate the white cuprous chloride ; in larger quantity, precipitate the yellow cuprous hydroxide, insoluble in excess. Am- monium hydroxide and ammonium carbonate, in excess, redissolve the hydroxide, and dissolve the oxide to a colorless solution, which turns blue on exposure. Potassa repre- cipitates the ammonia solution. Soluble carbonates precipitate the yellow cuprous car- bonate, Cu2C03.—lodide of potassium precipitates the white cuprous iodide, Cu2I2, without liberation of iodine (345 &).—Hydrosulphuric acid and sulphides precipitate Cu2S, black.—Phosphates, oxalates, cyanides, and ferrocyanides precipitate their re- spective cuprous salts, white; ferricyanides, brown-red. Ammoniacal solution of silver nitrate precipitates metallic silver.—With the blow-pipe, cuprous salts behave like cu- prous compounds (349). 338. Reactions of Cupric Salts.—Cupric suits, in crystals or solution, have a green or blue color; the chloride (2 aq.) in solution is emerald-gneu when concentrated, light blue when dilute ; the sulphate (5 aq.) is blue vitriol.” Anhydrous cupric salts are white. The crystallized chloride is deliquescent; the sulphate, permanent; the acetate, efflorescent. Cupric hydroxide, basic carbonate, oxalate, phosphate, borate, arsenite, sulphide, cyanide, ferrocyanide, ferricyanide, and tartrate are insoluble in water. The ammonio-oxide and most of the ammonio salts, the potas- sio and sodio cyanides, and the potassio and sodio tartrate, are soluble in water. In alcohol the sulphate and acetate are insoluble ; the chloride and nitrate, soluble. Ether dissolves the chloride. Copper is easily identified by reduction with iron to the lustrous metal- lic state (350) ; also, by the blue solution with excess of ammonium hydroxide (340), used as a separation from bismuth. 339. Fixed alkalies—KOH—added to saturation in solutions of cop- per salts, precipitate copper hydroxide, Cu(OH)s, deep blue, insoluble in excess, soluble in ammonium hydroxide (if too much fixed alkali is not present), very soluble in acids, and changed, by standing, to the black, basic hydrate, Cu30„(0H)2; by boiling, to CuO. If tartaric acid, citric acid, grape-sugar, milk-sugar, or certain other organic substances are present, the precipitate either does not form at all, or redissolves in excess of the fixed alkali to a blue solution. The tartrate alkaline solution may be boiled without change ; in presence of sugar, the application of heat pre- cipitates the yellow cuprous hydrate (346). The addition of alkali Copper. 87 hydroxides, short of saturation, forms insoluble basic salts, of a lighter blue than the hydrate. 340. Ammonium hydrate, added short of saturation, precipitates the pale blue basic salts ; added just to saturation, the deep blue hydroxide (in both cases like the fixed alkalies) ; added to supersaturation, the precipitate dissolves to an intensely deep blue solution. The blue solution consists of compounds of cuprammonium, (W2H6Cu)", a diammonium formed by the substitution of an atom of copper for an atom of hydrogen in each of two semi-molecules of ammonium, NH310I3Cu. The cuprammonium oxide is united with ammonium salt, as (W!H,Cu)O,(NHJ!SO. and (N,H.Cu)O.- (NH,OI), : CuSO4 + 4NH4OH = (N2HOCu)O.(NH4)2SO4 + 3H20 From this solution the fixed alkalies in strong solution precipitate the blue hydroxide, and on boiling the black oxide, CuO. 341. Ammonium carbonate, like ammonium hydroxide, precipitates and redissolves to a blue solution. Carbonates of fixed alkali metals—as K2C03—precipitate the greenish-blue, basic carbonate, Cu2(OH)2C03, of variable composition, according to conditions, and converted by boiling to the black, basic hydroxide and finally to the black oxide. Barium carbon- ate precipitates completely, on boiling, a basic carbonate. 342, Hydrosulphuric acid, and soluble sulphides, precipitate copper sulphide, CuS, black, formed alike in acid solutions (distinction from iron, manganese, cobalt, nickel) and in alkaline solutions (distinction from arsenic, antimony, tin).—Solutions containing only the one-hundred-thou- sandth of copper salt are colored brownish by the reagent. The precipitate, CuS, is easily soluble by nitric acid (a) (distinction from mercuric sul- phide); with difficulty soluble by strong hydrochloric acid (distinction from antimony) ; insoluble in hot dilute sulphuric acid (distinction from cad- mium) ; insoluble in fixed alkali sulphides, and but slightly soluble in am- monium sulphide (distinction from arsenic, antimony, tin) ; soluble in so- lution of potassium cyanide (b) (distinction from lead, bismuth, cadmium, mercury); soluble in solution of potassium carbonate. a. 6CuS + 16HNOs - 6Cu(NO3)2 + BS2 + BH2O + 4NO b. CuS + 4KCN = (KCN)2Cu(CN)2 + K2S 343. Phosphates—as Na„HPO4—give a bluish-white precipitate of cop- per phosphates; CuHP04, if the reagent is in excess; Cu3(P04)2, if the copper salt is in excess ; the precipitates slightly soluble by acetic acid.— Oxalates precipitate cupric oxalate, CuC.204, bluish-white, insoluble in acetic acid, and formed from mineral acid salts of copper by oxalic acid added with alkali acetates.—Normal potassium chromate precipitates brown-red basic cupric chromate, somewhat soluble in water.—Arsenites, as K3As03, or arsenious acid with just sufficient alkali hydroxide to neutralize it, pre- cipitate from solutions of cupric salts (not the acetate) the green copper 88 Copper. arsenite, chiefly CuHAs03 (Scheele’s green, “Paris green”), readily soluble in acids and in ammonium hydroxide, decomposed by strong potassium hydroxide solution. From cupric acetate, arsenites precipitate, on boiling, copper aceto-arsenite, (CuoAs203)3Cu(C2H302)2, Schwemfurt green or Im- perial green, “ Paris green,” dissolved by ammonium hydroxide and by acids, decomposed by fixed alkalies. 344. Alkaline cyanides—as KCN—precipitate at first the yellowish- green cyanide, Cu(CN)2, soluble in excess of the reagent by formation of potassium cupric cyanide, K2Cu(CN)4. The cupric cyanide precipitate is instable, becoming cuprous, or cuproso-cupric cyanide, Cn2(CN)s; the lat- ter unites with ammonium hydroxide, forming several green to blue salts, mostly soluble in water. Ferroeyanides—as K4Fe(CN)B'—precipitate the copper ferrocyanide, Cu2Fe(CN),, reddish-brown, insoluble by acihs, de- composed by alkalies. In highly dilute solutions a reddish coloration, without precipitate, is seen. Ferricyanides—as K6Fe2(Clsr)12—precipitate copper f'erricyanide, Cu3Fe2(CN)12, yellowish-green, insoluble in hydrochlo- ric acid. Thiocyanates, with sulphurous or hypophosphorous acid, preci- pitate cuprous thiocyanate, Cn2(COSrs)2, white (distinction from cadmium). 345. Soluble iodides precipitate, from concentrated solutions of copper salts, cuprous iodide, Cu2I2, white, colored dark brown by the iodine sepa- rated in the reaction (a). The iodine dissolves with color in excess of the reagent, or dissolves colorless on adding ferrous sulphate or soluble sul- phites, by entering into combination. Cuprous iodide dissolves in thiosul- phates (with combination). The cuprous iodide is precipitated, free from iodine, and more com. pletely, by adding reducing agents with iodides ; as, Wa2S03, H,SOs, FeS04 (h). a. 3CuSO4 + 4KI - Ou2I2 + I 2 + 3K2504 b. 3CuS04 + 2KI + 3FeS04 = Cu2l2 + K2S04 + Fe2(S04)3 2CuSO4 + 4KI + H2S03 + H.O = Cu.l2 + 3K,SO. + H2S04 + SHI 346. Metallic copper is reduced and separated from cupric solutions by iron, zinc, cobalt, nickel, lead, cadmium, bismuth, tin, and phosphorus. A bright slip of iron in solution of cupric salts acidulated with hydrochlo- ric acid, receives a bright copper coating, recognizable from solutions in 130,000 parts of water. Zinc acts most promptly in contact with platinum, as by use of a platinum dish, when the copper is deposited on the platinum; when minutely divided as a precipitate, the copper is dark brown to black. Finely divided zinc can be removed by solution in hydrochloric acid. Ni- tric acid and tartaric acid interfere with this reaction : CuSO4 + Fe = Cu + FeSO4 (For every 63 parts of copper deposited, 55.9 parts of iron are dissolved.) For detection of minute traces of copper, by metallic reduction, Hager directs as fol- lows : The material is obtained in solution acidulated with acetic acid. The end of a Coffer. 89 platinum wire is inserted just within the eye of a large sewing-needle, around which the wire is wound. The coil is left in the solution three or four hours, at a temperature of 25° to 30° C. (77° to 86° P.) The presence of copper is indicated by a black-brown coat- ing on the platinum wire, but more closely determined by further treatment. The needle is now withdrawn, the platinum wire is washed by gentle introduction into water, placed in a test tube, treated with four or five drops of nitric acid and a few drops of diluted sulphuric acid, warmed, boiled to expel all nitric acid, and an excess of ammonium hy- droxide added. Hydrobromie Acid. —To one drop of the copper solution add two drops of HBr, and concentrate by evaporation to one-half of its volume ; the rose-red color indicates copper. 0.01 of a milligramme of copper may thus be detected. Of the common metals only iron interferes. KBr and H2S04 may be substituted for the HBr. Arsenious acid, certain sugars, and many organic compounds, reduce cupric salts with fixed alkali hydroxide to a yellow precipitate of cuprous oxide and not to metallic cojrper. Sodium thiosulphate, Na2SiO:,, added to hot solutions of copper salts, gives a black precipitate. In solutions strongly acidulated (with hydrochloric acid) this is a separa- tion from cadmium. 347. Ignition with sodium carbonate on charcoal leaves metallic copper in finely divided grains. The particles are gathered by triturating the charcoal mass in a small mortar, with the repeated addition and decantation of water until the copper subsides clean. It is recognized by its color, and its softness under the knife. 348. Copper readily dissolves, from its compounds in beads of borax and of microcosmic salt, in the outer flame of the blow-pipe. The beads are green while hob, and hlue when cold. In tbe inner flame the borax bead becomes colorless when hot; the microcosmic salt turns dark green when hot, both having a reddish-brown tint when cold (Cu.,o) (helped by add- ing tin). 349. Compounds of copper, heated in the inner flame, color the outer flame green. Addition of hydrochloric acid increases the delicacy of the reaction, giving a greenish-blue color to the flame. 350. Estimation. —(1) It is precipitated by means of zinc in a platinum dish (or a small battery may be employed) and weighed as metallic copper. Or it maybe converted into CuH2 by treating with H3P02, and after reducing to metallic copper by ignition weighed as such. (2) It is converted into CuO and weighed after ignition. (3) It may be precipitated either by H„S or Ea,,S2- 03, and, after adding free sulphur and igniting in hydrogen gas, weighed as cuprous sulphide, or it may be precipitated by 'KCNS in presence of H2S03 or H3PO,„ and, after adding S, ignited in H and weighed as Cu2S. Cu„o, CuO, Cu(HO3)2,CuCO3, CuS04, and many other cupric salts, are converted 90 Bismuth. into Cu2S by adding S and igniting in hydrogen gas. (4) By adding KI to the cupric salt and titrating the liberated I by Ha2S2Os; not permissible with acid radicals which oxidize HI. (5) By precipitating as Cu, dissolving in Pe2Cl6, Ou + Fe2Cl6 = CuCl2 + 2FeClo and titrating the latter by K2Mns08. (6) By reducing with a solution of SnCl2 of known strength in presence of free HCI 2CuCl2 + Sno12 = Cu2C12 + SnCl4, the end of the reaction is known by the disappearance of the green color of CuCl2. 351. Oxidation.—Solutions of Cu" and (Cu2)" are reduced to the metal- lic state by Zn, Cd, Sn, Al, Pb, Pe, Co, Hi, Bi, and in presence of KOH by K2Sn02. Cu" is reduced to (Cu2)" by SnCl2 in presence of HCI, and in presence of KOH by As4Oe and by grape sugar. Metallic copper is oxidized to Cu" by solutions of Hg", (Hg2)", Ag', PtIV, and Au"', these salts being reduced to the metallic state. Copper is also oxidized by many acids. BISMUTH. Bi = 207 523. 352. Specific gravity, 9.759 (Schroder, 1859). Melting point, 268.3° C. (515° F.) (Rlemsdyk, 1869). Valence, a triad in Bi"'Cl3, a pentad in Biv 205, and in Bi202 a dyad. 353. Occurrence.—A comparatively rare metal, usually found native ; also found as bismuth ochre (8i.0.), bismuthite or bismuth glance (Bi2Ss), and in some other minerals. 354. Preparation.—On a large scale it is always reduced from its ores by fusing with carbon. It may also be produced by fusing with K, Ha, CO, KCH, H2C204, or HH4CI. 355. Properties.—A hard, brittle, reddish-white metal. Its melting point is lowered by alloying with other metals. Fusible metal consists of, in parts by weight, 2 Bi, 1 Sn, 1 Pb, and melts at 93.7° C., “ Wood’s Metal,” 15 Bi, 8 Pb, 4 Sn, 3 Cd, melts at 68° C. Bismuth may be distilled in an atmosphere of H above 1100° C. bismuth dioxide (8i202) is sometimes called hypobismuthous oxide ( Wails’ Dictionary, 1888); and more properly bisrnuthous oxide (Graham Otto’s Chemie, 1888). It is formed by the action of potassium stannite upon bis- muth hydroxide— -356. Oxides.—Four oxides are known—Bi202, Bi203, Bi204. Bi2Os. Di- A black powder easily oxidized to Bi'", A corresponding dibismuth tetra- chloride, Bi2Cl4, is formed when metallic bismuth is heated with BiCl3, Hg2Cl2, or Cl in the right proportions. It is instable. Bismuth peroxide, Bi2Os. is formed when Cl is passed into a hot solution of KOH, containing 28i(0H)3 + K2SnQ2 = Bi202 + K2SnG3 + 3H20. Bismuth. 91 Bi(OH),. It is a reddish powder. Non-reducing acids, such as H2S04, change it to Bi'", with evolution of oxygen. Reducing acids are them- selves oxidized, no oxygen being given off—e.g., BinOa + 10HCI = 2BiCls + 2C12 -l- 5H20. Dibismnth tetroxide, is little known ; formed like Bi206, but using less chlorine. It may be considered in its valence as a union of Bi'" and Bi% 2Bi204 = Bi203 + Bi205. Bismuth oxide and its corresponding salts are stable. It has a yellow- ish-white color (Bi(OH)3 is white); is formed when any other bismuth ox- ides, hydroxides, organic salts, or inorganic oxysalts with volatile acids, are ignited in the air. Bismuth hydroxide, Bi(OH)3, is formed by precipitat- ing bismuth salts with alkalies. 357. Solubilities.—Bismuth is but slightly oxidized in the air at ordi- nary temperatures, rapidly at a red heat ; it takes fire in chlorine, and unites readily with bromine, iodine, and sulphur. Hydrochloric acid scarcely attacks it ; boiling sulphuric acid salifies it with separation of sul- phurous anhydride, but it dissolves much the most readily in nitric acid, with evolution of lower oxides of nitrogen. The sulphide, hydroxide, basic carbonate, phosphate, chromate, borate, sulphite, oxalate, iodide (364), cyanide, ferrocyanide, ferricyanide, tartrate, citrate, tannate, and valerianate are insoluble in water. The chloride, bromide (364), nitrate, chlorate, and sulphate—when taken as normal salts—are soluble in water acidulated with their respective acids, or with other acids forming “ soluble” bismuth salts ; but are decomposed by pure water, with partial solution and partial separation of insoluble basic salts—(326 and equations in 358). The ammonio citrate is soluble in water without decomposition ; and the decomposition of the normal chlo- ride, nitrate, and sulphate is prevented by the addition of compara- tively small quantities of acetic, citric, and certain other organic acids. The acidulated, water-saturated solutions of the nitrate and chloride may be considerably further diluted with alcohol, without disturbance. In analysis, bismuth is precipitated alone, from the nitric acid solution of second-group sulphides, after removing lead (and silver), by adding ex- cess of ammonium hydroxide, a separation from copper and cadm um. The precipitation by water (358) suggests bismuth. 358. Water precipitates, from the acidulated bismuth solutions, white basic salts (see equations below), which contain less of their acid radicals in proportion as greater quantities of water are added, and some of which can be washed on the filter until almost pure hydroxide or oxide. The precipi- tation is most complete with the chloride, and with other salts .is promot ed by addition of hydrochloric acid or chlorides ; hence it may occur as a first- group precipitate. All the precipitates are readily soluble in hydrochloiie and nitric acids ; not in tartaric acid (distinction from antimony). Acidti- Bismuth. lation with certain organic acids (in accordance with the statement in 357) prevents the precipitation ; BIC13 + H.o = BiOOl + 2HCI 8i(N03)3 4- 2H20 = (a) BiON03.H20 + 2HNO-, 48i(N03)3 + 6H20 = (J) Bi40S(N03)2.H20 + 10HNO, Bi(NOs)3 + 3H20 = (c) Bi(OHj3 + 3HNOa 359. The alkali hydroxides precipitate from bismuth solutions—in absence of tartaric acid, citric acid, and certain other organic substances— the white bismuth hydroxide, Bi(OH)3, insoluble in excess of the reagents, converted by boiling to the oxide, Bi2Os, yellowish-white. Certain reducing agents turn the precipitate black (366). 360. The carbonates precipitate basic bismuth carbonate, Bi202C03, white, insoluble in excess. Barium carbonate forms the same precipitate, without heating. 361. Hydrosulphuric acid and sulphides precipitate bismuth sul- phide, Bi2S3, black, insoluble in dilute acids and in alkali hydroxides; insoluble in alkali sulphides (distinction from arsenic, tin, antimony), and in alkali cyanides (distinction from copper). It is soluble by moderately concentrated nitric acid (distinction from mercury), the sulphur mostly re- maining free. 362. Soluble chromates—both K2Cr04 and K2Cr207—precipitate the yellow, basic bismuth chromate, Bi20(Cr04)2, distinguished from that of lead by its insolubility in fixed alkali hydroxide. 363. Phosphoric acid and soluble phosphates precipitate bismuth phos- phate, BiP04, insoluble in five per cent, nitric acid, insoluble in dilute acetic acid, readily soluble in hydrochloric and sulphuric acids.—Arsenic acid and arseniates form a precipitate corresponding to the phosphate in composition, and having the same solubilities. 364. Oxalic acid and oxalates precipitate bismuth oxalate, Bi2(C204)3, white, insoluble in dilute acids. Potassium lodide produces in slightly acidulated solutions of bismuth .salts—not acidulated to excess with hydrochloric acid—a dark brown pre- cipitate of bismuth iodide, partly basic, soluble in excess of the reagent, in hydrochloric acid and in hydriodic acid—in each case with a brown tinge to the solution, not soluble in dilute nitric acid.* Bromides precipitate a basic salt, soluble in acid. * This precipitate, at the moment of its formation in concentrated solutions, is doubtless normal bis- muth iodide, Bil3, which is gradually decomposed by water, more rapidly in dilute solutions, forming basic iodide (oxy-iodide) with separation of hydriodic acid. The oxy-iodide of the composition BiOI is stated to be insoluble insolations of alkali iodides, while this precipitate is soluble in these solutions, even after de- ■composition by much water. The reaction of iodides, with bismuth solutions, differs in degree but not in kind from that of chlorides; the normal bismuth iodide only requiring stronger acidulation to hold it in solution than the normal chlor- ide Also, intermediate between the behavior of these two lies that of bismuth bromide. The aqueous ■iodides form a very delicate test for even quite strongly acidulated solutions of bismuth salts, and the bis- muth iodide may not improperly be classed as an “ insoluble ” salt (857). Bismuth. 93 365. Alkaline cyanides precipitate the white hydroxide Bi(OH)s, with formation of hydrocyanic acid. The precipitate is insoluble in the reagent. —Ferroeyanides form a white to yellow precipitate; ferricyanides a yel- low to brownish-yellow precipitate—botli normal bismuth salts, and both insoluble in acids. Tannic acid throws down bismuth tannate, yellow. 366. Metallic bismuth is reduced from bismuth solutions, mostly as a spongy precipitate, by zinc, iron, tin, lead, copper, and cadmium. Potassium or sodium stannite (K2Sno2), when added in excess to bis- muth solutions, causes a black precipitate, from reduction to dibismuth dioxide, Bi202, a very delicate reaction. The stannite is made, when wanted, by adding to stannous chloride solution, in a test-tube, enough sodium or potassium hydroxide to redissolve the precipitate at first formed. The basic bismuth nitrate is reduced by grape sugar, in a warm solu- tion of fixed alkali carbonate with formation of a blackish-brown liquid and dark-gray sediment containing dibismuth dioxide. Also, the recent bismuth hydroxide, in suspension with the excess of fixed alkali, is reduced by digestion with grape sugar or milk sugar to a black precipitate. 367. On charcoal, with sodium carbonate, before the blow-pipe, bis- muth is readily reduced from all its compounds. The globule is easily fusi- ble, brittle (distinction from lead), and gradually oxidizible under the flame, forming an incrustation (Bi2Os), orange-yellow while hot, lemon-yel- low when cold, the edges bluish-white when cold. The incrustation disap- pears, or is driven by the reducing flame, without giving color to the outer flame. With borax or microcosmic salt, bismuth gives beads, faintly yellowish when hot, colorless when cold. 368. Estimation.—(1) As metallic bismuth formed by fusion with po- tassium cyanide. (3) As Bi2Os formed by ignition of bismuth salts of or- ganic acids, or of the salts of volatile inorganic oxyacids. (3) By precipi- tation by H2S, and after drying at 100° C., weighing as Bi2S3. (4) By pre- cipitation by K2Cr20T, and after drying at 130° C., weighing as BIsO- (0rO4),, 369. Oxidation.—All compounds of bismuth having less than five bonds are oxidized to Biv by Cl, in presence of KOH, The triad is reduced to the dyad in presence of KOH by K2Sn02, and by grape sugar. Metal- lic bismuth reduces salts of Hg, Ag, Pt, and An to the metallic state. From bismuth salts the free metal is precipitated by Zn, Mg, Al, Cd, Pb, Fe, and Cu. 94 Cadmium. CADMIUM. Cd = 111.835. 370. Specific gravity, hammered, 8.667 (Schroder, 1859). Vapor density (H = 1), 55.8 (Deville and Troost, 18(51). Melting point, 830° C. (608° F.) (Riemsdyk, 1869). Boiling point, 768° to 772° C. (Carnelley and Williams, 1878). 371. Occurrence.—Found in greenoeldte (CdS), and with zinc in many of its ores. Preparation.—Reduced by carbon and separated from zinc (approximately) by distilla- tion, the cadmium being more volatile. It may be reduced by fusion with H, CO, or coal gas. 372. Properties.—A white crystalline metal, soft, but harder than tin or zinc; more tenacious than tin; malleable and very ductile, can easily be rolled out into foil or drawn into fine wire, but at 80° C. it is brittle. It may be completely distilled in a current of hydrogen above 800° C.; only slightly tarnished by air and water at ordinary tempera- tures. When ignited burns to CdO. When heated combines directly with Cl, Br, I, S, Se, P, and Te, It forms many useful alloys. Oxide and Hydroxide.—lts only hydroxide, Cd(OH)2, is formed from its salts by precipitation with KOH or NaOH; and its only oxide (CdO), by ignition of Cd(OH)2. Cd(NO3)2, Cd 0204, CdCO3, etc. 373. Solubilities.—lt dissolves slowly in hot, moderately dilute hydrochloric or sul- phuric acid, with evolution of hydrogen; in nitric acid, more readily with generation of nitrogen oxides.—Cadmium forms a single oxide, Cd"o, yellowish-brown, and a corre- sponding series of salts, from which it is reducible, in the wet way, only by strong reduc- ing agents. It forms numerous double salts, especially haloids.—The hydroxide, sul- phide, carbonate, oxalate, phosphate, cyanide, ferrocyanide, and ferricyanide are insolu- ble in water. The chloride and bromide are deliquescent, and soluble in alcohol as well as water; the iodide is permanent, and soluble in water and alcohol; very sparingly in ether. The ammonio-oxide and the potassio and sodio cyanides are soluble in water. All of its salts that are insoluble in water are soluble in hydrochloric and nitric acids and in ammonium hydroxide, except CdS. 374. Reactions of Cadmium Salts.—Fixed alkalies precipitate from solutions of cadmium salts—in absence of tartaric and citric acids, and cer- tain other organic substances—the white hydroxide, Cd(OH)2, insoluble in excess of the reagents (distinction from zinc). Ammonium hydroxide forms the same precipitate, which it redissolves. Alkali carbonates preci- pitate CdCC3, white, insoluble in excess of the reagents. Barium carbon- ate forms a complete precipitate, in the cold.—Hydro sulphuric acid and sulphides throw down the sulphide, CdS, yellow ; insoluble in cold dilute acids, in alkalies, and in alkali sulphides and cyanides, soluble in hot and dilute sulphuric acid (compare 375).—Alkali chromates precipitate yellow cadmium chromate, from concentrated solutions only, and soluble on addi- tion of water.—Phosphates form a white precipitate, readily soluble in acids ; oxalates and oxalic acid, cadmium oxalate, white, difficultly soluble in acids. Potassic cyanide precipitates Cd(CH)2, white, soluble in excess of the reagent, as K2Cd(CH)4; ferrocyanides form a white ; ferricyanides, a yellow precipitate—both soluble in hydrochloric acid, and in ammonium hydroxide. On charcoal, with sodium carbonate, cadmium is reduced before the Cadmium. 95 blow-pipe to metallic salt, and usually vaporized and reoxidized nearly as fast as reduced, thereby forming a characteristic brown incrustation (CdO). This is volatile by reduction only, being driven with the reducing flame.— Cadmium oxide colors the borax bead yellowish while hot, colorless when cold ; microcosmic salt, the same. 375. Cadmium may be separated from copper: (1) By the solubility of CuS in KCjM, CdS being insoluble in KCN ; or, better, by treating the blue ammomacal solution with KCN until the blue color disappears, then H2S will precipitate the cadmium as CdS, while the copper remains in solution. (3) By reduction of the copper to Cu2Cl2 with SnCl2 and its precipitation with milk of sulphur, removal of excess of tin with NH4OH and precipita- tion of the cadmium with H2S. (3) The solution is acidified with HCI and the copper precipitated with hot Na2S2C3, and the cadmium in the filtrate is precipitated by (M4)2S after neutralization with 35TH4OH. (4) From the mixed sulphides, CdS is dissolved by dilute hot H2SC4 (one part H2S04 to five of H2O). 376. Comparison of Certain Reactions of Bismuth, Copper, and Cad- mium. Taken in Solutions of their Chlorides, Nitrates, Sulphates, or Acetates. Bi. Cu. Cd. KOH or NaOH, in excess. Bi(OH)s, white. Cu(OH)2, dark blue. Cd(OH)2, white. NH4OH, in excess. Bi(OH)3, white. Blue solution. Colorless solution. Dilution of saturated solu- tions. BiOCl, etc., white. Iodides. Partial precip., in solutions not very strongly acid (304). Partial precipitate completed by re- ducing agents (345). No pre. Sulphides. Bi2S3, black, insol. in cyanide. CuS, black, sol. in cyanide (375). CdS, yellow, insol- uble in cyanide. Iron or Zinc. B' (spongy precip.) (378). Cu (bright coating) (351). Cd (gray sponge). Sugar, KOH and heat. BiO and Bi (black) (366). Cu2(OH)2 (yellow). K2SnQ2 + KOH. Bi202 (black). Metallic copper. Metallic cadmium. 377. Estimation.—(l) It is converted into, and after ignition weighed as, an oxide ; (2) converted into, and after drying at 100° -C., weighed as CdS ; (3) precipitated as Lead. CdC204 and titrated by K2Mn2OB; (4) precipitated as CdS and reduced with Fe2Cl6, and the amount of reduction determined by K2Mn2Os. 378. Oxidation.—Metallic cadmium precipitates the free metals from solutions of Au, Pt, Ag, Hg, Bi, Cu, Pb, Sn, and Co; and is itself reduced by Zn, Mg, and Al. LEAD. Pb = 206.471. 379. Specific gravity, 11,38 (Reich). Melting point, 326.2° C. (619.3° F) (Person, 1849). Valence, a dyad in Pb O, and in all lead salts ; a tetrad, in Pbo02 and in plum- bates. 380. Occurrence.—Lead is seldom found in the free state in nature. Its chief ore is galena (PbS). In smaller quantities it is found as cerussite, or white lead ore (PbCO3) ;as anglesite (PbSO4), and in very many other minerals. 381. Preparation.—Prom galena (1) It is roasted in the air, forming variable quan- tities of PbSO4, PbO, and PbS ; then the air is excluded and the temperature raised, and the sulphur of the sulphide reduces both the PbO and the PbSO4, SO2 being formed : PbS04 4- PbS = 2Pb + 2SO, 2PbO + PbS = 3Pb + SO, (3) Similar to the first except that some form of carbon is used to aid in the reduc- tion. (3) It is reduced by fusing with metallic iron (PbS -p Fe = Pb + FeS). Fre- quently these methods are combined or varied according to the other ingredients of the ore. 382. Properties —A bluish-white, soft metal ;it can be rolled out into sheets, but not drawn into wire ; nearly inelastic ; is a poor conductor of heat and electricity ; it forms alloys with most metals. Solder is lead one part, tin one part; type metal, lead two parts, tin and antimony each one part. Shot contains 0.5 per cent, of arsenic. Lead is slowly volatile at a white heat. It tarnishes in the air at ordinary temperatures by formation of diplurnbic monoxide, Pb2o, blackish gray. Pure water, free from air, does not affect lead, free from oxide or hydroxide, in the cold ; but granulated lead slowly decomposes boiling water, with evolution of hydrogen, and formation of lead hydroxide, Pb(OH)2. In water containing air, the hydroxide and basic carbonate are formed. This corrosion and solution are greatly promoted by nitrogenous organic mat- ters—ammonium salts, and nitrates and nitrites—and by chlorides ; hindered or pre- vented by carbonates, acid carbonates and sulphates. Above the melting point, lead gradually oxidizes in the air to “ litharge,” PbO. 383. Oxides.—Lead forms four oxides, Pb2o, PbO, Pb02l and Pb304. Lead sub- oxide (Pb2o) is little known ;it is the black powder formed when Pbo204 is heated to 300° C., air being excluded. Lead oxide (litharge, or massicot) is formed by intensely igniting in the air Pb, Pb2o, Pbo2. Pb304, Pb(OH)2, PbCO3, PbC204, or Pb(NO3)2. It has a yellowish-white color, melts at a red heat, and is volatile at a white heat. 384. Triplumbic tetroxide (red lead or minium), Pb304, is formed by heating PbO to a dull-red heat with full access of air for several hours. Strong, non-reducing acids, such as HNO3, H2S04, HClO3j etc., convert it into a lead salt and Pb02 (a). But con- centrated hot H2SO4 converts the whole into PbS04, oxygen being evolved (b). But with the dilute acid and reducing agents, such as glycerine, sugar, H2C204, H2C4H406, Zn, Al, Cd, Mg, As, Pb, etc., it is all reduced to the dyad lead without evolution of oxy- gen (c), (d). and (e). Hydracids usually reduce the lead and arc themselves oxi- dized (/). Lead. 97 (a) Pb304 + 2H2SO, (dilute) = Pb02 + 2PbS04 + 2H20 (b) 2Pb304 + 6H2504 (concentrated and hot) GPbSC>4 + GH2O +O2 (c) Pb304 + H2C204 + 6HNO3 = 3(Pb(NO3)2 + 4H20 + 2C02 (d) 10Pb,04 + As4 + 30H2SO4 = 80PbSO4 + 4H3AsO4 + 24H.0 (e) Pb304 +Zn + 4H.,504 = 3PbS04 + ZnS04 + 4H20 (/) Pb304 + BHCI = 3PbCI2 + Cl 2 + 4H20 The valence of Pb304 is best explained by the theory that it is a union of the dyad and tetrad (Pb" and Pbiv), Pb304 = 2Pb"O + PbivQ2). 385. Lead dioxide or peroxide (PbO-2), is formed (1) by fusion of PbO with KOlOs or KNOs; (3) by fusing Pb3o4 with KOH; (3) by treating any compound of Pb'' with Cl, Br, K6Fe2(CN)12, K2Mn2Os, or H202 in presence of KOH; (4) by treating Pb304 with non-reducing acids. Ignition forms first Pb3o4 and above a red heat PbO, oxygen being given off. It dis- solves in acids on same conditions as Pb304 (see 384). Very strong solution of potassium hydroxide, in large excess, dissolves it, with formation of “ potassium plumbate,” K2Pb03. Lead dioxide is a powerful oxidizing agent, one of the strongest known. Di- gested with ammonium hydroxide, it forms lead nitrate and water. Triturated with one-sixth of sulphur, or tartaric acid, or sugar, it takes fire ; with phosphorus, it deton- ates. 386. Solubilities.—Dilute nitric acid is the proper salifying solvent for metallic lead, forming plumbic nitrate with evolution of nitric oxide. Concentrated nitric acid acts more slowly. Lead does not dissolve in dilute sulphuric acid, cold or hot, or in con- centrated sulphuric or hydrochloric acid, in the cold; but hot sulphuric acid, containing less than twenty-five per cent, water, forms lead sulphate, sparingly soluble in the con- centrated acid; and hot concentrated hydrochloric acid forms, with evolution of hydro- gen, and dissolves, a limited proportion of lead chloride. Dilute hydrochloric acid forms chloride, but dissolves little of it. The oxide, and hydroxide (formed in water, 883), are soluble in 7,000 to 10,000 parts of water, to which they give the alkaline reaction. The sulphide, carbonate, phosphate, chromate, sulphite, borate, cyanide, ferrocyanide, and tannate are insoluble in water. The sulphate and oxalate are very slightly soluble in water ; the chloride, iodide, bro- mide, and ferricyanide are sparingly soluble in hot water, still more sparingly soluble in cold water. The sulphate and chloride are less soluble in dilute sulphuric and hydro- chloric acids than in pure water, but much more soluble in the same acids concentrated than in water. Nitric acid increases the solubility of the sulphate and chloride in water, more and more, as the nitric acid is stronger—the salts separating again on diluting the nitric acid solution. The sulphate and chloride are insoluble in alcohol. The iodide is moderately soluble in solutions of alkaline iodides, insoluble in alcohol, decomposed by ether. The basic acetates are permanently soluble (if carbonic acid is strictly excluded). The basic nitrates are but slightly soluble in water, and are precipitated on adding solu- tions of potassium nitrate to solution of basic lead acetate. In analysis, the solubility of the chloride, sparing as it is, enables lead to be sepa- rated from the other first-group metals. As a final precipitate, in both first and second groups, the sulphate is most used. The sulphide precipitate exceeds other tests in delicacy. Pb304 + 4HN03 = PbOj + 2Pb(NOs)2 + 2H20 387. Fixed alkalies precipitate, from solutions of lead salts, lead hy- droxides, Pb(OH)2, white, soluble in excess of the reagents, by combination, as potassium or sodium plumbite, K2Pb02 (distinction from silver, mercury. 98 Lead. bismuth, copper, cadmium). All the precipitates of lead hereafter given, except the sulphide and ferricyanide, are soluble in strong solutions of the fixed alkali hydroxides. The alkaline solution of lead is precipitated by alkaline solutions of chromic, stannic, stannous, antimonious, and arsenious oxides. Ammonium hydroxide precipitates white basic salts, insoluble in ex- cess (distinction from silver, copper, cadmium) : with the chloride, the precipitate is Pb2OCl2 ; with the nitrate, Pb3020HN03. With the acetate, in solutions of ordinary strength, excess of ammonium hydroxide (free from carbonate) gives no precipitate, soluble tribasic acetate being formed. Soluble carbonates precipitate lead basic carbonate, 'white, the carbon- ate and hydroxide combined in proportions varied by conditions. With excess of the reagent, in concentrated solution, the precipitate consists chiefly of Pb3(OH)2(C(D3)2. Barium carbonate on boiling precipitates lead salts completely. Free carbonic anhydride precipitates the basic acetate. 388. Hydrosulphuric acid and the sulphides precipitate—from neu- tral, acid, or alkaline solutions—lead sulphide, PbS, brownish-black, insol- uble in highly dilute acids, in alkalies, or alkali sulphides. Freshly pre- cipitated CdS, MnS, FeS, CoS, and HiS, give the same precipitate. H2S changes all freshly precipitated lead salts -'o PbS, Moderately dilute (15 to 25 per cent.) nitric acid dissolves lead sulphide, with separation of sulphur (equation a) ; concentrated nitric acid changes it mostly to the (insoluble) lead sulphate (equation b)—in both cases with evolution of nitric oxide. The oxidation of the sulphur always occurs in the action of nitric acid on sulphides, in degree proportioned to the strength of acid, temperature, and duration of contact ; a. fiPbS + 16HN03 = 6Pb(NOa)2 + 352 + 4NO + BH2O b. 3PbS + BHNO3 = BPbSO4 + BNO + 4H20 In solutions too strongly acidulated, especially with hydrochloric acid, the formation of brick-red basic sulphides, as Pb2SCl2, interferes with per- fect precipitation ; in solutions excessively dilute, only a brown coloration occurs without precipitation. Lead is revealed in solutions in 100,000 parts of water, by this test. 389. Sulphuric acid and sulphates precipitate, from neutral or acid solutions, lead sulphate, PbS04, white, not chemically changed or perma- nently dissolved by acids, except hydrosulphuric acid,yet slightly soluble in strong acids, as more particularly stated in 386. Soluble in boiling am- monium acetate, and in the fixed alkalies. For solution by transposition into soluble salts, see 396. Soluble in warm sodium thiosulphate solution, at temperatures not above 68° C. (154° F.) ; (in hot solution, decomposed, PbS being one of the products formed, insoluble in thiosulphate); dis- tinction and separation from barium sulphate,'which does not dissolve in thiosulphates. Lead. 99 This test is from five to ten times less delicate than that with hydrosul- phuric acid ; but lead is quantitatively separated as a sulphate, by precipi- tating with sulphuric acid in presence of alcohol, and washing with alco- hol. If the PbS04 is heated with K2orO4, transposition takes place, and the yellow PbCr04 is formed (393). The yellow precipitate is soluble in fixed alkali hydroxides, then reproduced by acetic acid. Also, excess of po- tassium iodide transposes lead sulphate, the yellow product (392) being a distinction of lead from barium. 390. Hydrochloric acid and soluble chlorides precipitate, from solu- tions not too dilute, lead chloride, PbCl2, white. This reaction constitutes lead a member of the FIRST GROUP—as it also is of the second. The solubility of the precipitate is such (386) that the filtrate obtained in the cold gives marked reactions with hydrosulphuric acid, sulphuric acid, chro- mates, etc.; and that it can be quite accurately separated from silver chlor- ide and mercurous chloride by much hot water. Also, small proportions of lead escape detection in the first group, while its removal is necessarily ac- complished in the second group. 391. Soluble Bromides precipitate lead bromide, Pbßr2, white, soluble in water to about the same extent as the chloride ; in concentrated solu- tions, the precipitate dissolves in excess of the potassium bromide, as (KBr)2Pbßr2, which is decomposed and precipitated by dilution with water. Also soluble in hot solutions of ammonium chloride and nitrate. 392. Soluble lodides precipitate lead iodide, Pbl2, bright yellow and crystalline, soluble in about 1,900 parts of cold or 200 of hot water ; soluble in hot moderately concentrated niiric acid, and in solutions of fixed alkalies not in cold hydrochloric acid ; soluble in excess of the alkali iodides, by formation of double iodides—with deficient excess of potassium iodide, forming KIPbI2 ; with superabundance of the same reagent, forming (KI)4PbI2, these double iodides requiring free alkali iodide to hold them in solution, and being partly decomposed by undue addition of water, with reprecipitation of the lead iodide. Lead iodide is not precipitated in pres- ence of sodium citrate ; alkaline acetates also hold it in solution to some extent, so that it is less perfectly precipitated from acetate than from nitrate of lead. 393. Soluble Chromates—both K2Cr04 and K2Cr2C7—precipitate lead chromate, PbCrC4, yellow, soluble in fixed alkali hydrates (distinction from bismuth), insoluble in chromic acid (distinction from barium), slightly soluble in acetic acid, decomposed by hydrochloric acid and by ammonium hydrate. 394. Disodium hydrogen phosphate precipitates trimetallie lead phosphate, Pb3- 1P04)2, white, insoluble in dilute acetic acid (compare 216), soluble in nitric acid and fixed alkalies: 3Pb(NOs)j + 3Na2HP04 = Pb3(PO4)2 + 6NaNO3 + H3P04 And H3P04 -f Na2HP04 = 3NaH2PO4 100 Lea d. Therefore, if there is excess of phosphate, the full reaction will bes Alkali oxalates precipitate Lead oxalate, PbC204, white, insoluble in acetic acid, solu- ble in potassium and sodium hydroxide solutions, and in nitric acid. Alkali sulphites—as Na2SO3—precipitate lead sulphite, PbSC3, white, less soluble in water than the sulphate, slightly soluble in sulphurous acid, decomposed by sulphuric, nitric, and hydrochloric acids. 395. Soluble cyanides—as KCN—precipitate lead cyanide, Pb(CN)2, white, soluble in a very large excess of the reagent, reprecipitated on boiling.—Ferrocyanides—as K4Fe(CN)6—precipitate ferrocyauide, Pb2Fe(CN)6, insoluble in dilute acids.—Ferri- cyanides form, in concentrated solutions, a dark brown precipitate, slightly soluble in water.—Sulphocyanates form, in concentrated solutions, a yellow crystalline precipitate of lead sulphocyanate, Pb(CNS)2, soluble in water, decomposed on boiling, with precipi- tation of basic sulphocyanate, PbOH(CNS), white. 396. Tannic acid precipitates solutions of lead acetate, and partly the nitrate, as yel- low-gray tannate of lead, soluble in acids. Solution of lead acetate precipitates a large number—and solution of lead subacetate a still larger number —of organic acids, color substances, resins, gums, and neutral principles. Indeed, it is a rule, with few excep- tions, that lead subacetate removes all organic acids (not acetic, formic, butyric, valeric, or lactic). Ammoniacal solution of lead acetate is used as a reagent, as a form of basic acetate (387). 397. Lead salts when fused on porcelain with Na2C03 are converted into PbO (a). If charcoal is added metallic lead is formed {b). Long continued fusion on charcoal may change the acid radical also (c). 3Pb(NO3)2 + 4Na2HPO4 = Pb3(P04)2 + ONaNOs + 2NaH2PO4 (a) PbCl2 + Na.COs = SNaOl + PbO + CO2 (b) 2PbS04 + 2Na2COs + O = 2Pb + 2Na2S04 + 8C02 (c) 2PbS04 + 2Na2C03 +5O = 2Pb -f 2Na2S + 7C02 After fusion the aqueous solution is tested for acids and the residue for bases after dis- solving in HN03 or hc2h302. 398. With borax and microcosmic salt, strictly in the older flame, lead oxide and oxi- dized compounds give a bead yellow when hot, becoming colorless when cold ; due to formation of lead borate or phosphate, fused in the glass. If the least reducing action is allowed to bear on the bead, the test is spoiled, and the platinum wire is spoiled likewise. (See under Platinum.) 399. Estimation.—(l) As an oxide into which it is converted by ignition (if a carbon- ate or nitrate), or by precipitation and subsequent ignition. (2) As a sulphate. Add to the solution twice its volume of alcohol, precipitate with H2S04, and after washing with alcohol ignite and weigh. (3) It is converted into an acetate, or sodium acetate is added to the solution, then precipitated with K2or207, and after drying at 100° C. weighed as PbCro4. (4) It is converted into PbS, free sulphur added, and after ignition in hydro- gen gas weighed as PbS. 400. Oxidation.—Pb" is oxidized to Pbiv, as stated in (385). Pbiv is reduced to Pb" in presence of dilute H2S04 by nascent hydrogen, and by all metals capable of pro- ducing nascent hydrogen (such as Al, Zn, Sn, Mg, Fe), and by soluble compounds of (Hg2)", Sn", Sb'", As"', (AsH3 gas), Bi”, (Cus)", Fe”, (Cr2)vi, Mn ", Mu'", Mmv, Mnvi. Also by H20204, HN02, H3PO2, H3P03, P, SO2, H2S, HCI, HBr, HI, HOy, HCNS, H4Fe(CN)6, glycerine, tartaric acid, sugar, urea, and very many organic com- pounds. In many cases the same reduction takes place in presence of KOH. Prom lead solutions Zn, Mg, Al, Co, and Cd precipitate metallic lead. And metallic lead precipitates the free metals from solutions of Au, Pt, Ag, Hg, Bi, and Cu. SIL VER. 101 SILVER. Ag = 107.675. 401. Specific gravity, precipitated, 10.5583 (G. Rose, 1848). Melting point, 954° C. (1749° F.)(Violle, 1877). Vaporizes at 1570° C. (V, and 0. Meyer, 1879). Valence, a monad in Ag20 and in silver salts. 402. Occurrence.—Found in a free state; oftener in combination ; its most impor- tant ores are argentite, or silver glance (Ag2S), pyrargyrite (Ag3SbS3), and horn silver (AgCl), and it is frequently found in paying quantities in galena (PbS), and copper py- rites, and in many other ores. 403. Preparation.—The limits of this work do not permit a description of the metal- lurgy of silver. Chiefly three methods are employed: (1) It is alloyed with lead by fusion and the lead separated by oxidation. (3) It is amalgamated with mercury and then the mercury separated by distillation. (8) It is brought into solution and the metal precipi- tated by copper. Silver is very easily reduced from its oxide by heat alone, and from all its compounds by ignition with H, C, CO, and by the organic carbon compound in which the carbon has less than four bonds. It is also reduced without the aid of heat by cer- tain metals, etc. (343). 404. Properties.—Silver is the whitest of metals, harder than gold and softer than copper; silver is hardened by copper. United States silver coin contains 90 per cent, silver and 10 per cent, copper. In malleability and ductility it is inferior only to gold ; and as a conductor of heat and electricity it exceeds all other metals. 405. Oxides.—Silver forms three oxides—AgoO, Ag4C, Ag2C2. Silver oxide, ar- gentic oxide (Ag20) is formed by the action of KOH or NaOH on silver salts or by heat- ing the carbonate to 300° C. It is a brown powder, a strong oxidizing agent, partially decomposed by light, completely decomposed by heat at 300° C. into metallic silver and oxygen. Argentous oxide (called also suboxide, quadrantoxide, and tetrantoxide), Ag4C, is formed by heating silver citrate to 100° C. in hydrogen gas, dissolving in water and precipitating with KOH. It is a black powder, easily decomposed by heat, soluble in NH.OH; decomposed by oxyacids forming metallic silver and a silver salt; with HCI forming argentous chloride (Ag2CI). Silver peroxide (Ag202) is a black powder formed by treating metallic silver or silver oxide with ozone or peroxide of hydrogen. Oxyacids reduce it, forming a silver salt and evolving oxygen. HCI reduces it to AgCl, evolving free Cl. Silver forms no definite hydroxides. 406. Solubilities.—Silver is not oxidized by water or air at any temperature, but is oxidized by ozone, is readily attacked by chlorine, bromine, or iodine, and is soon tar- nished in air containing hydrosulphuric acid, or in contact with sulphides or certain or- ganic substances containing sulphur, by formation of silver sulphide; also, by substances easily liberating phosphorus, as silver phosphide. -As silver is easily reduced from its salts, these act as oxidizing agents of considerable force. 407. The proper solvent of silver is nitric acid, most efficient when about fifty per cent., but active whether concentrated or dilute—with production of nitric oxide as the chief residual product. Hot concentrated sulphuric acid forms sulphate, which is spar- ingly soluble ; and hot concentrated hydrochloric acid forms silver chloride, slightly soluble in the concenti’ated reagent, but precipitated on dilution. The fixed alkalies do not act upon silver in the wet or dry way ; hence, silver crucibles are used instead of platinum for fusion with caustic alkali. Silver, in the form of a precipitate, is very slowly acted upon by strong aqueous ammonia, dissolving as a nitride.—There is but a single series of salts of silver—those represented by Ag', and sometimes designated argentic salts. 408. The nitrate, acetate, and sulphate form permanent anhydrous crystals. The salts of silver are chiefly colorless, except the ortho-phosphate and arsenite, yellow; the tSIL VER. arseniate, reddish-brown; the iodide, yellow; the bromide, yellow-white; the sulphide, black. Normal silver salts do not redden litmus. 409. Silver forms a greater number of insoluble salts than any other metal; though, in this respect, there is but little difference between the first-group bases. The oxide, sulphide, chloride, bromide, brornate, iodide, iodate, cyanide, ferrocyanide, ferricyanide, carbonate, oxalate, phosphate, arsenite, arseniate, sulphite, and tartrate are insoluble in water; the sulphate is soluble in 200 parts of cold, and less than 100 parts of boiling water. The acetate is soluble in 100 parts of water. The borate, thiosulphate, and citrate are very sparingly soluble in water. The ammonium silver oxide and the nume- rous ammonium silver salts, the double cyanides, iodides, and thiosulphates of silver and alkaline metals, are soluble in water. The chloride is sparingly soluble in strong hydro- chloric, nearly insoluble in nitric and dilute sulphuric acids ; soluble, to some slight extent, in solutions of all soluble metallic chlorides (except calcium and zinc chlorides), especially soluble with sodium chloride (double chloride being formed); also soluble with certain other alkali salts, and in concentrated solution of mercuric nitrate. The nitrate is sparingly soluble in alcohol and in ether, and soluble in glycerine. 410. Both the oxy-salts and haloid salts of silver, which are insoluble in water, are dissolved by ammonium hydroxide, except the sulphide and iodide; by cold dilute nitric acid, except the chloride, bromide, iodide, bromate, iodate, and the haloids of cyanogen and its compounds ; by solution of potassium cyanide, except the sulphide ; and by al- kali thiosulphates, almost without exception. In analysis, silver is completely precipitated as a chloride, in the first group (418), and the solubility of this precipitate in ammonium hydroxide separates it from the other first-group bases. Reduction to metallic silver is sometimes employed in analysis (443). 411, The fixed alkali hydroxides precipitate, from solutions of silver salts (in absence of citrates), silver oxide, Ag„o, grayish-brown (329), insol- uble in excess of the reagents ; easily soluble in nitric, acetic, and sul- pha ric acids, and in ammonium hydroxide ; somewhat soluble in ammoni- um salts ; soluble in alkali cyanides and thiosulphates; also, soluble in about 3,000 parts of water. Ammonium hydroxide, in neutral solutions of silver nitrate, forms the same precipitate, silver oxide, very easily dissolving in excess, by formation of ammonium silver oxide, NH4AgO.* In solutions containing much free acid, all precipitation is prevented bydhe ammonium salt formed. The ammoniacal solution of silver is not immediately precipitated by addition of excess of fixed alkalies in the cold, but on boiling a black precipi- tate is formed. From the cold ammoniacal solution containing fixed alka- lies a precipitate gradually forms. This precipitate contains fulminating silver—a black powder, which explodes with dangerous violence by friction or by drying above ordinary temperatures. Fulminating silver may also be deposited from ammoniacal solutions of silver, on standing, and by digest- * This formula accords with the results of a series of volumetric determinations made by Mr. D. E. Os- borne and the author (Jour. Am. Chem. Soc., 1880). If silver replaces hydrogen of ammonium, the formula would be (NH3Ag-)2o—the molecule of which, with a molecule of water, would make two molecules formed as given in the text. For the latter, we have : AgN03 + 2NH4OH = NH4AgO + NH4N03 + H^O Sjl ver. 103 ing oxide of silver with strong ammonium hydroxide. Its production, in the way first mentioned, is most favored by a slight excess of the fixed alkali.* ; 412. Hydrosulphuric acid and alkali sulphides precipitate from neu- tral, acid, or alkaline solutions, silver sulphide, Ag2S, black, soluble in moderately concentrated nitric acid (distinction from mercury), not in so- lution of potassium cyanide (distinction from copper) ; insoluble in alkali sulphides (distinction from tin, etc.) 413. Hydrochloric acid and the soluble chlorides precipitate silver chloride, AgCl, white, curdy, separating on shaking the solution ; turning violet to brown in the light (from formation of argentous chloride, Ag2CI), very easily soluble in ammonium hydroxide, as ammonio silver chloride, The precipitate, also, is slowly soluble in concentrated solution of ammonium carbonate; and is fusible without decomposition. For solubilities of the precipitate—indicating the conditions of delicacy in the test—see 409 and 410. This precipitation is the most delicate of the ordinary tests for silver; being recognized in solution in 250,000 parts of water, and enables us wholly to remove this metal IH THE FIRST GROUP of bases. 414. Soluble bromides precipitate silver bromide, Agßr, white, with a slight yellowish tint, but slightly soluble in excess of potassium bromide, and much less easily soluble in ammonium hydroxide than silver chloride. Soluble iodides precipitate silver iodide, Agl, pale yellow, easily soluble in excess of the reagents by formation of double iodides, as KIAgl. The double iodide is decomposed by dilution with much water, and all the silver reprecipitated as iodide. The precipitate is scarcely at all soluble in am- monium hydroxide (one part dissolving in 2,600 parts of ten per cent, solu- tion of ammonia). Concentrated nitric acid slowly dissolves it. Regard- ing other solubilities of argentic bromide and iodide, see 409 and 410. 415. Potassium cyanide, or hydrocyanic acid, precipitates, from neu- tral or slightly acid solutions, silver cyanide, AgCH, white, quickly soluble in excess of the reagent, as potassium silver cyanide, KAg(CH)2, By for- mation of these double cyanides, the various compounds of silver are ren- dered soluble through treatment with alkali cyanides; also, a soluble iodo- cyanide is formed. Silver cyanide is readily soluble in ammonium hydrox- ide, and promptly decomposed by hydrochloric acid. 416. Potassium ferrocyanide precipitates silver ferrocyanide, Ag4Fe- (cn)6. yellowish-white, difficultly soluble in ammonium hydroxide; on boil- ing metallic silver separates and a ferricyanide is formed, not decomposed by hydrochloric acid, changed by nitric acid to the ferricyanide. Exposure * The composition of this substance, known as Berthollet’s Fulminating Silver, has not been determined, but it contains nitrogen. It is distinct from the silver fulminate, A-g? CN , represented by fulminicacid, aud isomeric with cyanates. 104 SIL VER. to the air gives it a blue tinge.—Potassium ferricyanide precipitates silver ferricyauide, reddish-yellow, soluble in ammonium hydroxide. 417. Alkali carbonates precipitate silver carbonate, Ag2C03, white or yellowish- white, slightly soluble in water, somewhat soluble in excess of fixed alkali carbonates, quite soluble in ammonium carbonate; soluble in nitric acid and in ammonium hydrox- ide; changed by boiling to silver oxide. Barium carbonate does not affect solution of silver nitrate. 418. Oxalic acid, and oxalates, precipitate silver oxalate, Ag2C204, white, slightly soluble in water, sparingly soluble by dilute nitric acid, readily soluble in solution of ammonium hydrate. It detonates when heated. 419. Disodium hydrogen phosphate precipitates trimetallic silver ortho-phosphate, Ag3P04, yellow, soluble in dilute nitric acid, in phosphoric acid, and in ammonium hydrate; but little soluble in dilute acetic acid.—Pyrophosphates—as Na4P207—precipi- tate silver pyrophosphate, Ag4P2C7, white, insoluble in acetic acid, soluble in dilute nitric and phosphoric acids, and in ammonium hydroxide. 420. Arseniates—as Na3Aso4—-precipitate red-brown silve/)' arseniate, Ag3As04, having the same solubilities as the ortho-phosphate.—Arsenites—as Na3Aso3—precipi- tate silver arsenite, Ag3As03. yellow, quickly soluble in dilute acids and in ammonium hydroxide. Chromates—as K2Cr04—precipitate silver chromate, Ag2Cr04, dull-red, sparingly soluble in water, not much more soluble in dilute nitric acid. Thiosulphates—as Na2S203—precipitate silver thiosulphate, Ag2S203, white, very instable, and readily soluble in excess of the precipitants, by formation of double thiosul- phates. That formed by sodium thiosulphate is first NaAgS203, with excess of the thio- sulphate, Na4Ag2(S203)3; and corresponding thiosulphates of silver and potassium are formed. By standing or heating, the precipitate turns black, as Ag2S. AgsSaOs + H.o = AgoS + H2S04 Thiocyanates give AgCNS, insoluble in dilute HNOs. 421. By a gradual reduction of the silver with certain reagents, it is obtained as a bright silver coating upon the inner surface of the test-tube, or other glass vessel. A somewhat dilute solution of ammonio nitrate of silver, treated with a dilute alcoholic so- lution of oils of cloves and cassia—the latter solution not in excess—gives this result. The coating is also obtained by adding to solution of silver nitrate a very little aqueous solution of chloral hydrate, and then a slight excess of ammonia; the ammonium formi- ate, gradually produced by decomposition of the chloral with alkali, deoxidizes the am- monio silver nitrate. A silver deposit on glass may sometimes be made to assume the form of a compact and lustrous coating, by rubbing with a glass rod. In these deoxida- tions, generally, the nitric acid radical of silver nitrate is not decomposed, but nitric acid is left. 4AgN03 + 2H.0 = 4Ag + 4HN03 + 02 See, as an example, the statement of the reaction between arsenious hydride and sil- ver nitrate, under Arsenious Acid. Light decomposes most compounds of silver, with blackening from formation of metallic silver or of argentous oxide, Ag40, or of both. The nitrate in crystal or pure water solution, fhe phosphate, iodide, and cyanide, are not decomposed by light alone ; but light greatly hastens their decomposition by organic substances, or other reducing agents—as of solution of silver nitrate in rain-water, or written as an ink upon organic fabrics. The base of most indelible inks is silver. 422. Silver nitrate and chloride fuse undecomposed, but decompose at a higher heat. Mercury. 105 Most silver compounds, heated in the glass-tube, leave a metallic residue. On Charcoal, with sodium carbonate, silver is reduced from all its compounds in the blow-pipe flame, attested by a bright malleable globule. Lead and zinc, and elements more volatile, may be separated from silver by their gradual vaporization under the blow-pipe. Copper and iron are removed along with larger quantities of lead, previously added for this purpose, either as metallic lead or by reduction from litharge. (See descriptions of Cupellation, in works on general chemistry, and more fully in works on assaying of precious metals). 423. To identify the acid of silver salts which are insoluble in HNOs (AgCl, Agßr, Agl), (1) Add metallic zinc and a drop of H2S04; when the silver is all reduced test for the acid in the filtrate. (2) Fuse with Na2C03, add water, and test the filtrate for acids, (3) Add H2S, and proceed in the same manner. (4) Boil with KOH or NaOH (free from HCI), and test the filtrate in the same manner. It must not be overlooked that by the first three methods, and not by the last, bromates and iodates are reduced to bromides and iodides. 424. Estimation.—(l) As metallic silver, into which it is converted by direct ignition if it is the oxide or carbonate, or by ignition in hydrogen if the chloride, bromide, iodide, or sulphide. (2) It is precipitated as AgCl, and after igniting to incipient fusion (260° C.) weighed. (3) It is converted into Ag2S by H2S, and weighed after drying at 100° C. Inadmissible in case of an acid that might liberate free sulphur. (4) Add KCN until a solution of KAg(CN)2 is formed, precipitate with HNO3, and after drying at 100° C. weigh as AgCN. (5) Volumetrically, by adding a graduated solution of NaCl until a precipitate is no longer formed. This may be varied by adding the measured silver solu- tion to the graduated NaCl solution, containing a few drops of K2Cr04, until the red precipitate begins to form. (6) Volumetrically, add a graduated solution of ammonium thiocyanate, containing ferric sulphate, until the red color ceases to disappear. 425. Oxidation.—Metallic silver precipitates gold and platinum from their solu- tions, and is precipitated as metallic silver from its solutions by Zn, Mg, Al, Cd, Pb, Crq (Cu2)", Sn, Sn", Hg, Bi, Te, Sb, SbH3, As, AsH3, P, H3P02, H2S03, and is reduced in presence of KOH by (Hg2y', Sn'', As'", Sb"', Bi", Mn", and (Cu2)". Also by FeS04 when cold (incompletely), and redissolved on boiling, the ferric sulphate first formed being again reduced to FeSO4. MERCURY. Hg 199.712. 426. Specific gravity, liquid, 13.596 (Volkmann, 1881); solid, 14.193 (Mallet, 1877). Melting (freezing) point, -38.8° C. (-38° F.) Stewart. Boiling point, 357.25° C. (675° F.) (Regnault, 1860). Vapor density (H = 1), 200.9 5 (V. Meyer, 1880). Valence, a dyad mHg O and in mer- curic salts ; a pseudo-monad in (Hg2)"0, and in mercurous salts. 427. Occurrence.—Found native, but its chief ore is cinnabar (HgS). It is also found as Hg2Cl2 and as an amalgam with gold and silver. Preparation.—From HgS. (1) It is roasted in the air ; HgS -f 02 = Hg so„. (2) Lime is added and the mercury distilled ; 4HgS -f- 4CaO = 3CaS _f_ CaS04 + 4Hg. 428. Properties.—lt is the only metal which is a liquid at ordinary temperatures ; white when pure, and having a brilliant silvery lustre. It is slightly volatile at ordinary temperatures. Divided in globules invisible u> the unaided eye, and separated by minute films of liquid or solid foreign 106 Mercury. matter, mercury appears as a dark gray powder. It is not oxidized by agitation with air or oxygen—the tarnish acquired on the surface of com- mercial mercury, by exposure to the atmosphere, being due to intermixture of foreign metals ; but by agitation with water, or with various substances, the metal is “extinguished,” or divided to the gray pulverulent form, which contains some mercurous oxide when so prepared. Also, the gray pulver- ulent mercury is precipitated by reduction from salts in solution. Aqueous solutions of alkali chlorides, with access of the air, gradually act upon mercury by formation of mercuric chloride. Solution of potassium per- manganate oxidizes mercury—forming mercurous oxide, manganic hydrox- ide and potassium hydroxide. 429. Solubilities.—The most effective solvent of mercury is nitric acid. It dissolves readily in the dilute acid hot or cold ; with the strong acid, heat is soon generated ; and with considerable quantities of material, the action acquires an explosive violence. At ordinary temperatures, nitric acid, when applied in excess, produces normal mercuric nitrate, but when the mercury is in excess, mercurous nitrate is formed ; in all cases, chiefly nitric oxide gas is generated. Both mercurous and mercuric nitrates re- quire a little free nitric acid to hold them in solution. This free nitric acid gradually oxidizes mercurosum to mercuricum, making a clear solution of Hg(H03)2, if there is sufficient HNQ3 present, otherwise a basic mercuric nitrate may precipitate. A solution of mercurous nitrate may be kept free from mercuric nitrate by placing some metallic mercury in the bottle containing it; still after standing some weeks a basic mercurous nitrate crystallizes out, which afresh supply of nitric acid will dissolve. Chlorine in aqueous solution, or formed in nitro-hydrochloric acid—-dissolves mercury slowly, to mercuric chloride. Hydrochloric acid does not dissolve mercury. Bromine and iodine promptly unite with mercury. Dilute sulphuric acid does not act upon mercury ; but the concentrated acid, when heated, dis- solves it with moderate rapidity, evolving sulphurous anhydride. 430. Oxides.—Mercurous oxide (Hg„0) is formed by treating mercur- ous oxy-salts or Hg2Cl2 with KOH or HaOH. It is a black powder which a gentle heat changes to HgO and Hg, and a higher heat into Hg and O. Mercuric oxide, HgO, is made (1) by keeping Hg° at its boiling point for a month or longer in a flask filled with air ; (2) by heating Hg2(H03)2 or Hg(H03)2 with about an equal weight of metallic mercury, Hg(N03)3 +3H g = 4HgO + 2NO; (3) by precipitating mercuric salts with KOH or HaOH. Made by (1) and (2) it is red, by (3) yellow. On heating it changes to vermilion red, then black, and on cooling regains its original color. A red heat decomposes it completely into Hg° and o°. Mercury forms no hydroxides. 431. Mercury forms two well-marked classes of salts—mercurous and mercuric ;—mercurous compounds being permanent in the air, but changed Mercury. 107 by powerful oxidizing agents to mercuric compounds. The latter are some- what more stable, but act as oxidizing agents in many relations. Mercury as a noble metal is not strongly electro-positive; and many reducing agents change mercuric compounds, first to mercurous combinations, and then to metallic mercury. Solutions of mercurous and mercuric salts redden litmus. Mercuric chloride is permanent; nitrate, deliquescent. 432. MERCUROUS compounds, of ordinary occurrence, are insoluble in water, except the normal nitrate ; the sulphate and the acetate are spar- ingly soluble (that is, in 300 to 600 parts of water). And these require acidulated water for their full solution ; becoming decomposed by water, at a certain degree of dilution, with precipitation of basic salts (326). Mercurous chloride is very slowly soluble by cold concentrated solutions of alkali chlorides, somewhat more rapidly when healed, the solution being due to formation of mercuric chloride and mercury. Dilute hydrochloric acid, at ordinary temperatures, fails to dissolve mercurous chloride ; but when heated it gradually causes the formation of mercuric chloride and mercury, the action being very slow with dilute acid, tolerably rapid with concentrated acid. In presence of certain organic substances, the resolu- tion into mercuric chloride and mercury takes place at 38° to 40° C. (100° F.) Free chlorine, and nitric acid, quickly dissolve mercurous chloride, as mercuric salt. In analysis, mercurous compounds are precipitated, from solution, as chloride, in the first group, and this precipitate is distinguished from others in the group, by blackening with ammonium hydroxide. The iden- tification of mercury, by reduction to metallic state, is the same as with mercuric compounds. 433, Reactions of mercurous salts.—Fixed alkali hydroxides pre- cipitate from solutions of mercurous salts, mercurous oxide, Hg20, black, insoluble in alkalies. Solution of ammonium hydroxide produces black precipitates ; that from solution of mercurous nitrate being (NH,Hg2)N03, nitrogen dihydro- gen dimercurous nitrate,* black, insoluble in alkalies, soluble in acids : Hg,(N03)2 + 2NH,HO = NH.Hg,NO3 + NH,NO:i + 3HaO Mercurous chloride, white, is changed by ammonium hydroxide to (NH2Hg2)CI, nitrogen dihydrogen dimercurous chloride, or dimercurous * The compounds produced by action of ammonium hydroxide on mercury compounds are considered as substitutions of Hg for a certain number of atoms of Hin NH4 (ammonium). The substitutions formed fr°m mercurous compounds contain (200 parts by weight or) one atom of Hg (acting as a monad) for each atom (1 part) of H displaced; they are termed mercurous-ammoniums : mercurosammonium being di-mercnrosammonium, NHjHg, ; tri-mercurosammonium, NHHg3retc. The substitutions ormed in ammonium by mercury from mercuric compounds contain one atom of Hg (acting as a dyad) for atoms of H displaced; they are designated as mercurammoniums; mercnrammo-ium being ®H3)2Hg; di-mercurammonium, NH,H? : tri-mercurammonium, (Nfl'.2Hg'a ; tetra-mer •rra nmo- nium. KHgv 108 Meucury. ammonium chloride, black (distinction from lead), decomposed by acids, insoluble in ammonium hydroxide (distinction from silver) : Hg2Cl2 + 2NH4HO = NH2Hg2CI + NH,OI + 3H20 434. Solutions of the carbonates of the fixed alkali metals precipitate an instable mercurous carbonate, Hg2C03, gray, blackening to basic car- bonate and oxide when heated. Ammonium carbonate reacts like ammo- nium hydroxide. BaC03, SrC03, CaC03, and MgCQ3 precipitate mercur- ous salts, in the cold. 435. Hydrosulphuric acid, and soluble sulphides, precipitate, not Hg,2S, but HgS —J— Hgj msoluble in but normal K„S, in presence of KOH, dissolves the HgS, and leaves the Hg as a residue. 436. Hydrochloric acid and soluble chlorides form a white precipitate of mercurous chloride, Hg2Cl2, “calomel”—placing the mercurous base IH THE FIRST GROUP. For relations of the precipitate to solvents, see 432 : to ammonium hydrate, see 433 ; fixed alkalies blacken it by formation of Hg20 (433). 437. Soluble bromides precipitate mercurous bromide, Hg.,Br2, yellow- ish white, insoluble in water and in alcohol, insoluble in dilute nitric acid. 438. Soluble iodides precipitate mercurous iodide, Hg2I2, greenish-yel- low—“ the green iodide of mercury.” The precipitate from mercurous ni- trate contains more or less mercuric iodide ; that from the acetate is nearly pure Hg2I2. Mercurous iodide is nearly insoluble in water, insoluble in alcohol (distinction from mercuric iodide), somewhat soluble in ether, slowly soluble in part by aqueous solutions ■of alkali iodides (excess of the precipitants), being first decomposed to mercuric iodide .and mercury, which last remains undissolved : Hg,I2 + 2KI = Hg + (Kl).jHgI2 Ammonium hydroxide solution slowly decomposes and partially dissolves mercurous iodide. By sublimation, and to some extent by exposure to light, mercurous iodide is changed to mercuroso-mercuric iodide, Hgl.HgI2, yellow—with separation of metallic mercury. When the precipitate by iodide of potassium, in solution of mercurous nitrate, is made in very dilute solutions or is allowed to stand for some time, it consists chiefly of this—“the yellow iodide of mercury.” It is strictly insoluble in alcohol; melts and sublimes unde- composed, and is affected by alkali iodides like mercurous iodide. 439. Alkali cyanides, also hydrocyanic acid, resolve mercurous salts into metallic mercury, a gray precipitate, and mercuric cyanide, which remains in solution.—Ferro- cyanides form a white, gelatinous; ferricyanides, a red-brown precipitate. Alkali phosphates—as Na2HP04—precipitate the white mercurous phosphate, Hg3P04, when the reagent is added in excess; the yellow mercurous phosphate-nitrate, HgsPOi.HgNOs, when mercurous nitrate is in excess.—Chromates precipitate the orange-yellow mercurous chromate, basic; changed by dilute nitric acid to the normal Hg2Cr04; by strong nitric acid changed to mercuric chromate, and dissolved.—Oxalic acid and oxalates precipitate the white mercurous oxalate, Hg2C204, slightly soluble in dilute nitric acid. Mercury. 109 Soluble sulphates precipitate, from solutions not dilute, the white mercurous sulphate, Hg2S04, sparingly soluble in water (433); decomposed by boiling water with precipita- tion of a basic sulphate ; more soluble in dilute nitric acid ; blackened by ammonium hydroxide and fixed alkalies (distinction from other sparingly soluble sulphates). 440. Mercurous compounds are reduced to metal by the same reducing agents that reduce mercuric compounds to metal; but not by all the reduc- ing agents capable of converting mercuric to mercurous combinations, as more fully specified in 448. As to oxidadon of mercurous compounds, see 451.—The reactions in the dry way are nearly the same as those for mer- curic compounds (449). 441. MERCURIC oxide, sulphide, iodide, iodate, basic carbonate, oxal- ate, phosphate, arseniate, arsenite, ferrocyanide, and tartrate are insoluble in water. The bromide is soluble in 250 parts of cold, or one-tenth that pro- portion of boiling water. The acetate and cyanide are freely, the chromate and citrate sparingly, soluble in water. The double iodides of mercury, and the metals of the alkalies and alkaline earths, are soluble in water—that is, mercuric iodide is soluble in aqueous solutions of alkali iodides. The double bromides dissolve in a smaller proportion of water than the bromide. Except the chloride, the ordinary mercuric salts which are soluble in water are so only by presence of free acid being partially decomposed by water, with separation of basic salts (326). In work with solution of mercuric nitrate, some of the reactions are modified by the free acid, always pres- ent.—Mercuric sulphate is soluble in very dilute sulphuric acid.—The chloride is soluble in about 12 parts of cold, or two to three parts of boiling water ; freely soluble in alcohol and in ether. In analysis, the second-group precipitate of mercury sulphide is separat- ed by its insolubility in dilute nitric acid. The final form, in determina- tion of mercury, is usually the metallic state (448 a, or 449). 442. Reactions of mercuric salts.—Solutions of the fixed alkali hy- droxides, added, short of saturation, to solutions of mercuric salts, pre- cipitate reddish-brown basic salts ; when the reagent is added to snpersatu- ration, the orange-yellow mercuric oxide, HgO, is precipitated. Prepared m the dry way, mercuric oxide is obtained red—the “red precipitate” of the shops. From very acid solutions, the precipitate is incomplete or does not form at all, owing to its solubility in alkali salts. It is very slightly soluble in water. In presence of an ammonium salt, the white precipitate (443) is formed. Certain organic acids interfere with the precipitation. 443. Ammonium hydroxide produces a “ white precipitate,” recog- nizable in very dilute solutions; that with neutral solution of mercuric chloride being (NH2Hg)CI, nitrogen dihydrogen mercuric chloride (a) ; that with hot dilute solution of mercuric nitrate and excess of ammonia being (NHg2)NOa, nitrogen dimercuric nitrate (b). The precipitates are easily soluble in hydrochloric acid ; sparingly soluble in strong ammonium 110 Mercury. hydroxide, which should not be used in excess in precipitation. They are also more or less soluble in ammonium salts, and especially in ammonium nitrate. Therefore, the precipitation by ammonium hydroxide is always in some degree incomplete ; and that of (he acid mercuric nitrate is decidedly diminished, and in very dilute solutions prevented altogether, by the am- monium salt formed in the reaction (as shown in equations a and b). A soluble combination of ammonium chloride with mercuric chloride, (M(- Cl)2HgCl2, or ammonium mercuric chloride, called “sal alembroth,” is not precipitated by ammonium hydroxide, but potassium hydroxide precipitates therefrom the white mercurammoniurn chloride, (H2HcHg)CI2 (c). a. HgCl2 + 2NH4OH = (NH2Hg)CI + NH4CI + 2H20 b. 2Hg(NQ3)2 + 4NH.OH =: (NHg.)NO3 + 3NH4N03 + 4H20 c. (NH4OI)2HgCI2 + 2KOH = ([NH3]2Hg)CI2 + 2KCI + 2H20 Ammonium carbonate reacts like ammonium hydrate. 444. Potassium and sodium carbonates precipitate first red-brown basic salts, which, by excess of the precipitants with heat, are converted into the yellow mercuric oxide. The basic salt formed with mercuric chlo- ride is an oxychloride, HgCl,2.(HgO)2, 3, or 4 ; with mercuric nitrate, a basic carbonate, (Hgo)3HgC03 or (HgO)4CO,2—-Barium carbonate precipitates a basic salt in the cold, from the nitrate, but not from the chloride. 445. Hydrosulphuric acid, gradually added to solutions of mercuric salts, forms at first a white precipitate, soluble in acids and in excess of the mercuric salt; by further additions of the reagent, the precipitate becomes yellow-orange, then brown, and finally black, insoluble in hydrochloric or dilute nitric acid. This progressive variation of color is characteristic of mercury, and is also produced by ammonium sulphide. The final and stable precipitate is mercuric sulphide, HgS; the lighter colored precipi- tates consist of unions of the original mercuric salt with mercuric sulphide, as HgCl2.HgS, the proportion of HgS being greater with the darker pre- cipitates. When sublimed and triturated, the black mercuric sulphide is converted to the red (vermilion), without chemical change. Mercuric sulphide is soluble by free chlorine (nitro-hydrochloric acid) (a) ; not affected by dilute nitric acid (distinction from all other metallic sulphides) or by hydrochloric acid ; insoluble in ammonium sulphide (dis- tinction from tin, antimony, arsenic); insoluble in HaOH, and in Na2S, but dissolves in a mixture of the two (M. C. Mehu, J. 8., 1876). 446. Soluble bromides precipitate, from concentrated solutions of mer- curic salts, the white Hgßr2, soluble in 25 parts of hot and in 250 parts of cold water. Also soluble in excess of mercuric salts, and in excess of the bromide by which the precipitate is formed ; hence, unless added in suit- able proportions, no precipitate will be produced. The precipitate is de- composed by strong or hot nitric acid. a. HgS + Cl 2 = HgCl2 + S Mercury. Soluble iodides precipitate mercuric iodide, Hgl2, first reddish-yellow, then red ; very slightly soluble in water, soluble in concentrated nitric and hydrochloric acids ; quickly soluble in solutions of the iodides of all the more positive metals—that is, in excess of its precipitants, by formation of soluble double iodides ; as variable to KXHgI2>* The dipotassium mercuric tetraiodide (K2Hgl4) (sometimes designated the iodo-hydrargyrate of potassium) is precipitated by ammonium hy- droxide, and by the alkaloids (see Kessler’s Test, 40). Dilute acids cipitate the mercuric iodide. 447. Soluble normal chromates precipitate, from very concentrated solutions, basic mercuric chromates, orange yellow to red ; considerably sol- uble in water, more soluble in solution of mercuric chloride or nitrate. Soluble phosphates, as Na2HP04, precipitate mercuric phosphate, Hg3- (P04)„ white, soluble in acids, including phosphoric acid, and in ammo- nium salts. Soluble oxalates, and oxalic acid, precipitate—from the nitrate, but not from the chloride—mercuric oxalate, HgC204, white, readily soluble in dilute hydrochloric acid, difficultly soluble in nitric acid. HgCl2 boiled in the sunlight with (NH4)2CsO4 gives Hg2Cl2 and C02. Po- tassium Perrocyanide precipitates mercuric ferrocyanide, white, becoming blue on standing. 448. Reducing agents precipitate, from the solutions of mercuric and mercurous nitrates, dark-gray Hg°; from solution of mercuric chloride, or in presence of chlorides, first the white, Hg2Cl2, then gray Hg. Strong acidu- lation with nitric acid interferes with the reduction, and heating promotes it. By digestion with hot concentrated hydrochloric acid—-and a little so- lution of stannous chloride—the gray precipitate of divided mercury is con- verted into liquid globules of metallic lustre. This somewhat tardy result is hastened by trituration with a glass rod in the test-tube; or first wash and then dry, when trituration will accomplish the object at once. The reducing agent most frequently employed is stannous chloride (a). Boiling solution of sulphurous acid (b) effects the reduction. A clean strip of copper, placed in a slightly acid solution of a salt of mercury, becomes coated with metallic mercury, and when gently rubbed with cloth or paper presents the tin-white lustre of the metal (c), the coating being driven off by heat. Zinc and iron, also, reduce mercury, and from mercuric chloride or in presence of chlorides, first precipitate calomel. Formic acid reduces mercuric to mercurous chloride, and in the cold does not effect further reduction. Dry mercuric chloride, moistened with alcohol, is reduced by metallic iron, a bright of which is corroded soon after immersion into the powder tested (a delicate distinction from mercurous chloride). * A hot concentrated solution of potassium iodide dissolves 3HgT2 for every 3KI. The first crystals from this solution are KI.HgT3. These are decomposed by pure water, and require a little free iodide for perfect water solution, but they are soluble in alcohol and in ether. 112 Mercury. a. 2HgCl2 + SnCl2 = Hg2Cl2 + SnCl4 Hg2Cl2 + Sno12 =7 3Hg + SnOI4 Or: HgCl2 + SnCl2 = Hg + SnCl4 Also: 2Hg(N03)2 + SnCl2 = Hg2Cl2 + Sn(NO3)4 b. HgCl2 -4- H2SO3 + H2O =Hg + H2SO4 + 2HCI c. SHgNOs +Ou = 2Hg + Cu(NO3)2 (compare 346). 449. All compounds of mercury, in glass tubes or on charcoal, are quickly volatile before the blow-pipe. Mercurous chloride and bromide and mercuric chloride and iodide sublime (in glass tubes) undecomposed—the sublimate condensing (in the cold part of the tube) without change. Most other compounds of mercury are decomposed by vaporization, and give a sublimate of metallic mercury (mixed witli sulphur, if from the sulphide, etc.) All compounds of mercury, dry and intimately mixed with dry so- dium carbonate, and heated in a glass tube closed at one end, give a sub- limate of metallic mercury as a gray mirror coat on the inner surface of the cold part of the tube. Under the magnifier, the coating is seen to consist of globules, and by gently rubbing with a glass rod or a wire, globules visi- ble to the unaided eye are obtained. 450. Estimation.—(l) As metallic mercury. The mercury is reduced by means of CaO in a combustion-tube at a red heat in a current of C02. The sublimed mercury is condensed in a flask of water, and, after decanting the water, dried in a bell-jar over sulphuric acid without application of heat. The mercury may also be reduced from its solution by SnCl2 (or H3POs at 100° 0.) and dried as above. (2) As mercurous chloride. It is first re- duced to (HgJ" by H3P03, which must not be heated above 60° C., other- wise metallic mercury will be formed ; and after precipitation by HCI and drying on a weighed filter at 100° C. it is weighed as Hg,2Cl2. Or enough HCI is added to combine with the mercury, then the Hg" is reduced to (HgJ" by EeS04 in presence of NaOH. 2HgO + 2FeO + 3H.O=: Hg.O + Fe2(0H)6 h2so4 is added, which causes the formation of Hg2Cl2, which is dried on a weighed filter at 100° C. (3) As HgS. It is precipitated by H2S, and weighed in same manner as the chloride. Any free sulphur mixed with the precipitate should be removed by CS2 or Na2S03. (4) As HgO, by heat- ing the nitrate in a bulb-tube in a current of dry air not hot enough to de- compose the HgO. (5) Volumetrically, by Ha2S203 ; from the 'nitrate the precipitate is yellow, from the chloride it is vdiite. BHg(NO3)2 + 2Na25,03 + 2H20 = Hg3S2(N03)2 + 2Na2SO, + 4HNOs BHgO12 + 2Na2S203 + 2H20 = Hg3S2Cl2 + 2Na2SO4 + 4HCI (6) Volumetrically, HgCl2 is reduced to Hg20 by PeS04 in presence of KOH, and after acidulating with H2SO, the excess of FeS04 is determined by K2Or2O7 or K2Mn20, (7) By iodine. It is converted into Hg2Cl2 and then dissolved in a graduated solution of I dissolved in KX HgaCl2 + 6KI + I„ = 2K2HgI4 + 2KCi Mercury. The excess of iodine is determined by Na2S203. (8) The measured solu- tion of HgCl2 is added to a graduated solution of KI; 4KI + HgCl„ =■ K2HgI4 ■4- 2KCI. The instant the amount of HgCl2 shown in the equation is exceeded a red precipitate of Hgl2 appears. 451. Oxidation.—Free mercury (Hg°) precipitates the free metals from solutions of Ag, Au, and Pt; and is precipitated as Hg° by Zn, Al, Mg, Cd, Bi, Co, Sn, Sn'', Cu, and (Cu2)", also by free H3PO„, H3P03, and H2S03. (Hga)" is oxidized to Hg" by Br, Cl, I, HNOs, HaS04 (hot), and HCIO,. 114 First- Group Metals. Pb Ag (Hga)" Hg" KOH or NaOH in excess Solution (887). Ag2C, grayish brown (411). Hg.O, black (438). HgO, yellow to red (443). White precipitate. Not formed in the acetate (387). PbCl2, sparingly soluble in water, more freely if hot (390). PbBr-j, slightly soluble in water (391), Solution (411). AgCl, white (413). AgBr, pale yellowish- white (414). (NH2Hg2)N03,black(435). Hg2Cl2, white, dissolved by oxidizing acids (436). Hg2Br2, yellowish-white (437). (NHg2)N03, white ; (NHaHg)01, white (443). Solution. (446). Iodides . . . . Pbl2, orange-yellow, spar- ingly soluble in alkali iodides (393). PbSCb, least soluble sul- phate except barium (389). PbCr04, yellow (893). Agl, pale-yellow, freely soluble in alkali iodides (414). No precip., except in con- centrated solutions(338). Ag2Cr04, purple-red (430). Hg2I2, yellowish-green, decomposed by alkali iodides (438) No precipitate in dilute solutions (439), Hg2Cr04, orange-yellow (439). Hgl2, red, freely soluble in alkali iodides (446). Solution requires a little free acid (441). Precipitate only in con- cent. solutions (447). H3SO1 and Sulphates Chromates . . . . Zinc Pb (400). Ag (435). Hg (448). Hg. Hg"2 salts, then Hg (448). Hg''2 salts, then Hg (448). Lead Ag- Ag- Stannous Salts Ag- Hg. (Hg2)" salts, then Hg. 452. Comparison of Certain Reactions of First-Group Metals. Taken in Solution of their Nitrates or Acetates. A RSENIC. ARSENIC. As == 74.918. 453. Specific gravity, pure crystalline, 4.71 (Bettendorf, 1867); after fusion, under great pressure, 5.71 (Mallet, 1872). Vapor density (H =l) 147.2 (Deyille and Troost, 1863); therefore the molecule is assumed to contain four atoms (As 4), Volatilizes in an atmosphere of coal-gas without melting at 450° C. (Conechy, 1880). Fuses under great pressure between the melting point of Ag and Sb (Mallet, 1872),. Valence, a pentad in H3Asv04 ; a triad in As4Oe; oxidation valence in free arsenic, zero (As0), and in arsenious hydride a negative triad (As-/"H3). 454. Occurrence.—Found native; also as an alloy Avith other metals, as PeAs2, NiAs, CoNiAs2; as realgar (As2S2), orpiment (As2S3), and mispickel (Pe„AsS2). It is sometimes found as As4Oe, as an arsenate in cobalt bloom (C03(A50.)2), and in a great variety of minerals. Preparation.—(l) Reduced from its oxide by carbon, As406 +3O = As4 + 3C02. (2) From Pe2AsS2 by simple ignition, air being excluded, 4Pe2AsS2 = BPeS + As4. (3) From As2S3 by fusion Avith Ha2C03 and KCN. 6Na2CO3 + 2As.S3 + 6KCN = As4 + GNa2S + 6KCNO + 0CO2 455. Properties.—Arsenic is by some chemists classed with metals, by others with the non-metallic elements. Its failure to act as a base with oxyacids determines definitely its non-metallic character. The amorphous arsenic is black, the crystalline a steel-gray, brittle and pulverizable. Its vapor is yellow, with a strong oppressive and poisonous alliaceous odor. It is slowly oxidized in moist (not in dry) air at ordinary temperatures ; when heated in the air, it burns with a bluish flame, and becomes the white arsenious anhydride, As4Oe. It readily combines with chlorine and bromine upon contact, and with iodine and sulphur by aid of heat. It is not attacked by aqueous hydrochloric acid at ordinary tempera- tures, and but slightly when hot and concentrated and with air ; it is slowly oxidized to arsenic acid by hot concentrated sulphuric acid, or more I‘eadily by nitric acid ; but its proper solvent is nitro-hydrochloric acid, or chlorine with Avater, by which it is oxidized to arsenic acid with violent rapidity (a). Hot solution of potassium or sodium hydrate dissolves it as arsenite (t>): a. As4 + 10C12 + 16H20 = 4H3As04 + 20HC1 h. As4 + 12KOH = 4K3AsO3 + 6H2 Arsenic forms tAvo oxides, both acidulous : arsenious anhydride, ■^■s; 4^6, representing a series of arsenious compounds and arsenites of Petals ; and arsenic anhydride, Asv 206, forming arsenates of metals, and arsenic acids, and representing other arsenic compounds. Both these classes °f compounds possess considerable stability ; the arsenious bodies acting aa efficient reducing agents, and the arsenic substances, with less activity, as oxidizing agents. 116 A useNic. 456. Arsenious Acid and Anhydride.—The vapor density of the anhydride (arsenious oxide) is 198 at a white heat (V. Meyer, Ber. Chem. Ges., 1879, 1117), therefore its molecule is assumed to be As4Oe. Its solu- tion in water is supposed to be arsenious acid, but on evaporation only As406 remains. Many arsenites are known, but the instability of the greater number of them has prevented an accurate determination of their composition. Their composition is best explained by regarding them as derived from one of the three hypothetical arsenious acids. (1) Orthoar- senious acid, H3As03; forming normal, acid, and basic ortlwarsenites—■ e.g., Ag3AsQ3, Mg3(As03)2, Co3H6(As03)4, CuHAs03, Fe809(As03)2, etc. (2) Pyroarsenions acid, H4As,205 ; forming pyroarsenites—e.g., (NH4)4- As206, Ca2As206, K4As205, etc. (3) Mefarsenous acid, HAs02; forming metarsenites—e.g., WH4As02, KAs02, Ca(Aso2)2, Pb(Aso2)2, etc. The triad arsenic forms some combination with nearly all metals, but many of its salts are so instable as to prevent the determination of their composition. 457. Arsenious anhydride—having both crystalline and amorphous modifications— is very slowly and sparingly soluble in cold water, much more quickly but quite spar- ingly soluble in hot water, the solution feebly reddening litmus; freely soluble in hydro- chloric acid, and somewhat soluble in sulphuric acid without combination; readily solu- ble in alkali hydroxides with combination, forming, perhaps, the hypothetical K3As03, but only KH(AsO2)2, KAsO2, and K4As205 have been isolated; slightly soluble in alco- hol, and soluble in glycerine.—Arsenious chloride is wholly decomposed by water, with formation of arsenious oxide and hydrochloric acid (equation a); arsenious sulphide is very slightly soluble in pure water, insoluble in acidulated water, but soluble by combi- nation in solutions of alkalies (6), alkaline carbonates and alkaline sulphides (c), (d) and (e).—Arsenites of the alkali metals are soluble in water; of the alkaline earth metals sparingly soluble, of magnesium insoluble; of all other metals, insoluble. The arsenites are decomposed—and, except those of first-group metals, dissolved—by hydrochloric acid, and are decomposed and dissolved by nitric acid, without exception: a. 4AsC13 + 6H20 = As 406 + 12HC1 h. 2As2S3 + 4KOH = KAsQ2 + 3KAsS2 + 2H20. Watts’ Diet., 2d ed., 1, 315. c. As2S3 + K2S = 2KAsS2 d. As2S3 + 6NH4HS = 2(NH4)3AsS3 + 3H2S* e. As2S3 + 2(NH4)2S = (NH4)4As2SS In analysis, the second-group precipitate of arsenious sulphide is sepa- rated with antimony and tin, by solution with ammonium sulphide. The final determination and separation from antimony is usually effected by the action of the hydrides upon solution of silver nitrate. 458. Alkali hydrates and carbonates do not precipitate arsenious compounds from solution ; whereby arsenic is distinguished from the bases. * Dibasic and monobasic as well as tribasic thioarsenites are formed in different conditions. Accord- ing to Nilsson (Bericht. d. deut. chem. Ges., IV., 989 ; Jour. Chem. Soc., X., 1872, 599), (NET,iHS always dissolves As3S3 as (NH4i2S.(As2S3)3. For the action of alkalies and alkaline carbonates upon As3S2, As2S3, As2Sb, and the formation of many thioarsenites and thioarsenates according to Nilsson, see Jour. prakt. Chem., 1876 [2], 14, 1-60 and 145-172. Arsexic. 117 459. Hydrosulphuric acid precipitates the lemon-yellow arsenious sul- phide, As2S3. The precipitate forms promptly in acidulated solutions, the most perfectly with hydrochloric acidulation ; being complete even in strong hydrochloric acid solution, but diminished by too strong nitric acid. It forms slowly in simple aqueous solution of arsenious acid, as a color rather than a precipitate ; being slightly soluble in pure water, but insol- uble in acidulated water. It is not formed in solutions of alkali arsenites, except by acidulation. Citric acid and other organic substances hinder, but, in presence of much hydrochloric acid, do not wholly prevent its for- mation. Alkali sulphides produce, and by further addition dissolve, the precipitate. The arsenious sulphide is soluble in solutions of alkali hydrates, carbon- ates, and sulphides, as severally explained in 457. From all these alkaline thioarsenites, acids reprecipitate the sulphide (a and b). By its solubility in solution of ammonium sulphide, it is separated with antimony and fin from the other members of group second ; and by its solubility in solution of ammonium carbonate, it is approximately separated from antimony and tin, in a process of separation which has been in common use : a. 2(NH4)BAsS3 + 6HCI = As2S3 + 6NH4CI + 3H2S b. (NH4)4As2SS + 4HCI = As2S3 4- 4NH4CI + 2H2S The color of the As2S3 distinguishes it from S, derived thus : The arsenious sulphide is also soluble in solutions of alkali sulphites with free sul- phurous acid (distinction, and a method of separation, from antimony and tin): c. 2(^H4)252 + 4HCI = S2 + 4NH4CI + 2H2S d. 4As2S3 + 82KHS03 = BKAsO2 + 12K2520s + 352 + 14SO, + 16HaO Like raetalloidal arsenic, the arsenious sulphide is insoluble in hydrochloric acid— another means of separation from antimony and tin. It is insoluble in dilute but dis- solves in strong nitric acid, and by free chlorine or nitro-hydrochloric acid, as arsenic anhydride, As2Os, or arsenic acid, H3AsO4—(equations e and /. Usually, however, a large portion of the S is oxidized to H2SO4, and completely if the Cl or HNC3 is used hot and in excess (g). Compare equation a, 455). Arsenious sulphide is not changed to arsenious oxide by any solvents. /. 6As2S3 + 20HN03 ■+■ BHsO = .2H3AsO4 + 9S2 + 20NO g. As253 + 14C12 + 20H20 = 2H3AsO4 + 3H2SQ4 + 28HC1 Thiosulphates—as Na2S2Q3—also precipitate, from boiling hydrochloric acid solu- tion of arsenious acid, the arsenious sulphide (distinction from tin): 2H3AsO3 + 3Na25203 = As2S3 + 3Na2SO4 + 3H20 e. 2As253 + lOCI 2 + l(iH2O = 4H3AsO4 + 352 + 20HC1 460. Silver nitrate solution precipitates from neutral solutions of ar- -B(mites, or ammonio silver nitrate* from water solution of arsenious oxide, silver arsenite, Ag3As03, yellow, readily soluble in dilute acids or in am- monium hydroxide or ammonium salts (420) : H3As03 + 3AgNO3 + 3NH4OH = Ag3As03 + BNH4NO3 -1- 3HsO * Prepared by adding ammonium hydroxide to the solution of silver nitrate, till the precipitate at first Produced is nearly all redissolved. 118 Arsenic. 461. Copper sulphate solution precipitates from solutions of neutral arsenites, or aramonio copper sulphate (prepared as directed in note under 460) precipitates from water solution of arsenious oxide, the green copper arsenite, OuHAsO3 (Scheele’s green), soluble in ammonium hydroxide and in dilute acids. Copper acetate in boiling solution precipitates the green copper aceto-arsenite (CuOAs2O3)3Cu(C2H302)2 (Schweinfurt green), soluble in ammonium hydroxide and in acids. Both these salts are often desig- nated as Paris green (343). For the reaction of Copper Salts with fixed alkali arsenite, see 475. 462. In general, solutions of arsenites are precipitated by solutions of normal salts of the metals, except those of the alkalies, and barium, stron- tium, and calcium (457). Normal magnesium salts form a white precipi- tate of magnesium arsenite. The precipitate is soluble in ammonium hy- droxide and ammonium chloride (distinction from arsenates). Ferric salts precipitate from arsenites, and recent ferric hydroxide {used as an antidot ) forms, with arsenious anhydride, variable basic fer- ric arsenites, scarcely soluble in acetic acid, soluble in hydrochloric acid. Water slowly and sparingly dissolves from the precipitate the arsenious an- hydride; but a large excess of the ferric hydroxide holds nearly all the arsenic insoluble. To some extent, the basic ferric arsenites are transposed into basic ferrous arsenates, insoluble in water, in accordance with the reducing power of arsenious oxide. 463. Arsenic is reduced to the elemental state by several methods of great analytical importance. By the action of hydrogen generated in acid solution (Marsh’s Method) it is reduced from all its soluble compounds, when it enters into a combination with hydrogen as arsenious hydride, AsH3, gaseous. The latter can be identified by numerous reactions, and from it the arsenic can readily be obtained free. The hydrogen is generated by sulphuric acid diluted with 6 to 8 parts water, and zinc (both free from arsenic). Compare 288. In general nascent hydrogen (or any metal with any acid which produces hydrogen) gives like results. The hydrogen removes the oxygen, from either oxide of arsenic, by forming water, and then combines with the arsenic ; two atoms of hydrogen taking the place of one atom of oxygen ; or, including the zinc and sulphuric acid ; As4O0 + 12H2 = 6H20 + 4AsH3 H3As03 + 3(Zn + H2S04) = 3ZnSO4 + 3H20 + AsH3 H3AsQ4 + 4(Zn + HoS04) = 4ZnSO4 + 4H20 + AsH3 It will be seen that arsenious hydride cannot be formed in presence of free chlorine or other oxidizing agents, such as nitric acid, nitrates, chh - rates, and hypochlorites. Sulphides and sulphites interfere; also mercury salts (by amalgamation of the zinc), and most organic substances. Arseni- Arsenic. 119 ous sulphides are not acted on by the nascent hydrogen. With zinc, strong potassium hydroxide or sodium hydroxide may be used instead of acid, the action being slower ; but quite rapid if finely divided iron, platinum, or some other metals are present. Sodium amalgam alone, in solutions neutral, acid, or alkaline, causes an abundant generation of arsenious hy- dride if arsenic is present. In the test made by sodium amalgam, in alka- line solution, cane sugar and some other organic bodies do not interfere. The generation of arsenious hydride, by metallic magnesium, when done in strong solution of ammonium chloride, is a separation from anti- mony. The solution may be neutral or alkaline, but, for the separation, not acid. Metallic aluminium, in strong potassium hydroxide solution, on warm- ing, generates arsenious hydride from arsenical compounds (distinction and separation from antimony). 464. Arsenious hydride (arsine) burns when a stream of it is ignited where it enters the air, and explodes when its mixture with air is ignited, like other combustible gases. It burns in a stream, with a somewhat lumi- nous and slightly bluish flame (distinction from hydrogen) ; the hydrogen being first oxidized, and the liberated arsenic becoming incandescent, and then undergoing oxidation ; the vapors of water and arsenious anhydride passing into the air (a). If a piece of cold, porcelain is held in the flame, the reduction of temperature prevents the oxidation of the arsenic, which is deposited in dark steel-gray spots, adherent to the porcelain, about which ft little of the water of combustion condenses (i) : a. 4AsH3 + 602 = As4o6 -+- 0H2O b. 4AsH3 + 302 = As 4 + 6H20 In many particulars above mentioned, the combustion of arsenious hydride resembles that of the hydrocarbons of illuminating gas. Arsenious hydride is an exceedingly poisonous substance, the inhalation of the un- fixed gas being quickly fatal. Its dissemination in the air of the laboratory, even in the small portions which are not appreciably poisonous, should be avoided. Furthermore, as lfc is recognized or determined, in its various analytical reactions, only by its decomposi- tion, to permit it to escape undeeomposed is so far to fail in the object of its production. The evolved gas should be constantly run into silver nitrate solution, or kept burning. 465. Arsenious hydride is decomposed by heat alone. In passing through glass tubes, heated to incipient redness, the gas is decomposed, the arsenic adhering to the lnner surface of the tube, beyond the heated part, as a steel-gray mirror coaling. This coating is readily driven by the heat, is gradually dissipated by hot hydrogen gas, and miparts the garlic odor to the escaping hydrogen gas. The latter, if ignited, will gene- rally deposit arsenic spots on porcelain, showing that the arsenic is not wholly retained ln the tube. 466. Both the mirror and the spots exhibit the properties of free arsenic (455). Jiquid reagents are most convenient for application to the spots. The reactions of these eposits having analytical interest are such as distinguish arsenic from antimony. 120 Arsenic. Comparison of Arsenic and Antimony, deposited from AsH3 and SbH3. Of a steel gray to black lustre. Volatile at 450° C.; as arsenious acid, at 318° C. Dissolve in hypochlorite (a). Warmed with a drop of ammonium sul- phide, form yellow spots (459), soluble in ammonium carbonate, insoluble in hydro- chloric acid. With a drop of hot nitric acid, dissolve clear (455) The clear solution, with a drop of solu- tion of silver nitrate, when treated with vapor of ammonia (from a glass rod moist- ened with ammonium hydrate and held near), gives a brick-red or a yellow color (391). With vapor of iodine, color yellow, by formation of arsenious iodide, readily vola- tile when heated. Arsenic Spots. Of a velvety brown to black surface. Volatile in vacuo at white heat; by oxida- tion, at a red heat. Do not dissolve in hypochlorite. Warmed with ammonium sulphide, form orange-yellow spots, insoluble in ammonium carbonate, soluble in hydrochloric acid. With a drop of hot dilute nitric acid, turn white. The white fleck, treated with silver ni- trate and vapor of ammonia, gives no color until warmed with a drop of ammonium hydrate, then gives a black color. With vapor of iodine, color more or less carmine red, by formation of antimonious iodide, not readily volatile by heat. Antimony Spots. Arsenic Mirror. Antimony Mirror. Deposited beyond the flame ; the gas being decomposed by a red heat (465). The mirror is driven- at 450° C.; it does not melt. By vaporization in the stream of gas, escapes with a garlic odor. By slow vaporization in a current of air (the tube open at both ends and held in- clined over the heat), a deposit of octahedral crystals is obtained above—if abundant, forming a white coating (457), soluble in water, the solution giving reactions for ar- senic. Deposited before or on both sides of the flame; the gas being decomposed consider- ably below a red heat. The mirror melts to minute globules at 43311 C., and then is driven at a red heat. The vapor has no odor. By vaporization in a current of air, a white amorphous coating is obtained—in- soluble in water, soluble in hydrochloric acid, and giving reactions for antimony. a. The hypochlorite reagent—usually NaCIO—decomposes in the air and light, by keeping. It should instantly and perfectly bleach litmus-paper (not redden it). It dis- solves arsenic by oxidation, to arsenic acid: As 4 + IONaCIO + 6H20 = 4H3AsO4 + IONaCI 467. When arsenious hydride is passed into solution of silver nitrate, the silver is reduced to metal by the oxidation of both elements in the gas —the hydrogen to water, and the arsenic to arsenious acid, which remains in solution along with the liberated nitric acid (distinction from antimo- nious hydride, which precipitates antimonious argentide) : AsH- + 6AgN03 + BH2O = 6Ag + H3As03 + 6HNOs A RSENIC. 121 The reactions for the arsenious oxide formed in solution should be ob- tained after filtering out the brown-black precipitate of silver, then adding a very little hydrochloric acid, that the silver in the undecomposed nitrate may be removed as chloride. From the filtrate, hydrosulphuric acid preci- pitates the sulphide, and arsenic may be quantitatively determined from the weight of this precipitate, after Marsh’s Test. Reliance should not be placed on blackening of the silver nitrate alone, as this may be due to SbH3, or to H2S or to PH3. H2S would be generated in the test, from sul- phides ; and PH3 from hypophosphites or phosphites. If the material treated with zinc and dilute sulphuric acid be placed in a flask or large test-tube, and a paper moistened with silver nitrate be tied over the mouth, it will (on standing) be blackened by arsenious hydride. The interference of hydrosulphuric acid may be avoided by causing the gas to pass through cotton-wool, moistened with solution of lead acetate, and carefully placed to fill the neck of the vessel, then left seve- ral hours. This operation may be relied on for negative results, in testing the purity of reagents, etc. The yellow silver arsenite (460) may be obtained as a distinctive test, with the silver nitrate left in solution, undecomposed by the arsenious hydride, after filtering out only the metallic silver, by the careful addition of ammonium hydroxide, in repeated small portions, by the glass rod, till the nitric acid and arsenious anhydride are just neu- tralized : H3AsOs + 3AgNOs + 6HNO3 + 9NH4OH = Ag3As03 + 9NH4NO3 + 9H30 Arsenious hydride received in nitric acid is changed to H3AsO4, soluble in water (separation from antimony). 468. Stannous chloride, SnCl2, reduces arsenious and arsenic oxides, from hot con- centrated hydrochloric acid solution, as flocculent, black-brown, metalloidal arsenic, con- taining three or four per cent, of tin (Bettendorf’s Method). The arsenic, in solution with the concentrated hydrochloric acid, acts as arsenious chloride (457 a) : The hydrochloric acid should be 25 to 33 per cent.: if not over 15 to 20 per cent., the reaction is slow and imperfect. Sulphuric acid with sodium chloride may be taken instead of hydrochloric acid : In a wide test-tube place 0.1 to 0.2 gram. (2 or 3 grains) of the (oxidized) solid or solution to be tested, add about 1 gram. (15 grains) of sodium chloride, and 2 or 3 cub. centim. (about one fluid drachm) of sulphuric acid, then about 1 gram. (15 grains) of crystallized stannous chloride; agitate, and heat to boiling several times, and set aside for a few minutes. Traces of arsenic give only a brown color; notable proportions give the flocculent precipitate. A dark-gray precipitate may be due to mercury (448 a), capable of being gathered into globules. If a precipitate or a darkening occurs, obtain conclusive evidence whether it contains arsenic or not, as follows : Dilute the mixture with ten to fifteen volumes of about 12 per cent, hydrochloric acid (equal parts of Fre- senius’s Reagent and water) ; set aside, decant ; gather the precipitate in a wet filter, wash it with a mixture of hydrochloric acid and alcohol, then with alcohol, then with a little ether, and dry in a warm place. A portion of this dry precipitate is now dropped into a small hard-glass tube, drawn out and closed at one end. and heated in the thune : arsenic is identified by its mirror (466), easily distinguished fioin mercury (449). Anti- niony is not reduced by stannous chloride ; other reducible metals give no mirror 111 the 4AsCl* -f- 6511C12 AS4 -f- 6511C14 Arsenic. reduction-tube. Small proportions of organic material impair the delicacy of this reac- tion, but do not prevent it. It is especially applicable to the hydrochloric acid distillate, obtained in separation of arsenic, according to 474. 469. Metallic copper reduces arsenious oxide, from hydrochloric acid solution, as an iron-gray film or crust of arsenic with copper, 33 per cent, arsenic, or CusAs2 (Reinsch’s- Method). The copper should be in bright strips, the solution hot, and the reaction awaited for some time. If much arsenic is present, the crust peels off in black scales. The crusts are not evidence of arsenic without further examination—according to 470, etc.—as antimony, silver, and other metals are reducible by copper. The film may be obtained and afterwards determined as arsenic, when but the 0.0005 gram is taken in pure hydrochloric acid solution. 470. lii Marsh’s Test, a portion of the arsenic, reduced by tbe zinc to the elemental state, remains for a short time, while the arsenic is in excess in the solution, as a grayish-black film upon the zinc. If the generation of hydrogen be continued after the arsenic is all reduced, all the latter soon forms arsenious hydride. The deposition of antimony, in Marsh’s Test, is much greater than that of arsenic. Also, if the operation be conducted in a platinum vessel or with platinum foil, in contact with the zinc, the re- duced arsenic does not adhere to the platinum as firmly as the reduced an- timony. 471. Potassium cyanide, with sodium carbonate, reduces arsenic from all ift com- pounds, in the dry way; As406 + GKCN = As 4 + GKCNO 2As,S3 + GKCN = As4 + GKCNS 2As253 + GNa,CO3 + GKCN = As 4 + GNa.S + GKCNO + 6C02 If this reduction be performed in a small reduction-tube with a bulb at th<* end, the reduced arsenic sublimes and condenses as a mirror (406) in the cool part of the tube. The presence of compounds of manganese, bismuth, zinc, or antimony hinders this reac- tion, but does not prevent it. The test can be performed in presence of mercury com- pounds, but more conveniently after their removal ; in presence of organic material, it is altogether unreliable. If much free sulphur is present, H. Rose recommends that the arsenic should be removed from it, by dissolving in ammonia, evaporating the solution to dryness, oxidizing to arsenic acid with hydrochloric acid and potassium chlorate (175), precipitating with ammonium hydroxide and magnesium solution as arsenate (478), and washing and drying the latter for the test. The thoroughly dried substance is mixed with six times its bulk of a dry mixture of equal parts of anhydrous sodium carbonate and potassium cyanide, and introduced into the bulb of the reduction-tube, which should not be over half filled. Heat the bulb very gently over the flame, and if water rises and condenses in the tube, thoroughly dry the bulb and tube—wiping the inside of the tube with twisted paper. Then heat strongly, while the tube is held inclined, finally to a full red heat. If arsenic is present, the mirror will be seen above the bulb, and can be tested, as stated in 466, etc. This operation becomes a more delicate test, and excludes antimony from the mirror, if the mixture be placed in a larger horizontal reduction-tube, drawn out narrow at one end, and connected at the other with an apparatus for generating and drying carbonic anhydride, which is passed over the substance during the reduction (Method of Fhese- Arsenic. 123 nius and Babo): Three parts of anhydrous sodium carbonate, with one of potassium cyanide, are taken, and ten or twelve parts of this mixture to one part of the substance tested, the whole well mixed and thoroughly dried (in the water-oven). The reduction- tube should be about 1.25 centimeters (one-half inch) wide and 10 to 15 centimeters (four to six inches) long, besides the drawn-out part. At the end not drawn out it is con- nected with a small wash-bottle, for sulphuric acid, and this connected with the flask for generating carbonic anhydride with marble and dilute hydrochloric acid. The dried mixture is introduced in the middle of the reduction-tube, by aid of a paper gutter ; the connections made, and the substance again dried by gentle heat. When the atmosphere is expelled and a steady stream of carbonic anhydride is passing through the apparatus, heat the tube between the mixture and the drawn-out end to redness, and then heat the mixture gradually to redness of the tube, driving the mirror to the narrowed portion of the tube. Finally, detach the tube, close the small end in the flame, and advance the heat up to the mirror. 472. Charcoal reduces arsenious oxide very readily, by heat in the glass tube, A small hard-glass tube is drawn out at one end, the extremity closed in the flame, and a particle of the well-dried material dropped into the tube, so that it will fall to the end of the narrow part. A fragment of recently burned charcoal is pushed down nearly to the substance, and heat applied, first to the charcoal and then to the substance, to redness. The mirror forms just above the heated part, and may farther be tested as stated in 466. During the reduction, the garlic odor is observed. All compounds of arsenic, heated with sodium carbonate on charcoal, and all oxi- dized compounds heated on charcoal alone, present the odor of arsenic. Non oxidized forms of arsenic, heated in air, as in a glass tube open at both ends, oxidize to arsenious anhydride (455); and the latter substance sublimes in the tube, pro- ducing a white coating of microscopic octahedral crystals. 473. If dry arsenious anhydride is heated with dried sodium acetate, in the bulb of a small reduction-tube, arsen-dimethyl oxide, or cacodyl oxide, As2(CH3)40, is produced and recognized by its intensely offensive odor : As4Os + 8KC2H302 = 2As2(CH3)4O + 4K2C03 + 4COs 474. Arsenic is removed from mixture with metallic salts and non-volatile acids, and obtained in a concentrated form, by distilling the mixture with concentrated hydro- chloric acid—or sodium chloride and sulphuric acid—when arsenious chloride passes over at 132° C. (270° F.) and condenses with hydrochloric acid. A flask over a sand-bath, With a tube passing through the stopper and then inclined downwards to a small receiv- es flask set in a vessel of cold water, constitutes a sufficient apparatus. The distillate eay be examined according to 468. 475. Arsenious compounds are oxidized to arsenic compounds by a large number of oxidizing agents. As already stated (455 and 459 e,f), the sol- Vei'ts of elemental arsenic, and of arsenious sulphide, produce pentad ar- senic compounds. Among the oxidations of arsenious compounds most llsed in analysis are those by action of chlorine or bromine (a), iodine with sodium carbonate (b), nitric acid (c), copper sulphate with free fixed alkali (°0? and permanganates (e). «• H3AsQ3 + Cl 2 -)- H2O = H3AsQ4 + 2HCI b- 2H3As03 + 212 + SNa2CO3 = 2Na3AsQ4 4- 4NaI + 3H20 4- 5C02 c- 3H3As03 4- 2HN03 = 3H:iAsO4 4- 2NO 4- H2O d- H3AsQ3 4- 2CuSO4 4- 7KOH = K3AsQ4 + Ou2(0H)2 4- 2K2S°ved hy oxalic acid, before the first and second-group precipitations. 581. Estimation.—Gold is always weighed in the metallic state, to ;vbieh form it is reduced: (1) By ignition, if it contains no fixed acid; U) by adding to the solution some reducing agent, usually FeS04, H2C204, 'ydrate of chloral, or some easily oxidized metal, such as Zn, Cd, or Mg. Cold is also estimated volumetrically by H2C204 and the excess of H2C204 used, determined by K.,Mn,,Os. 156 Platinum. 582. Oxidation.—Gold is reduced to the metallic state by very many reducing agents, among which may be mentioned the following : Pb, Ag, Hg, (Hg2)", Sn, Sn", As, As"', AsH3, Sb, Sb", SbH3, Bi, Bi, Cu, (Cu2)", Pd, Pt, Te, Fe, Fe", Al, Co, Ni, (Cr2)vg Zn, Mg, H2C204, HNOs, P, H3P02, H3POs, PH3, H2S03, and a great number of organic substances. PLATINUM. Pt = 194.415. 583. Specific gravity, 21.50 (Deville and Debra y, 1875). Melting point, 1775° C. (3227° F.) (Violle, 1879), Valence, a dyad in Pt"o, and in all platinous compounds ; a tetrad in PtIVOa, and in all platinic com- pounds. 584. Occurrence.—Found in nature only in the metallic state, gene- rally alloyed with palladium, iridium, osmium, rhodium, ruthenium, etc. The Ural Mountains furnish the largest supply of platinum. 585. Properties.—Pure platinum is softer than silver, hut it generally contains iridium, which hardens it and increases its elasticity. In ductility it is surpassed only by gold and silver. It is a tin-white metal, nearly as lustrous as silver, fusible in the oxy-hydrogen blow-pipe flame, by which it can be vaporized. It is obtained as a soot-black powder—“ platinum- black”—by reduction from solution of alkali and alcohol ; and as a gray, porous, slightly coherent mass—“platinum-sponge”—by ignition of am- monio-platinic chloride. Both these bodies are remarkable for adhesive power ; and both, by strong compression, become compact, malleable, and lustrous, in the ordinary form of the metal. Platinum is not affected by air or water, at any temperature ; is not sensibly tarnished by hydrosul- phuric acid vapor or solution ; and is not attacked at any temperature by nitric acid, hydrochloric acid, or sulphuric acid, but dissolves in nitro-hy- drochloric acid, to platinic chloride, less readily than gold. 586. The Preservation of Platinum Vessels requires that it he remem- bered : (J) That free chlorine and bromine attack platinum at ordinary temperatures (forming platinic chloride, bromide); and free sulphur, phos- phorus, arsenic, selenium, and iodine, attack ignited platinum (forming platinous sulphide, platinic phosphide, platinum-arsenic alloy, platinic selenide, iodide). Hence, the fusion of sulphides, sulphates, and phos- phates, with reducing agents, is detrimental or fatal to platinum crucibles. The ignition of organic substances containing phosphates acts as free phos- phorus, in a slight degree. The heating of ferric chloride, and the fusion of bromides, and iodides, act to some extent on platinum. (2) The alkali hydroxides (not their carbonates) and the alkaline earths, especially baryta and lithia, with ignited platinum in the air, gradu- Platinum. 157 ally corrode platinum (by formation of platinites ; 3Pt + 2BaO -f 02 = SBaO.PtO). Nickel crucibles are recommended for fusion with alkali hy- droxides. (3) All metals which may he reduced in the fusion—especially com- pounds of lead, bismuth, tin, and other metals easily reduced and melted —and all metallic compounds with reducing agents (including even alkalies and earths) form fusible alloys with ignited platinum. Mercury, lead, bis- muth, tin, antimony, zinc, etc., are liable to be rapidly reduced, and im- mediately to melt away platinum in contact with them. (4) Silica with charcoal (by formation of silicide of platinum) corrodes ignited platinum, though very slowly. Therefore, platinum crucibles should not be supported on charcoal in the furnace, but in a bed of mag- nesia, in an outer crucible of clay. Over the flame, the best support is the triangle of platinum wire. (5) The tarnish of the gas-flame increases far more rapidly upon the already tarnished surface of platinum—going on to corrosion and cracking. The surface should be kept polished—preferably by gentle rubbing with moist sea-sand (the grains of which are perfectly rounded, and do not scratch the metal). Platinum surfaces are also cleansed by fusing horax upon them, and by digestion with nitric acid. 587. Platinum forms two oxides, platinous oxide, Pt"o, and platinic oxide, Pt "'O2 ■—both of which represent corresponding classes of oxy salts and haloid salts. The oxy- salts are instable.—None of the platinous compounds are permanently soluble in pure "'vater; the chloride is soluble in hydrochloric acid, the sulphate in water acidulated with sulphuric acid.—Platinic chloride (PtCl4) and bromide, all the platinicyanides (as pbPt(CN) 8), and the platinocyanides of the metals of the alkalies and alkaline earths (as K2Pt(ON)4), are soluble in water. The platinous and platinic nitrates are soluble in water, but easily decomposed by it, with the precipitation of basic salts. The larger number of the metalln-platijiic chlorides or “ chloroplatinates ” are soluble in water, including those with sodium [Na2PtClo or (NaCl)2PtCl4], barium, strontium, magnesium, zinc, alumi- nium, copper; and those with potassium, and ammonium, are sparingly soluble in water, ar>d owe their analytical importance as complete precipitates to their insolubility in alco- hol. Of the metallo-platinous chlorides (the “ chloroplatinites ”)—those with sodium, [Na2PtCl4], and barium, are soluble; zinc, potassium, and ammonium, sparingly solu- ble; lead and silver, insoluble in water. Platinic sulphate, Pt(SO4)2, is soluble in water. 588. Platinous chloride, PtCl2, is a greenish-brown powder, soluble in hydrochloric a°id without change, as a dark-brown solution, which remains platinous if protected b‘°in the air, but becomes platinic in contact with the air. The purely platinous chlo- ride solution is precipitated by potassium hydroxide and sodium hydroxide as platinous hydroxide, Pt(OH)2, dark-brown, soluble in excess, as alkali platinite, K2PtO2, etc,; b'orn which alkaline solutions, alcohol precipitates “platinum black ’’ (585).—Ammo- -11111111 hydroxide gives a green crystalline precipitate of platino-diammonium chloride, 2H6PtCl2, insoluble in cold water and in alcohol (compare 433). Hydrosulphuric a°id very slowly precipitates platinous sulphide, PtS, black, sparingly soluble in water, llot affected by acids, sparingly soluble by ammonium sulphide.—lodide of potassium slowly precipitates platinous iodide, Ptl2, red-brown (o black.—Oxalic acid produces no change; ferrous sulphate (slowly), and zinc (quickly), reduce the metal. 158 Palladium. 589. Platinic Chloride, PtCl4, is a brown-red solid, dissolving in water, or alcohol, as a reddish-yellow solution, permanent in the air.—Potassium hydroxide, and ammonium hydroxide, give, in solutions not very dilute, a yellow crystalline precipitate, as potassium platinic chloride, or potassium cbloroplatinate, K2PtCl6, etc., slightly soluble in water, insoluble in alco- hol, soluble in excess of the alkalies, and reprecipitated by hydrochloric acid. Chlorides of potassium and ammonium give the precipitate, the most nearly complete. Carbonates of potassium and ammonium form the same precipitates, insoluble in excess. Sodium carbonate gives no precipi- tate. Sodium hydroxide slowly precipitates, in moderately concentrated solutions, after warming, the (brownish-yellow) sodium platinate, Na2Pt03. —Potassium iodide colors the solution brown-red, and precipitates the black platinic iodide, Ptl4; with excess of the reagent, forming the spar- ingly soluble potassium platinic iodide, K2PtI6, brown. Sodium iodide the same.—Hydrosulphurie acid, and ammonium sulphide, slowly precipitate the black platinic sulphide, PtS„, slightly soluble in water, soluble by chlo- rine, and soluble in ammonium sulphide, as ammonium sulphoplatincite ; sodium sulphoplatinate is likewise soluble, reprecipitated by acids. Phos- phates form no precipitate. Reduction is not effected by oxalic acid (dis- tinction from gold); is slowly accomplished by ferrous sulphate, and rapidly by zinc ; also by chloral hydrate, and excess of alkali with heat (formiate); the reduced metal being, in each case, in black powder.—By the reducing blow-pipe flame, the compounds of platinum are reduced to spongy plati- num. 590. Estimation.—Platinum is in variably weighed in the metallic stale. It is brought to this condition : (1) By simple ignition ; (2) by precipita- tion as (NH.),PtCI.. K.PtCI., or PtS2 and ignition ; (3) by reduction, using Zn, Mg, or PeS04. 591. Oxidation.—Solutions of platinum are reduced to the metallic state by the following metals : Pb, Ag, Hg, Sn, (Sn" to Pt"), Bi, Cu, Cd, Zn, Pe, Ee", Co, and Hi. Very many organic substances reduce platinum compounds to the metallic state. PALLADIUM. Pd = 105.737. 592. Specific gravity, 11.4 (Deyille and Debray, 1859). Melting point, 1500° 0. (2732° F.) (Yiolle, 1878). Valence, a dyad in Pd'O and in all palladons compounds. A tetrad in Pd,v02, and in all palladic compounds. Found in platinum ores. 593. Properties.—Spongy palladium absorbs 935 times its volume of hydrogen, and still retains 600 volumes at 100° 0. At very high tempe- ratures the whole of the hydrogen is given off. The absorbed hydrogen acts in some respects like nascent hydrogen ; reducing mercuric chloride to Palladium. 159 metallic mercury, etc. Palladium is a white metal, more lustrous than platinum, with which it is classed, in accordance with its general properties. It is, however, a little more fusible and volatile (in the oxyhydrogen flame), and much more oxidizable than platinum. In the air, it is little tarnished at ordinary temperatures, but at a red heat it covers with oxide.—lt is slightly attacked by boiling hydrochloric or sulphuric acid ; dissolves in nitric acid, with formation of palladous nitrate, and if in the cold, with separation of nitrous acid, which remains in solution ; more readily in mtro-hydroehlorie acid, as palladic chloride, PdCl4. It is blackened by alcoholic solution of iodine (distinction from platinum). 594. Palladium forms one stable oxide, palladous, PdO, and two chlorides, pallad- ous, PdCl2, and palladic, PdCl4. The latter is the most stable of the palladic combina- tions, but is reduced to palladous chloride by boiling in water, and by dilution with much cold water. Palladous chloride is readily soluble in water with a brownish-red color ; with metal- lic chlorides, it forms double chlorides, as potassio-palladous chloride, K2PdCl4, all of which are soluble in water.—Palladous iodide is insoluble in water, alcohol, or ether ; insoluble in dilute hydrochloric acid or hydriodic acid ; slightly soluble by iodides and by chlorides.—Palladous nitrate, Pd(NOs)2, is soluble in water with free nitric acid ; the solution being decomposed by dilution, evaporation, or by standing, with precipitation of variable basic nitrates. Palladous sulphate, PdSC4, dissolves in water, but decom- poses in solution on standing. The instable palladic chloride, brown-black in solution, forms double chlorides with the metals—as calcium palladia chloride, CaPdCl6—these being mostly stable in water, and mostly soluble in water and in alcohol. The potassium palladic chloride (red), as an exception, is slightly soluble in water and insoluble in alcohol, but partially decomposed by both solvents. 595. Palladous Chloride is precipitated by potassium hydroxide or sodium hy- droxide ; as brown basic salt or as brown palladous hydrate, Pd(OH)2, soluble in ex- cess of the hot reagents. Ammonium hydroxide gives a flesh-red precipitate of pal- tadio-diammonium chloride, N2HcPdCI2. The flesh-red precipitate is soluble in excess °f the ammonia, and from this solution reprecipitated by hydrochloric acid, with a yel- low color. The fixed alkali carbonates precipitate the hydroxide ; ammonium carbon- ate acts like the hydroxide.—Hydrosulphuric acid and sulphides precipitate the dark- brown palladous sulphide, PdS, insoluble in the ammonium sulphides, soluble in nitro- bydrochlonc acid.—Potassium iodide precipitates palladous iodide, Pdl2, black, visible Hi 500,000 parts of the solution, with the slight solubilities stated in 594, an important separation of iodine from bromine. In very dilute solutions, only a color is produced, or the precipitate separates after warming. At a red heat, the precipitate is decomposed. Potassium cyanide precipitates palladous cyanide, Pd(CN)2. white, soluble in excess of the reagent.—Chloride of potassium precipitates, from highly concentrated solutions, the golden yellow, crystalline, potassium palladous chloride, K2PdOlO (594). Phos- phates give a brown precipitate. Palladous nitrate gives most of the above reactions ; no precipitate with ammonia, ar*d a less complete precipitate with iodides. 596. Palladium is reduced, in dark-colored precipitate, from all compounds in solu- tion, by sulphurous acid, stannous chloride, phosphorus, and all the metals which pre- cipitate silver. Ferrous sulphate reduces palladium from its nitrate, not from its chlo- 160 Molybdenum. ride. Alcohol, at boiling heat, reduces it ; oxalic acid does not (distinction from gold). —Nearly all its compounds are reduced by heat, before the blow-pipe, to a “ sponge.” If this be held in the inner flame of an alcohol-lamp, it absorbs carbon at a heat below redness ; if then removed from the flame, it glows vividly in the air, till the carbon is all burnt away (distinction from platinum). 597. Estimation.—(l) As metallic palladium, to which state it is re- duced by mercuric cyanide or potassium formate ; it is then ignited, first in the air and then in hydrogen gas. (2) As K2PdCl0. Evaporate the solu- tion of palladia chloride with potassium chloride and nitric acid to dryness, and treat the mass when cold with alcohol of .833 sp. gr., in which the double salt is insoluble. Collect on a weighed filter, dry at 100° C., and weigh. MOLYBDENUM. Mo = 95.537. 598. Specific gravity, 8.56 (Loughlin, 1868). Valence, a dyad in Mo 'O and in Mo'C12 ; a triad in Mo"203 and in Mo '2CI6 ; a tetrad in Moiv02 and in MoivC14 ; a hexad in MoVI0,( and in molybdates. 599. Occurrence.—Not found native. But occurs chiefly as molyb- denite, MoS2, as an oxide in molybdenum ochre, MoOa, and as wulfenite, PbMo04. The metal is reduced from its oxides by heating with carbon, or in a current of hydrogen. 600. Properties.—lt is a silver-white, hard and brittle metal, fusible at the highest furnace heat. It is not oxidized in the air at ordinary tem- peratures ; but when slowly heated, it gains a brownish-yellow, then a blue tarnish ; and at a higher heat, it burns to Mo03. It is oxidized by water vapor at a red heat. It is quickly dissolved by nitric acid, as molybdic anhydride (MoOa), with evolution of nitric oxide ; slowly by hot, strong sulphuric acid, with liberation of sulphurous anhydride. Molybdenum forms three classes of compounds, viz., molyhdous oxide (sometimes called hypomolybdous oxide), Mo 'O ; chloride, MoC12, and other molyhdous salts ; molybdic oxide (sometimes called molyhdous oxide), Mo""0„; chloride, MoC14, and corresponding salts; and molybdic anhydride, MoVIOs, which combines with bases, to form stable molybdates—also feebly unites with strong acids. Each of these classes includes stable salts ; the two bases are converted into molybdic acid or molybdates by strong oxidizing agents ; while molybdates are reducible to one or the other of the bases by deoxidiz- ing agents. 601. The molybdom salts are not generally very permanent in solution. The chlo- ride is soluble (in dilute hydrochloric acid?); the bromide and iodide decomposed in water to oxybromide and oxyiodide; the sulphate decomposed by water to a soluble and an insoluble salt ; the nitrate soluble in water. The solutions are dark-brown and opaque. M OL YBDENUM. Water dissolves the tetrachloride, bromide (yellow-brown solution), iodide (red solu- tion), nitrate and M0(504)2 (reddish-brown solutions). Molybdic anhydride is very sparingly soluble in water (800 parts). It does not form an acid in the solid state. The normal molybdates of the alkali metals (as K2Mo04) are soluble, of the remaining metals insoluble in water. The molybdenum trisulphate, M0(504)3, and molybdenum hexa-chloride, MoCI6, are soluble in water; the correspond- ing nitrate, M0(N03)6, soluble in dilute nitric acid. Molybdous salts, as M0(N03)2, with alkali hydroxide and carbonates, precipitate the dark-brown molybdous hydroxide, becoming blue in the air by oxidation to molybdic molybdate, Mo(MoD4)2 and M 0205. The hydroxide is insoluble in alkalies, sparingly soluble in alkali carbonates, readily soluble in alkali hydrogen carbonates. With hy- drosulphuric acid and sulphides, a brown precipitate of molybdous sulphide, MoS. solu- ble in ammonium sulphide. 602. Molybdic salts, as MoC14, with alkali hydroxides and sulphides, give reac- tions corresponding with molybdous salts, and likewise turn blue in the air, by formation of intermediate oxides. The precipitated hydroxide is reddish-brown ;it dissolves in al- kali carbonates by formation of alkali molybdates. Zinc precipitates molybdous hydrox- ide,, by reduction. 603. Molybdic Anhydride is a white powder, or is in needle-form crys- tals, turning yellow when hot and again white on cooling ; melting at a red and vaporizing at a white heat. It is soluble by acids, and by alkali- hydroxides, especially ammonia, in formation of molybdates. Water solutions of MOLYBDATES, with acids, precipitate molybdic anhydride, MoOa, white, soluble in excess of the acids.* Hydrosulphuric acid colors the molybdate in neutral or alkaline solutions, yellow to brown, without precipitation ; but from the acid solutions it precipitates the brown MoS„, sulphomolybdic acid, the supernatent liquid appearing blue. The precipitate is soluble in ammonium sulphide, better when hot and not too concentrated, as ammonium thiomolybdate, (NH4)2MoS4, from which acids I’eprecipitate MoSa. 604. Tribasic phosphoric acid and its salts precipitate, from strong nitric acid solutions of ammonium molybdate, somewhat slowly and on Warming, ammonium pliosphomolybdate, yellow, .of variable composition, soluble in ammonia and other alkalies, sparingly soluble in excess of the Phosphate. Hydrochloric acid may be used instead of nitric. The sodium Vbospbomo1 ybdatc is soluble in water, and precipitates ammonium from its salts ; also, it precipitates the alkaloids—for which reaction it has some ’mportance as a reagent.f * In making the solution of molybdate with nitric acid, used as a reagent, the slightly alkaline solution °f the molybdate is poured into the nitric acid, slowly, with stirring. + Sodium Phosphomolybdate—Sonnenschein’s reagent for acid solutions of alkaloids—is prepared as follows; The yellow precipitate formed on mixing acid solutions of ammonium molybdate and sodium Phosphate—the ammonium phosphomolybdate -is well washed, suspended in water, and heated with sodium carbonate until completely dissolved. The solution is evaporated to dryness, and the residue gently ignited all ammonia is expelled, sodium being substituted for ammonium. If blackening occurs, from reduction °I molybdenum, the residue is moistened with nitric acid, and heated again. It is then dissolved with water ®nd nitric acid to strong acidulation ; the solution being made ten parts to one part of residue. It must be hept from contact with vapor of ammonia, both during the preparation and when preserved for use. Molybdenum. Arsenic acid and arsenates give the same reaction ; ammonium arseno- molyhdate being formed. Recent molybdic anhydride, well washed (with alcohol and then water), on digestion with aqueous phosphoric acid, forms a lemon-yellow salt, insoluble in water, but slowly soluble in excess of the hot phosphoric acid, as acid permolybdic phosphate or phospho- molyhd-ic acid. It is soluble in alcohol as well as in water ; the solution giving, with ammonia and with alkaloids, the yellow precipitate of ammonium phosphomolybdate, etc. 605. The. alkaline solutions—normal molybdates of the alkali metals—give, with all non-alkali salts, precipitates of non-alkali molybdates, the latter being insoluble (or sparingly soluble) in water. The following are some of the reactions giving precipitates: K2MoQ4 + BaClo = BaMoO, -f 3KCI K2MoOi + Pb(NO3)2 = PbMoO, (wulfenite) -f 2KNOs K2MoOj -f- OuS04 CuMoo4 (yellow-green) -f- K2S04 606. Reducing agents convert molybdic acid either into the blue inter- mediate oxides, or, by further deoxidation, into the black molybdous oxide, MoO. In the (hydrochloric) acid solutions of molybdic acid, the blue or black oxide formed by reduction, will be held in solution with a blue or brown color. Nitric acidulation is, of course, incompatible with the reduc- tion. Certain reducing agents act as follows : Ferrous salts (in the hydrochloric acid solution) give the blue oxide solution. Cane sugar, in the feebly acid boiling solution, forms the blue color—seen better after dilution ; a delicate test. Stannous chloride forms first the blue, then the brown, or the greenish-brown to black-brown, solu- tion of both the intermediate oxide and the molybdous oxide. Zinc gives the blue, then green, then broicn color, by progressive reduction. Formic and oxalic acids do not react.—Dry molybdates, heated on platinum foil with concentrated sulphuric acid to vaporization of the latter, form, on cooling in the air, a blue mass—a very delicate test.* 607. With microcosrnic salt, in the outer blow-pipe flame, all compounds of molyb- denum give a bead which is greenish while hot, and colorless on cooling; in the inner flame, a clear green bead. With borax, in the outer flame, a bead, yellow while hot, and colorless on cooling; in the inner flame, a brown bead, opaque if strongly saturated (mo- lybdous oxide). On charcoal, in the outer flame, molybdic anhydride is vaporized as a white incrustation; in the inner flame (better with sodium carbonate), metallic molybde- num is obtained as a gray powder, separated from the mass by levigation. 608. Estimation.—(l) Molybdic anhydride and ammonium molybdate may be reduced to the dioxide by heating in a current of hydrogen gas. * A solution of 1 milligram of sodium (or ammonium) molybdate in 1 c.c. of concentrated sulphuric acid (about 1 part to 1,840 parts) is in use as Fecehdb’s Reagent for alkaloids. The molybdenum in this so- lution, which must be freshly prepared for use each time, is reduced by very many organic substances; and with a large number of alkaloids, it gives distinctive colors, blue, red, brown, and yellow. G ERMANIUM. —Nor we g ium. 163 The heat must not be permitted to rise above dull redness. Or the tem- perature may rise to a white heat, which reduces it to the metallic state, in which form it is weighed. (2) Lead acetate is added to the alkali molyb- date, the precipitate washed in hot water, and after ignition weighed as 3?bMo04. (3) Volumetrically. The molybdic acid is treated with zinc and HCI, which converts it into Mo2Cl6. This is converted into molybdic acid again by standard solution of potassium permanganate. GERMANIUM. Ge = 72.32. 609.—Specific gravity, 5.469 (Winkler, 1886). Melting point, 900° C. (1652° F.) (Winkler, 1886). Valence, a dyad in Ge''o, and in all ger- manous salts ; a tetrad in GeIV02, and in all germanic salts. Found in argyrodite. A white, lustrous metal, brittle and easily powdered. It dis- solves in H2S04, not in HCI. HI403 converts it into white Ge02. It forms two series of compounds, the germanous and the germanic. GeCl2 is formed by passing HCI over the heated metal, and is a colorless, fuming liquid, boiling at 72° C. GeCl4, formed by the direct union of its elements, is also a liquid boil- ing at 86° C. It fumes in the air, and is deconrposed by water. GeO is formed by treating GeCl2 with KOH, and heating the precipitate in a cur- rent of C02. Ge02 is formed when the metal burns in oxygen. It is white, and sparingly soluble in water, from which it may be crystallized. It acts as an oxide toward stronger bases. When GeCl4 is treated with H2S, GeS2 is precipitated. It is somewhat soluble in water, unless HCI is present. It dissolves in alkali sulphides. Heating it in hydrogen gas converts it mto GeS. GeS is soluble in KOH, which converts it into germanium and 2. For description of other germanium compounds see Jr. Prakt. Chem., 1886, (2), 34, 177. Also other articles in same journal by C. Winkler. NORWEGIUM. Ng = 218.93 ? 610.—Discovered by Dahll in a specimen of Norwegian nickel glance. specific gravity, 9.441. Melting point, 254° 0. Chemically it resembles bismuth. Its chloride is precipitated by water, forming an oxychloride. I'iie oxide (probably Ng203) is fusible, and gives, with the blow-pipe, a Metallic incrustation. But unlike bismuth, its hydroxide is soluble in both alkali hydroxides and alkali carbonates. Reactions of Certain Rare Metals. 164 611. Reactions of Ruthenium, Iridium, Rhodium, and Osmium. In solution as Rti2CI6, orange. (NaCl)2IrCl4, black. K3R1i(S04)3, rose-colored. H20sO4. H2S, in acid sol. Ztu«S3,brown, form- ed slowly, with blue sol. Ir2S3 -f S, brown. Kh2S3, brown, sol. in hot nitric acid. Pre. OsS4, black. (NH4)2S2, in ex- cess. Solution formed with difticuliy. Solution. Precipitate. Pre. OsS4, black. KOH or NaOH, in excess. Ruj(OH)6, black. Brovrn-black pre., turning blue. Solution, pre. by al- cohol. SnClj. Zn. (NaCl)8Ir2Cl6. Ir, black. Pre., metallic Ru. Eih, black. Os. kno2. Pre , sol. in excess, turned to dark red by (NH4!2S. Ir(OH)4 + KN02, boiled with S02, green pre. Pre., orange. Pre. K20s04. (1) (3) (3) (4) (1) Potassium thiocyanate, in absence of other platinum metals, slowly forms a red color, turning to violet when boiled. (3) The metal is insoluble in all acids. After fusion with sodium hydroxide and oxidizing agents, it dissolves by nitro-hydrochloric acid. (3) The metal is insoluble in all acids; but by fusion with potassium hydrogen sulphate, forms the solu- ble salt above taken. (4) The metal vaporizes at white heat, and burns to 0s04, an acidulous anhydride, having an irritating and offensive odor, and forming instable osmates of great oxidizing power. Osmates separate iodine from iodides, decolorize indigo solution, and, with sulphites, give a deep violet color or blue precipitate. 612. Reactions of Tellurium, Selenium, and Tungsten. In solution as KqTgOs. K2TeQ4. Na2Se03. Na2Se04. Na2WO,. H2S in acid sol. TeS2, brown, Na2WS4. ’ (NH4)2S2, excess. HC1. No pre. H2Te03, white pre. No pre. No pre. No pre. h2wo4, white pre., then yellow. Boiled, H2Te03, and Cl. Boiled, HoSe- 03, and Cl. SnCl2. Te, black. Se, red. Yellow pre., heated with HC1, blue. Fusion on Ch, ■with Na2C03. Na2Te, staining Agr, and with acids giving H2Te, having odor of H2S, and dissolving red. Na3 ' + 2H ,0 + 4NO HNO2 + hi = I + H.o + NO 710. Nitrites with very dilute acids—and with acetic acid—form the brown liquid with cold solutions of ferrous salts (distinction from nitrates, provided the reagent, PeS04, contains no free sulphuric or other free acid). Nitrites, with iodic acid, or ioclate and slight acidulation, give free iodine—a good distinction from nitrates. A concentrated solution of nitrites, treated with a drop or two of aniline sulphate solution, gives the vapor of phenol, recognized by its odor (COH7N oxidized to CbH6O) —a distinction from nitric acid. Indigo solution in sulphuric acid is bleached by nitrites. As a reducing agent, a nitrite decolors potassium permanganate solution acidulated with sulphuric acid—an easy distinction from nitrate. 711. Nitrites are all soluble in water—argentic nitrite being but very sparingly soluble, and nitrites generally requiring for solution a larger proportion ot water than nitrates. In solutions not very dilute, silver nitrate precipitates stiver nitrite, AgN02, white. In moderately concentrated solutions of potassium nitrite, cobalt nitrate precipi- tates pntassio cobaltic nitrite, reddish-yellow, sparingly soluble in water (370). 6Fe304 + 2H,504 + 2KNO2 = Fe2(S04)3 + K2S04 + 2(FeSO4)2NO + 2H20 Nitrous Acid. In analysis, nitrites respond to the common test for nitrates (724); from which they are distinguished as stated in 709 to 711. By ignition, metallic nitrites are resolved into metallic oxides, nitrogen, and oxygen; except silver and mercury nitrites, and the nitrites of some of the rarer metals which, like their oxides, ai’e reduced to the metallic state by heat alone. Ammonium nitrite is decomposed into nitrogen and water. Heated with oxidizable bodies, nitrites deflagrate or detonate, like nitrates. If cyanide of potassium be added to an alkaline nitrite, then some neutral solution of cobaltous chloride and a little acetic acid, the liquid becomes of a rosy-orange color from the formation of nitrocyanide of cobalt and potassium (C. D. Braun). 712. Production of nitrites.—The nitrites are all soluble; the silver and lead salts sparingly. Nitrite of potassium may be made by fusion from the nitrate, oxygen being evolved, or by passing peroxide of nitrogen, N204, into potassium hy- drate. Silver nitrite can be made from this by precipitation, and purified by recrysrallization. Basic lead nitrite can be made by boiling lead nitrate with metallic lead. The other nitrites are made by transposition ; adding to silver nitrite, the chloride of the metal, which we wish to change to a nitrite, care being taken to add just enough. The nitrites of mcrcurosura, tin, antimony, bis- muth, aluminium, iron, and chromium have not been made. 713. Oxidation.—Nitrous acid acts sometimes as an oxidizing and some- times as a reducing agent. When it oxidizes nitric oxide is formed. When it reduces nitric acid is produced. H3P02 h. Nitric oxide and phosphoric acid are formed. H2S c. Hydrosulphuric acid has no action upon nitrite of potassium, but on addition of acetic acid sulphur separates, and if the solu- tion is hot nitric oxide is formed, the fumes of peroxide of ni- trogen being clearly visible. H2S03 d. Forms sulphuric acid and nitric oxide [Gmelin-Kraut, Handhuch der Chemie, I. 2, 458]. HCIO3 (j. First forms peroxide of chlorine, C1204 [Millori], then hydro- chloric acid [ Tons saint; Gmelin-Kraut, Handhuch der Chemie, I. 2, 458]. Br h. Free bromine seems to have no action on nitrous acid. HI j. Forms iodine and nitric oxide. HI03 k. Forms iodine and nitric acid. HCNS I, Forms sulphuric and hydrocyanic acids, and nitric oxide Some- times traces of other cyanogen products are formed. H4Fe(CN)6 m. Forms first, ferricyanic acid, then nitroferricyanic acid. H6Fe2(CN)12 n. Forms nitroferricyanic acid. PbO„ 0. Forms Pb" and nitric acid f Gmelin-Kraut, Handhuch der Chemie, I. 2, 458]. Nitrogen Peroxide—Nitric Acid. 203 Hg p. Becomes Hg° [ Watts’ Dictionary, 4, 70 ; Wurtz’s Dictionnaire de Chimie, 1. 489 ; Gmelin-Kraut, Handbuch der (Jhemie, I. 2, 460], Mercuric salts are not reduced [Heppe, Chemisette Reaction en, 336]. Mn" + W r. Dorms Mn'' and nitric acid. Co ' s. Changes Co" to Co "' (270) [ Watts’ Dictionary, I. 1045]. ■Nit- Nitrites acidulated with phosphoric acid reduce Ni'" to Ni", even in the cold. CrVI u. Becomes (Cr2)VI [ Watts’ Dictionary, 4, 70]. Urea v. On warming the nitrogen of both compounds is set free. CH4N20 + 3HN02 = C02 + 2N2 + BH2O NITROGEN PEROXIDE. NOa. Oxidation valence NivO~"2 714. Vapor density at 140° C. (H=l) is 23 (Deyille and Troost), showing that at that temperature the molecule is NC2. At 60.2° C. the density seems to indicate that the gas is made up of a mixture of equal parts of NOa and N2C4. Production.—(l) By submitting one volume of oxygen and two parts of NO to a freezing mixture, NO -f O = N(D2. (2) By ignition of Pb(NC3)2. 2Pb(NO3)2 = 2PbO + 4NO,2 + 02. The NOa is separated from the oxygen by passing the mixed gases through a freezing mixture. (3) By adding oxygen to the oxides evolved by the action of nitric acid on metals and other reducing agents. It is formed when NO comes in contact with the air. Properties.—A brownish-red gas, very irritating and poisonous when inhaled. Does not support ordinary combustion. It solidities at -9° C. (Peligot) ; and boils at 21.6° C. (Thorpe, 1880). At 0° C., water forms HNO, and UNO,,. At higher temperatures, HN03 and NO. Potassium hydroxide forms a nitrate and a nitrite. 2N0.. + HoO = HN03 + HNO2 3NOo + HoO = 2HNOj + NO 2N02 + 2KOH = KN03 + KN02 + HaO NITRIC ACID. HN03. Oxidation valence H'NV 0~"t _ O ii Structural valence H-0-N=o 715. Preparation of nitric anhydride, N205.—(1) Bypassing dry chlo- rine over AgN03 at 93.5° C.. condensing the gas in a cold receiver (De- yille). (2) By passing N02CI over AgNOa (Odet and Vignoh). (3) By treating HN03 witli P205 (Weber). 204 Nitric A cid. 1. 4AgNOa + 2C12 = 4AgCI -f 2N205 + 02 2. AgNOs -f N02CI = AgCl + N205 3. 6HNO3 + P2Oa = 2H3P04 + 3N205 A colorless solid, melting at 30° C., and boiling at 45° to 50° 0. (De- ville). It is instable, decomposed by beat above its boiling point. It unites with water to form the corresponding acid. 716. Absolute nitric acid—HNv o3—is a colorless, transparent, mobile liquid, of the specific gravity of 1.58 at 15° C. [Millon], boiling at 86° C. (187° F.) with partial decom- position, leaving nitric acid mixed with water. Aqueous nitric acid having 70 per cent, of HNO3, and corresponding to (HN03)2(H20)3, specific gravity 1.42, appears to be a definite hydrate, as both stronger and weaker acids are, by boiling, reduced to this composition, which boils at 128° C. (258° F.) The reagent designated in this work as nitric acid has a specific gravity of 1.2, and about 35 per cent, of HNOs (Fresenius’ standard). 717. By heating, by action of the light, and by organic particles from the air, strong nitric acid parts with oxygen and generates nitrous anhydride and nitric peroxide, N2Os and N204, which remain dissolved with a yellow color. The tendency to this change is very strong in absolute nitric acid, which cannot well be preserved colorless ; and the acid of 70 per cent, colors far more readily than that of 35 per cent. The nitrogen ox- ides may be expelled by boiling ; or, with less waste of nitric acid, by passing pure air through it, by means of a bellows, a wash-bottle, and, to avoid dilution, a drying- tube. Nitric acid is a strong oxidizing agent, and, as such, its reactions with oxidizable elements and compounds are in constant requisition in analysis. Unless heated, nitric acid does not generally oxidize substances as quickly as chlorine with water. 718. In oxidizing and dissolving metals or metalloids, and in oxidizing lower oxides, nitric acid most frequently disengages water and nitric oxide (d); but, with cer- tain substances and under certain conditions, other residues are chiefly produced, as dinitrogen tetroxide (6), nitrous acid (c), nitrous oxide (e), nitrogen (/), hydrogen (n), ammonium nitrate {g). Examples of several of these results, as varied by conditions, are seen in the case of zinc (290 c, d, e), iron (202 b, c, d, e), tin (526 d, e), arsenious acid (475 c). Further, in the study of oxidations, see 733. a. HNC3 = NO3 (combined) + H b. 2HNOs = O “ + H2O + N204 c. HNOa = O “ + HNO,(2HNO, = H2O and N203) d. 2HN03 = BO “ + H2O + 2NO e. 2HN03 = 40 “ + H2O + N.O /. 2HNO3 = 50 “ + H..0 + 2N (combined). g. IOHNOa = BNO3 “ + BH2O + NH4N03 719. The metallic sulphides (except mercuric sulphide) dissolve as nitrates by action of nitric acid, more or less readily ; the sulphur being at first mostly left as a residue. But as fast as the sulphur is oxidized, metallic sulphates are formed, soluble or insoluble (equations in 388). 720. Nitric acid is formed by transposition between sulphuric acid and nitrates : kno3 + h2so4 - hno3 + khso4 (With solid nitrates, short of a higli heat.) 2KNO3 + H2SO. = 2HNO3 + k2so4 (In solution ; or with solids at a red he.-it.) Nitric Acid. 721. The Nitrates are all soluble in water. There are a few basic ni- trates—basic bismuth nitrate, basic mercurous and mercuric nitrates—insol- uble in water. Many of the nitrates are insoluble in alcohol. Most of the tests for the identification of nitric acid are made by its deoxidation, disengaging a lower oxide of nitrogen (724), or even, by com- plete deoxidation, forming ammonia (726). 722. Sulphuric acid is transposed with metallic nitrates, with but little decomposition of the nitric acid formed (720), The colorless or slightly reddish gas does not rise till the mixture is very hot—absolute nitric acid not being, like hydrochloric acid, a gas at ordinary temperatures. It red- dens litmus, and has a characteristic acrid odor. 723. If, with the sulphuric acid, a bit of copper turning, or a crystal of ferrous sulphate, is added to a concentrated solution or residue of nitrate, the mixture gives off abundant brown vapors; the colorless nitric oxide, NO, which is set free from the mixture, oxidizing immediately in the air to dinitrogen trioxide and peroxide, N2Oa and N02 : 2KN03 + 4H2SQ4 + BCu = K..504 + 30u504 + 4H20 + 2NO 2KN03 + 4H2504 + CFeS04 = K2SC4 + eFe2(S04)3 + 4H2C + 2NO 2NO + O = N203 2ND + 20 = N204 The three atoms of oxygen furnished by two molecules of nitrate (as in 718 d), suffice to oxidize three atoms of copper; so that 3CuO with 3H2S04, may form 3CuS04 and 3H20. The same three atoms of oxygen (having- six bonds) suffice to oxidize six molecules of ferrous salt into three mole- cules of ferric salt; so that 6FeS04 with 3H2S04, can form 3Fe2(S04), and 3H„O. 724. Now if, by the last-named reaction, the nitric oxide is disengaged in cold solution, with excess of ferrous salt and of sulphuric acid, instead of passing off, the nitric oxide combines with the ferrous salt, forming a black brown liquid, (FeSO4)2NO, decomposed by heat and otherwise in- stable. And 2NO require 4FeS04, in addition to the proportion of ferrous salt in the equation in 723. 2KNO3 + 4H2S04 + 10FeSO4 = K2S04 + 3Fe2(S04)3 + 4H2Q + 3(FeSO4)2NO This exceedingly delicate test for nitric acid or nitrates in solution may be conducted as follows : Take sulphuric acid to a quarter of an inch in depth in the test-tube ; add without shaking a nearly equal bulk of solu- tion of ferrous sulphate, and cool the liquid; then add slowly of the solu- tion to be tested for nitric acid, slightly tapping the test-tube on the side, but not shaking it. The “ brown ring ” forms, between the two layers of liquid—violet, red, brown, or black, according to proportions and condi- tions. The color disappears on heating, with evolution of nitrous gas, yel- lowish ferric solution remaining. The test is somewhat more delicate if a. crystal of ferrous sulphate be added, instead of the solution, and the test- tube be set aside for several hours. 206 Nitric Acid. 725. Slight traces of nitrate (as in rain or river-waters) are detected, according to the above reaction, by first reducing to nitrite by heating for some time with zinc amalgam, or less readily with finely divided zinc. Nitrites previously found to be absent, by the same test, viz.: To a thick layer of the clear filtered water the solution of ferrous sul- phate is added, and the brown coloration obtained, if nitrites have been formed. Or, a drop or two of potassium iodide solution with fresh starch-paste, and a drop or two of very dilute sulphuric acid, is added (709). The reduction to nitrite may also be effected by zinc or cadmium, in acidulated so- lution, as follows (Stoker):* Boil the solution, slightly acidulated (by addition of sul- phuric acid, if necessary) with metallic cadmium (or zinc), for about five minutes, in a tall vessel—or, better, in a retort with raised condenser—and filter or decant from the metal. Then add a mixture of potassium iodide and starch-paste—or, better, a mixture of zinc iodide, zinc chloride, and starch-paste. The iodine-color indicates nitrous acid, reduced from nitric acid. Without the boiling in acidulate solution, hydrogen peroxide may be formed, giving a fallacious indication for nitric acid. 726. Reduction to Ammonia, by strong reducing agents (718 g), is a valuable re- source in identifying nitric acid. The tests based on this principle are delicate, but do not distinguish nitric acid from nitrous acid or cyanogen compounds. Ammonia, if found already present, may be distilled off. In those tests requiring use of strong al- kali, nitrogenous organic substances will give ammonia. The ammonia obtained is, in most methods, identified by potassium mercuric iodide—a reaction so delicate as to show the frequent presence of ammonia in distilled water and many reagents. Hence, all these must be first tested, if necessary, after distillation. Larger quantities of am- monia are recognized in vapor, by litmus, etc. In neutral solutions, sodium amalgam is used as follows—a method for the total of nitric and nitrous acids and ammonia, in potable waters, and otherwise applied: To 6 or 8 cub. cent. (or 2 fluid drachms) of the solution (in a carefully cleaned test-tube) add 100 to 200 grains of sodium-amalgam which is % per cent, sodium ; cork the tube lightly and leave for twelve hours. Hydrogen is always slowly evolved, and escapes. Decant and rinse into a glass cylinder one inch wide, and at least six inches high, and add water to about 60 cub. cent. (2 fluid ounces). Nessler’s test solution is now applied. The test cannot be made before decantation from the amalgam, as the nascent hydrogen inter- feres. The nascent hydrogen developed by dissolving zinc in solution of potassium hy- droxide also reduces nitrates in the alkaline solution, and evolves ammonia. This is a convenient and efficient test by reduction to ammonia. The solution should be strongly charged with potassium hydrate, the zinc finely divided, and mixed with half its weight of iron filings. The mixture is then distilled at a boiling rate, and the distillate tested for ammonia by potassium mercuric iodide. The reagents, including the water, should be first tested in the apparatus. Metallic magnesium may be used for the reduction, as follows: Acidulate with phosphoric acid; add magnesium wire, and leave, cold, a few minutes. Then test for ammonia, by the potassium mercuric iodide solution with potassium hydroxide. If in- terfering acids are present, add potassium hydroxide, distill, and test the distillate for ammonia. Stannous chloride and hydrochloric acid, heated with a nitrate, form stannic chlo- ride, and convert nitric acid to ammonia (which remains as ammonium salt). 727. With hydrochloric acid, nitric acid forms free chlorine, etc. (nitrohydrochloric acid), applied as a test for nitric acid—in absence of other oxidizing agents—as follows : Heat a little hydrochloric acid in a test-tube to boiling; color it (slightly) with a drop or * Amer. Jour. Science, [3] XII. 176 (•Sept., 1876). Nitric A cm. two of very dilute indigo solution (in sulphuric acid), and boil again. If the hydrochlo- ric acid was pure, the color remains unchanged. The addition of a nitrate, with boiling, now quickly bleaches the solution. 728. Phenol, ChHsOH, gives a deep red-brown color with nitric acid, by formation of nitrophenol (mono, di, or tri), C6Ht(NOs)OH to CcH2(NOL.)sOH, “picric acid” or nitrophenic acid. A mixture of one part of phenol (cryst. carbolic acid), four parts of strong sulphuric acid, and two parts of water, constitutes a reagent for a very delicate test for nitrates (or nitrites), a few drops being sufficient. With unmixed nitrates, the action is explosive, unless upon very small quantities. The addition of potassium hy- droxide deepens and brightens the color of the nitrophenic acid solution. 729. Brucia, dissolved in concentrated sulphuric acid, treated (on a porcelain sur- face) with even traces of nitrates, gives a tine deep-red color, soon paling to reddish-yel- low. If now stannous chloride dilute solution be added, a fine red-violet color appears. (Chloric acid gives the same reaction) Aniline Sulphate solution, with a half volume of concentrated sulphuric acid, treated (on a porcelain surface) with traces of nitrates, gives a rose-red color, commencing with red lines, and when concentrated appearing brown-red. 730. By slight ignition, nitrates of the fixed alkali and alkaline earth metals are reduced to nitrites, recognized as shown in 709. Stronger ignition changes them to caus- tic bases, with formation of brown vapors. Nitrates of heavy metals are mostly changed to oxides by heat; ammonium nitrate, wholly to nitrous oxide and water. Heated on Charcoal, or with potassium cyanide, or sugar, sulphur, or other easily oxidizable substance (as in gunpowder), nitrates are reduced with deflagration or explo- sion, more or less violent. With potassium cyanide, on platinum foil, the deflagration is especially vivid. In this reaction free nitrogen is evolved, as by equation f, 718. Strongly heated with excess of potassium hydrate and sugar or other carbonaceous compound, in a dry mixture, nitrates are reduced to ammonia, which is evolved, and may be detected. In this carbonaceous mixture, the nitrogen of nitrates reacts with alkalies, like the unoxidized nitrogen in carbonaceous compounds (compare 705, 718 g, and 736) 731. Free Nitric Acid may be distinguished from nitrates, by giving the brown liquid with ferrous salt, on reduction by zinc, without addition of sulphuric acid, as stated in 735, and by coloring quill-cuttings or white woollen fabrics yellow when the solution is evaporated with them on the water-bath. The yellow color substance contains xanthoproteic acid, and is formed by action of nitric acid upon any gelatinoid substance —as the skin—and upon ordinary alburnenoid substances. 732. Nitrates are all soluble, a. They are formed by the action of nitric acid upon metals, as described in 733. h. By dissolving the oxides, hydrox- ides, or carbonates of the metals in nitric acid. Mercurous, stannous, manganous, and ferrous nitrates should, not be evaporated to expel excess of nitric acid, since a higher metallic form would result. The crystals may be washed in cold water, to free them from the uncombined acid. All nitrates are decomposed by heat; a few, the alkalies and alkaline earths, first evolve oxygen and form nitrites, afterward a mixture of oxygen and nitrogen, leaving the oxides; others either free nitric acid or a mixture of the oxides of nitrogen, until only the oxide of the metal remains. There are two exceptions, silver and mercury, in which cases only the free metals remain. Nitric Acid. 733. Of course nitric acid can never act as a reducing agent. Acting as an oxidizing agent, it may form NHa, W, W2O, WO, W2Oa, or W204. If the nitric acid is in excess, and a boiling heat be used, the product is nearly all WO, while excess of the reducing agent and a low temperature favor the formation of WH3. A convenient test for WO is made by passing the gas into ferrous sulphate. (See 718 and Acworth and Armstrong, Jour. Chem. Society, 32, 54.) The action upon H2C204 and HOWS has been given (659, 702). H3P02 b. Becomes phosphoric acid. H2S c. Forms sulphur, and may be further oxidized to sulphuric acid. The nitric acid must be stronger than sp. gr. 1.18 [Gmeliu- Krauf, Handbuch, 1 2, 219]. H2S03 d. Becomes sulphuric acid. HCI e. Forms nitrohydrochloric acid. HCI03 /. Nitric acid added to a chlorate liberates chloric acid, which de- composes, but the nitric is not changed [Penny, Jour. PraJct. Chem., 23, 296]. HBr g. Gives free bromine. I h. Forms iodic acid; action slow, and strong nitric acid should be used— at least sp. gr. 1.42. This is a practical method for making iodic acid, if acid of sp gr. 1.48 is used [Bousson, Comptes Rendus Academic des Sciences, 13, 1111]. HI i. Forms first free iodine, then iodic acid. HOWS j. Forms hydrocyanic and sulphuric acids. With strong nitric acid, traces of carbonic acid are formed [ Vogel, Gmelin's Hand-book, 8, 75]. H4Fe(CW)6 h. Forms ferricyanic acid [ Watts’ Dictionary, 2, 250], and then nitroferricyanic acid. HBFe2(CW)12 I. Forms nitroferricyanic acid. m. Nitric acid oxidizes all ordinary metals. (It does not act upon gold or platinum.) It forms nitrates, except in the case of tin, antimony, and arsenic, with which it forms HIOSn5O16, Sb205, and H3As04. With the respective metals it forms Hg' or Hg", Sn" or Sn"", As"' or Asv, Sb'" or Sbv, Fe" or Fe'", according to the amount of nitric acid employed. With copper it forms cupric nitrate (never cuprous); with cobalt it forms cobaltous nitrate. Pb304 n. Becomes plumbic nitrate and Pb02. The nitric acid is not re- duced. Hg' o. Becomes Hg". SnO p. Becomes Sn02. (Not stannic nitrate.) Oxygen. 209 H3As03 q Becomes H3As04. Sb.,03 r. Becomes Sb2Oa. Cu' s. Becomes Cu", t. MnO, NiO, CoO, 0r2O3 are not oxidized. Fe" u. Becomes Fe'". OXYGEN. 0 = 15.9633. 734. The vapor density (H = 1), 16, shows that the molecule of oxygen has two atoms. Occurrence.—The rocks, clay, and sand constituting the main part of the earth’s crust contain from 44 to 48 per cent, of oxygen ; and as water contains 88.87 per cent., it has been estimated that one-half of the crust is oxygen. Except in atmospheric air, it is always found combined. Preparation.—(l) By igniting HgO. (2) By heating KCI03 to 350° C., KCI04 is produced and oxygen is evolved ; at a higher temperature the KCI04 becoming KCI. But in presence of Mn(D2 the chlorate is com- pletely changed to KCI at 200° C., without forming KCI04, the Mn02 not being changed. Spongy platinum may be substituted for Mn02. (3) Ac- tion of beat on similar salts furnishes oxygen—e.g., KCIO and KCI02 form- ing KCI, KBr03 forming KBr, KI03 and KI04 forming KI, KN03 forming KNO, (at a white heat forming K2O, NO, and O). (4) By action of heat on metallic oxides, as shown in the equations below. (5) BaO, heated in the air to 550° C., becomes Ba02, and at 800° 0. is resolved into BaO and free oxygen, making theoretically a cheap process. (6) By heating higher oxides or their salts with H2S04 : CrVI is changed to (Cr2)yi, (Co 2)VI to Co”, (Ni2)VI to Ni”, Biv to Bi'”, FeVI to (Fe2)VI, PbIV to Pb”. Mn" +» to Mn” ;in each case a sulphate is formed and oxygen given off. (7) By passing sul- phuric acid over red hot bricks ; the S02 is separated by water, and after conversion into H2S04 is used over again. (8) The following cheap process is now employed on a large scale. Steam is passed over sodium manganate at a dull red heat; Mn2Q3 and oxygen are formed. Then, without change of apparatus or temperature, air instead of steam is passed over the mixture of Mn203 and NaOH. The Mn2Oa is thus again oxidized to Na2Mn04, and free nitrogen is liberated. 1. 3HgO (at 500° C) = 2Hg + 02 2. a. 3KCI03 (at 350° C.) = KCIO4 + KCI + 02 b. 2KCI03 (at red heat) = 2KCI + 302 c 2KCIO-, -I- JtMnO. (at 300° C.) = »Mno3 + 3KCI + 30* 3. a. KCIO2 + ignition = KCI -f- 02 b. SKBrOa 4- “ = 3KBr + 3Q2 c. 2KI03 + “ = 2KI + 302 d. KI04 + “ = KI + 302 e. 3KNO3 + “ = 3KNO2 + 02 210 Ozone. f. 4KNO.J (white heat) 2K,G + 4NO + CL 4. a. 3Pb304 “ = 6PbO + O, h. 3Sb205 (red heat) = 3Sb204 4- 02 c. Bi2Oa 11 = Bi203 -f- 02 d. 40r03 (about 200° C) 2Cr203 4- 302 e. 4K2Cr207 (red heat) = 2Cr203 + 4K2Or04 4- 302 /. 6Fe203 (white heat) = 4Fe304 4- 02 c/. 3MnOa 6i Mn304 -f- 02 h. 6Co203 (dull red heat) = 4C0304 + 02 fi. 3Ni203 u 4NiO -j- 02 j. 3Ag20 (300° C.) = 4Ag 4- 02 5. 3Ba02 (800° C.; = 2BaO +O2 6. 3K2Cr207 4- 8H2504 = 3K2Cr2(S04)4 + 3Q2 + BH2Q 3K2Mn208 4- 6H2S04 = 3K2504 4- 4MnSO4 4- 502 4- OH2O 2Pb304 + 6H2S04 = OPbSO4 4- 0H2O 4- 02 7. 3HnSO4 (upon red hot bricks) = 3S02 4- 3H20 4- 02 8. 4Na.MnO4 4- 4H20 (dull red heat) = BNaOH 4- 3Mn203 4- 302 SNaOH 4- 3Mn203 + air, 3(Q2 4- 4N2) = 4Na2MnO4 4- 4H2Q 4- 13N2 735. Properties.—By pressure and cold it is condensed to a colorless liquid. This liquid boils under ordinary atmospheric pressure at -184° C.; and under a pressure of 50 atmospheres at -11-3° C. (Weoblewski, 1884, Compt. Rend. 98, 983). Pure oxygen may be breathed for a short time without injury. A rabbit placed in pure oxygen at 34° 0. (75° F.) lived for three weeks, eating voraciously all the time, but nevertheless becoming thin. The action of oxygen at 7.3° C. (45° F.) is to produce narcotism and eventually death. When oxygen is cooled by a freezing mixture it induces so intense a narcotism that operations may be performed under its influ- ence. Compressed oxygen is “ the most fearful poison known.” The pure gas at a pressure of 3.5 atmospheres, or air at a pressure of 33 atmospheres, produces violent convulsions, simulating those of strychnia poisoning, ulti- mately causing death. The arterial blood in these cases is found to contain about twice the quantity of its normal oxygen. Further, compressed oxy- gen stops fermentation, and permanently destroys the power of yeast (Paul Bert). At varying temperatures oxygen combines directly with all metals ex- cept silver, gold, and platinum, and with these it may be made to combine by precipitation. It combines with all non-metals except fluorine ; the combination occurring directly, at high temperatures, except with Cl, Br, and I, which require the intervention of a third body. OZONE. 03 (P). 736. Ozone is the name of a modified form of oxygen, of the true na- ture of which there is still some doubt, as it has never been obtained free from ordinary oxygen. It is found in small quantities in the air, especially Hydrogen Peroxide. 211 at the sea-shore and in the open country. It is never found in the atmos- phere of crowded cities. Production—(1) By the action of electricity on moist oxygen or air, better at a low temperature. (2) By the electrolysis of dilute sulphuric acid. (3) By the combustion of phosphorus in moist air. (4) Formed during the evaporation of water. Properties.—When inhaled in small quantities, supposed to be health- ful. In larger quantities irritates the eyes and nose. Has proved fatal to animals that have been made to breathe it. Condenses to a liquid at a higher temperature than ordinary oxygen. Theory represents the mole- cule as having three atoms, Oa, ordinary oxygen two atoms, 02. When heated to 300° C. it is decomposed, forming ordinary oxygen, with an in- crease in volume ; perhaps thus, 203 = 302. It liberates iodine from KI, probably thus, 2KI -f- H2O +O3 = 2KOH +I2 + Oa. It is an energetic oxidizing agent, oxidizing where ordinary oxygen does not. HYDROGEN PEROXIDE. H2Q2. 737. This remarkable substance is usually made by treating the peroxide of the alkalies or alkaline earths with dilute acids. Either the oxyacids or hydracids will answer ; even carbonic will complete the result. It is gene- rally manufactured by treating Ba02 with C02, HCI, H2S04, or HFI ; on a large scale the latter is more frequently used. Ba02 + CO2 + H2O 83.C03 + H202 BaQ2 + 2HF = BaF2 + H202 Properties.—The pure anhydrous H202 is a transparent, colorless liquid, sp. gr. 1.4. Does not freeze at -30° C. Volatilizes slowly in a va- cuum without decomposition, if the temperature is not allowed to rise above 15° 0., otherwise explosion occurs. Does not redden but slowly bleaches litmus paper. Bleaches most organic colors, but more slowly than Cl or S02. It is odorless, and tastes bitter. The concentrated H202, placed upon the skin, soon produces a blister. It decomposes slowly, at a low tempera- ture, rapidly at a high temperature, into oxygen and water. The dilute so- lution in water is more stable. It is soluble in ether, making solution more stable than in water. It may be removed from the aqueous solution by shaking with ether. The acidified aqueous solution is more stable than the neutral or the alkaline. The best test for H202 is to add K2Cr207 and a little dilute HoS04. The chromium is oxidized to perchromic acid, H2Cr2Oa ; on adding ether and shaking, it is all absorbed by the ether, and, rising to the surface, is recognized by its intense blue color. Perchromic acid is instable in water, decomposing into oxygen and chromic salt. Its ethereal solution is permanent. K2Cr207 + H2S04 + H202 = K2S04 -f- H2O + H2Or2Og H2Or2Os + 3H2504 = Cr2(SO4)3 + 4H20 + 202 212 Fluorine. It generally oxidizes. Thus it changes Pb" to PbIV, (Hg2)" to Hg", Sn" to SnIV, As'" to Asv, Bi'" to BY, H3P02 to H3PQ4, H2S to S, H2S03 to H2S04, HI to I, etc. H2Oo + SKI = 2KOH + I 2 H3As03 -f H202 H3As04 + H2O H3P02 + 2H202 = H3P04 + 2H20 Yet it sometimes strangely acts as a reducing agent; and when it does so, ordinary oxygen is evolved from both substances. It reduces K2Mn2Oa to a manganous salt in acid mixture ; but in alkaline mixture to Mn02. It reduces gold, silver, and platinum from their oxides. K2Mn208 + 3H2504 + 5H202 = K2S04 + 2MnSO4 + BH2O + 502 Ag2o -f H202 2Ag + H2O -f* O2 It is estimated volumetrically by measuring the oxygen evolved on adding K2Mn2Og. Uses—lt has been used in medicine, externally as a lotion, and inter- nally for diabetes and oxaluria. It constitutes the golden hair-dye of the shops, the dyeing being really a bleaching (oxidizing) action. It is also used for bleaching and cleaning oil paintings and stained engravings (con- verting PbS into PbSOj. FLUORINE. F = 18.984. 738. Since Davy’s experiments in 1813, many others have attempted the isolation of fluorine. In his zeal the unfortunate Louyet fell a victim to the poisonous fumes which he inhaled, Faraday, Gore, Fremy, and others took up the problem in succession, but it was not ultimately solved until H. Moissan, in 1886, produced a gas which the chemical section of the French Academy of Sciences decided to be fluorine. Many ingenious ex- periments had been made in order to obtain fluorine in a separate state, but it was found that it invariably combined with some portion of the material of the vessel in which the operation was conducted. The most successful of the early attempts to isolate fluorine appears to have been made, at the suggestion of Davy, in a vessel of fluor-spar itself, which could not, of course, be supposed to be in any way affected by it. Moissan’s method was as follows : The hydrofluoric acid having been very carefully obtained pure, a little potassium hydrofluoride was dissolved in it to improve its conduct- ing power, and it was subjected to the action of the electric current in a U-tube of platinum, down the limbs of which the electrodes were inserted ; the negative electrode was a rod of platinum, and the positive was made of an alloy of platinum with 10 per cent, of iridium. The U-tube was pro- vided with stoppers of fluor-spar, and platinum delivery tubes for the gas- es, and was cooled to -23° 0. The gaseous fluorine, which was extricated at the positive electrode, was colorless, and possessed the properties of chlo- rine, but much more strongly marked. It decomposed water immediately. H YDR OFL UOIIIC A CID. 213 seizing upon its hydrogen, and liberating oxygen in the ozonized condition ; it exploded with hydrogen, even in the dark, and combined, with combus- tion, with most metals and non-metals, even with boron and silicon in their crystallized modifications. As, Sb, S, I, alcohol, ether, benzene, tereben- thene, and oetrolenm took fire in the g-as. Carbon was not attacked by it. (H. Moissan, 1886 ; Compt. Rend., 103, 202, and 256 ; Jour. Chem. Soc., 50, 849, and 976.) 739. Fluorine, in several characteristics, appears as the first member of the Chlorine Series of Elements. It cannot be preserved in the elemental state, as it- combines with the materials of vessels (except fluor-spar), and instantly decomposes water, forming hydrofluoric acid, HF, an acid prepared by acting on calcium fluoride with sulphuric acid (a). Fluorine also combines with silicon as SiF4. silicon fluoride, a gaseous com- pound, prepared by acting on calcium fluoride and silicic anhydride with sulphuric acid (b). On passing silicon fluoride into water, a part of it is transposed with the water, forming silicic acid and hydrofluoric acid (c); but this hydrofluoric acid does not at all remain free, but combines with the other part of the fluoride of silicon, as hydrojluosilicic acid, (HF)2SiF4, or H2SiF6 {d). This acid is used as a reagent ; forming metallic sihco- fluorides, soluble and insoluble (742). a. CaF2 + H2S04 = CaSO4 4- 3HF b. 2CaF2 + SiC2 + 2H2SQ4 = 2CaSO, + 2H2C + SiP4 c. SiF4 + 2H20 = Sio2 + 4HF (not remaining free) d. 2HF + SiF4 = H2SiF6 candd. 3SiF4 + 2HsO = SiC2 + 2H2SiF6 HYDEOFLUOHIO ACID. 740. A colorless, intensely corrosive gas, soluble in water to a liquid that reddens litmus, rapidly corrodes glass, porcelain, and the metals, except platinum and gold (lead but slightly). Both the solution and its vapor act on the flesh as an insidious and viru- caustic giving little warning, and causing obstinate ulcers, the anhydtous acid at 25° C. (77° F.) has a vapor density (H = 1) of 20, indicating that the molecule at this temperature is H2F2. But at 100° C. it is only 10, indicating that at that temperature the molecule is HF. The anhydrous liquid acid boils at 19.44° C. (Gore, 18G9), and does not solidify at -84.5° C. The Fluorides of the alkali metals are freely soluble in water, the solutions alka- line to litmus and slightly corrosive to glass: the fluorides of the alkaline earth metals are insoluble in water; of copper, lead, zinc, and ferrieum, sparingly soluble ; of silver and mercury readily soluble. Fluorides are identified by the action of the acid upon glass. Calcium chloride solution forms, in solution of fluorides or of hydrofluoric acid, a gelatinous and transparent precipitate of calcium fluoride, CaF2, slightly soluble in cold hydrochloric or nitric acid and in ammonium chloride solution. Barium chloride pre- cipitates, from free hydrofluoric acid less perfectly than from fluorides, the voluminous, white, barium fluoride, BaF2. Silver nitrate gives no precipitate. Sulphuric acid transposes with fluorides, forming hydrofluoric acid, HF(<39 a). The gas is distinguished from other substances by elchivg hard glass previously pre- pared by coating imperviously with (melted) wax, and writing through the coat, the 214 Silicon. operation may be conducted in a small leaden tray, or cup formed of sheet lead; the pul- verized fluoride being mixed with sulphuric acid to the consistence of paste. 741. If the fluoride be mixed with silicic acid, we have, instead of hydrofluoric acid, silicon fluoride, SiF4 (739 b); a gas which does not attack glass, but when passed into water produces hydrofluosilicic acid, H2SiF6 (739 c, d) (742).* Also, heated with acid sulphate of potassium, in the dry way, fluorides disengage hydrofluoric acid. If this operation be performed in a small test-tube, the surface of the glass above the material is corroded and roughened: CaF2 + 2KHSO.i = CaSO4 + K2SO4 + SHF By heating a mixture of borax, acid sulphate of potassium, and a fluoride, fused to a bead on the loop of platinum wire, in the clear flame of the Bunsen gas-lamp, an evanes- cent yellowish-green color is imparted to the flame. 742. HYDROFLUO SILICIC ACID, (HF)2SiF4, or H2SiF6, prepared as directed in the note to 741, forms metallic silicofluorides, mostly soluble in water; those of barium, sodium, and potassium, being only slightly soluble in water, and made quite insoluble by addition of alcohol. The silicofluorides are precipitated translucent and gelatinous. With concentrated sulphuric acid, they disengage hydrofluoric acid, HP. By heat, they are resolved into fluorides and silicon fluoride : as, BaF2 -(- SiF4. SILICON. Si = 28.195. 743. Occurrence.—It is never found in nature in a free state, but al- ways combined with oxygen in the form of silicon dioxide, Si02, as quartz, opal, flint, sand, etc. All geological formations except chalk contain sili- con, as the dioxide or as a silicate. Preparation and Properties.—There are three modifications of silicon. (1) Amorphous silicon, made by fusing K.jSiF,, with K, or SiCl4 with ISTa. K2SiF6 + 4K = 6KF + Si SiCl4 + 4Na = 4NaCI + Si It is a dark-brown powder, sp. gr. 2.0. Non-volatile and infusible ; burns in the air, forming Si02 ; and in chlorine, forming SiCl4. Not attacked by acids, except HF ; but dissolved by KOH. Si + 6HF = HoSiFe + 2H2 Si + 4KOH = K4SiQ4 + 2H2 (2) Graphoidal silicon is made by fusing the amorphous with Al, and then dissolving the Al with HCI. It fuses in, but is not oxidized by, oxy- gen ; HF has no action on it ; but is dissolved in a mixture of HF and HNOs, H2SiF6 being produced. Fused KOH acts slowly upon it. (3) * Hydrofluosilicic acid is directed to be prepared by taking one part each of fine sand and finely-pow- dered fluor-spar, with six to eight parts of concentrated sulphuric acid, in a small stoneware bottle or a glass flask, provided with a wide delivery-tube, dipping into a little mercury in a small porcelain capsule, which is eet in a large beaker containing six or eight parts of water. The stoneware bottle or flask is set in a small sand-bath, with the sand piled about it, as high as the material, and gently heated from a lamp. Bach bub- ble of gas decomposes with deposition of gelatinous silicic hydrate. When the water is filled with this de- posit, it may be separated by straining through cloth and again treating with the gas for greater concentra- tion. The strained liquid is finally filtered and preserved for use. Silicon. 215 Adamantine silicon is formed by fusing the graphoidal. It looks like hematite, and is oxidized Avith greater difficulty than the other modifications. 744. Silicic Anhydride, or silica, Si""02, is a stable, non-volatile, infusi- ble solid ; insoluble in water or acids, soluble in fixed alkali hydrate solu- tions, by formation of silicates. These are formed as normal salts, qua- dribasic ; metasalts, dibasic ; and in many other proportions of base and acid. Si02 + 4KOH = K4Si04 + 2H20 745. Of the Silicates, only those of potassium and sodium are soluble in water. The solutions of alkali silicates somewhat resemble, in the nature of their union, the alkaline solutions of zinc, aluminium, and lead. These silicates *iu solution are decomposed by all acids, including carbonic, with separation of silicic acid, H4Si04, gelatinous. Silicic acid is soluble, silicic anhydride insoluble, in the mineral acids. Some of the insoluble silicates are also dissolved by sulphuric and by hydrochloric acid, with separation of gelatinous silicic acid. Other silicates are dissolved by these acids, when heated in closed tubes at about 300° C. Soluble silicates precipitate, from salts of non-alkali metals, silicates insoluble in water, but mostly soluble in acids. Silicates are determined, qualitatively or quantitatively, by the separa- tion of the anhydride, 746. 746. Silicic acid is obtained as H2SiOs, H4Si04, and other hydrates of Si02. It is decomposed by evaporation to dryness, with a residue of silica, insoluble in acids.—Hence, when an alkali silicate is acidulated (with hydro- chloric or nitric acid), and evaporated strictly to dryness on the water-bath, and again treated with water and the same acid, the silica is left behind in- soluble. This behavior is characteristic for silicic acid, and serves for its separation as well as detection. The formation of the alkali silicate, from silica or an insoluble silicate, as the first step in this operation, is generally obtained in the dry way, by fusion with three or four parts of mixed carbonates of sodium and potas- sium in a porcelain crucible. (These carbonates, mixed in about molecular proportions, fuse at a much lower temperature than either alone.)—Also, the soluble silicates may be formed by boiling with solution of potassium or sodium hydroxide, as stated in 744; the operation being performed in a silver or platinum vessel. Silica dissolves, with more or less readiness, m boiling solution of potassium or sodium carbonate. Silica is not soluble in ammonium hydroxide, and the salts of ammo- nium separate gelatinous silica from solutions of potassium or sodium sili- cate. Therefore, in operating with these solutions, silica is precipitated with the bases of the third group : 747. In the fused bead of microcosmic sale, particles of silica swim uu- K4SiQ4 + 4NH4CI + 4H20 = H4SiQ4 + 4KCI + 4NH4OH 216 Phosphorus. dissolved. If a silicate be taken, its base will, in most cases, be dissolved out, leaving a “skeleton of silica” undissolved in the liquid bead.—But with a bead of sodium carbonate, silica (and most silicates) fuse to a clear glass of silicate. For the reactions of silica with fluorides, see 739. 748. Silica is separated from the fixed alkalies in natural silicates, by mixing the latter in fine powder with three parts of precipitated calcium carbonate, and one-half part of ammonium cldoride, and heating in a platinum crucible to redness for half an hour, avoiding too high a heat. On digesting in hot water, the solution contains all the alkali metals, as chlorides, with calcium chloride and hydrate. [J. Lawrence Smith.] PHOSPHORUS. P = 30.958. 749. Specific gravity of the solid, 1.814 ; of the liquid, 1.7555 (Damien, 1881); of the red, 2.34 (Teoost, 1875). Vapor density (H = l), 62, indi- cating that its molecule is P4. Yellow phosphorus melts at 44.4° 0. (112° F.), and boils at 287.3° C. (549° F.) (Pisati, 1875). Tinder pressure of 26.2 atmospheres boils at 511° C. (952° F.) (Teoost, 1873). lied phosphorus melts at 255° 0. (491° F.) (Doling). Occurrence.—lt is never found free in nature. It is found in the primitive rocks as calcium phosphate, occasionally as aluminium, iron, or lead phosphate, etc. Plants extract it from the soil, and animals from the plan ts. Hence traces of it are found in nearly all animal and vegetable tissues; more abundantly in the seeds of plants and in the bones of animals. Preparation.—(l) From bones. They are first burned, which leaves a residue, consisting chiefly of Cas(P04)2 ; then H2S04 is added, producing soluble calcium tetrahydrogen diphosphate {a). After filtering from the insoluble calcium sulphate it is ignited, leaving calcium metaphosphate (b). Then fused with charcoal, reducing two-thirds of the phosphorus to the free state (c). The admixture of sand, Si02, with the charcoal is preferred, in which case the whole of the phosphorus is reduced (cl)- Hydrochloric acid passed over red hot calcium phosphate and charcoal reduces the whole of the phosphorus. This process works well in the laboratory, and has also been successfully employed on a larger scale. Either of the calcium phos- phates may be used (e) and (/). a. Ca3(PO4)2 + 2H2S04 = 2CaSO( + CaH4(PO4)2 b. CaH4(P04)2 + ignition = Ca(PO3)2 + 3H20 c. 3Ca(PO3)2 + lOC = Ca3(PO4)2 + 10CO +P4 d. 3Ca(PO3)2 + lOC + 2Si02 = 2CaSiOs +P4 + 10CO e. 2Ca3(PO4)2 + 100 + 12HC1 = GCaOI, + P4 + 16CO + GH2 /. 2Ca(PC3)2 + 12C + 4HCI = 2CaCI2 +P4 + 12CO + 2H2 750. Properties of Yellow or Ordinary Phosphorus,—lt has a trans- parent, slightly yellowish color, but turns white at the surface after some Phosphor us. time. Its odor, sometimes described as garlic, which it scarcely resembles, is well known. It is very poisonous ; one-half a grain has proved fatal. When inhaled in small quantities, as in match-making, it causes toothache, followed by decay of the teeth ; later aching of the jaw, followed by decay of the jaw. A trace of it dissolves in water. Alcohol dissolves 0.4, ether 0.9, olive oil 1.0, and turpentine 2.5 per cent, of it, while carbon disulphide dissolves 10 to 15 times its own weight. It oxidizes in the air at high tem- peratures to H3P04 ; at lower temperatures, or with an insufficient supply of air, to H3P03. Phosphorus is luminous in the dark. This luminosity is the result of slow combustion. It is not luminous at 0° C. Its luminosity increases with the temperature, and if exposed freely to the air the heat of combus- tion raises its temperature, until at 44° C. it bursts into flame. It is best preserved under water. In a fine state of division, such as is produced by its evaporation from the carbon disulphide solution, it takes fire at once at temperatures in which the compact phosphorus can be kept for days. Boiled with alkali or alkaline earth hydroxides, it forms a hvpophos- phite and phosphorus hydride, PH3. Phosphorus is largely used in match- making. Yellow phosphorus is used in the ordinary match, and the red (amorphous) in the safety matches, the phosphorus being on a separate sur- face. Phosphorus is usually estimated as a phosphate after oxidation with HNO, 751. Oxidation.—In presence of water, Cl, Br, and I form HCI, HBr, HI, and H3P04. Chloric, bromic, and iodic acids give the same products. When dry, finely-divided phosphorus is mixed with substances which readily part with oxygen, such as potassic chlorate or metallic peroxides ; very slight friction is sufficient to cause the explosive oxidation of the phos- phorus. When sodic carbonate is heated to redness with phosphorus, the carbonic anhydride is reduced and carbon is set free. Owing to its affinity for oxygen, phosphorus acts as a powerful reducing agent. Platinum, gold, silver, and copper are deposited in a metallic state when white phos- phorus is left in contact with the solution of their salts. 752. Red or Amorphous Phosphorus.—When ordinary phosphorus is heated for 40 hours at 240° 0. to 250° C., the red variety is formed ; but if the heat is allowed to rise above 260° C. it is again changed to ordinary phosphorus; but under pressure ordinary phosphorus heated to 300° 0. is immediately and completely changed to the red. If finely divided it has a scarlet-red color ; in larger particles it has a dull or brownish-red color. It is odorless, tasteless, non-poisonous, and non-luminous in the dark. It is insoluble in water, ether, or carbon disulphide. It requires no special precaution for its preservation ; does not (if free from traces of the other variety) oxidize in the air, and does not ignite until heated to 240° C. (404° F.) HNOs, Cl, Br, and I act as upon the ordinary phosphorus. 218 Phosphorus Hydride—Rypophosfhorous Acid. but more slowly. A black variety of phosphorus is formed by suddenly cooling the melted ordinary phosphorus to 0° 0. This and the white opaque crust which forms on ordinary phosphorus have not been much studied. PHOSPHORUS HYDRIDE. PH3. Oxidation valence P~'"H'a H ! Structural valence H-P-H Vapor density, 17, indicating that the molecule is PH3. 753. Preparation.—(l) By boiling phosphorus in an alkali or an alkaline earth. (2) By ignition of H3P02 or H3P03. (3) By ignition of hypophos- phite of the alkalies or alkaline earths. (4) By action of calcium or mag- nesium phosphide on water or HCI. 1. 3KOH +P4 + 3H20 = 3KH2PO2 + PH3 3. 2H3POo = H3PQ4 + PHS 4H3P03 = 3HsP04 + PHS 3. 4NaH2PO2 = Na4P207 + 2PH3 + H2O 4. Ca3P2 + 0H2O = 3Ca(OH)2 + 2PH3 Ca3P2 + OHCI = 3CaCl2 + 2PH3 Properties.—A colorless gas having a very disagreeable odor. It is lique- fied by pressure and cold. As usually prepared, it is spontaneously inflam- mable. It is a strong reducing agent; transposes many metallic solutions—• e.g., CuS04. It reduces solutions of gold and silver to the metallic state. It burns in the air, forming water and metaphosphoric acid. It is oxidized by hot H2S03, H2S04, chlorine-water, hypochlorites, N2O, 140, HN03, H3AsQ4, etc. 3CuS04 + 2PH3 = Cu3P2 + 3H2504 BAgNO„ + PH3 + 4H20 = H3P04 + BHNOs + BAg 2PH3 + 402 (ignition in the air) = 2HPO3 + 2H20 A liquid phosphorus hydride, P2H4, and a solid, P4H2, are known. HYPOPHOSPHOROUS ACID. H3P02. Oxidation valence H'3PO""a H * Structural valence H-0-P=o i H 754. The anhydride, P2O, has not been obtained. The acid is formed r (1) By transposing lead hypophosphite with H2S ; or the barium salt with * The structural formula represents the atom of phosphorus as holding three positive bonds (toward oxygen) and two negative bonds (toward hydrogen), so that the sum of its valence is one See paragraph 614. Hypophosphorous Acid. 219 H2S04 added in exactly molecular proportions ; or by treating the calcium salt with a like quantity of H2C204. Pb(H2PO2)2 + H2S = PbS + 2H3P02 8a(H2P02)2 + H2S04 = BaS04 + 3H3P02 Properties.—By evaporation in a vacuum the liquid H3P02 is ob- tained, which, on cooling to 174° 0., solidifies (Thompson, 1874). In forming salts only one-third of the hydrogen of the acid is displaced—e.g.y sodium hypophosphite is NaH2P02, not Na3POs. It rapidly oxidizes if ex- posed to the air, forming phosphoric acid. Heat decomposes it, forming PH3 and H3P04 (or HPOa if at a red heat). On ignition hypophosphites leave pyrophosphites, evolving PH3. Hypophosphorons acid may be known from phosphorous acid by adding cupric sulphate to the free acid and heat- ing the solution to 55° 0. (131° F.) With hypophosphorons acid a red- dish-black precipitate of copper hydride (Cu2H2) is thrown down, which, when heated in the liquid to 100° C. (313° F.), is decomposed with the de- position of the metal and the evolution of hydrogen, whilst with phos- phorous acid the metal is precipitated and hydrogen evolved, but no Cu2H2 is formed. Further, hypophosphorons acid reduces the permanganates imme- diately, but phosphorous acid only after some time. Phosphites precipitate barium, strontium, and calcium salts, while hypophosphites do not. When hypophosphorons acid is treated with zinc and sulphuric acid it is converted into phosphoretted hydrogen. On boiling hypophosphorons acid with ex- cess of alkali hydroxide, first a phosphite then a phosphate is formed, with evolution of hydrogen. Estimation.—(1) By oxidation with nitric acid and then proceeding as with phosphoric acid. (3) By mercuric chloride acidulated with HCI; the temperature must not rise above 60° 0., otherwise metallic mercury will be formed. The precipitated Hg2Cl2, after washing and drying at 100° 0., is weighed. NaHjPOo + 4HgCl2 + 2HaO = 3Hg2Cl2 + HBP04 + NaCl + 3HCI 755. Production of Hypophosphites. —All ordinary metals form hypo- phosphites except tin, copper, and mercurosum. Silver and ferric hypo- phosphites are very instable. (1) A few metals, such as zinc and iron, dis- solve in H3P02, giving off hydrogen and forming a hypophosphite. (3) The alkali and alkaline earth salts may be formed by boiling phosphorus with the hydroxides. (3) As all hypophosphites are soluble, none can be formed hy precipitation. All may be formed from their sulphates by transposition with barium hypophosphite. ,(4) All may be made by adding to the carbonates or hydroxides of the metals. 756. Oxidation.—ln all cases an excess of the oxidizing agent produces phosphoric acid or a phosphate. Perhaps the best method of proving that H3PO„ is all changed to H3P04 is that it fails to blacken argentic nitrate. HYPOPHOSPHOR OUS A CID. The reduction of mercuric chloride is sometimes preferred, especially if hydrochloric acid is present. Where the oxidation is not fully complete, first remove any phosphate which may be present as an impurity, by addition of magnesium sulphate in presence of ammonium chloride and ammonium hydroxide ; then, after oxidation, repeat the process to prove partial oxidation. HN02 and HN03 a. Become NO. H„S03 h. Becomes free sulphur. H2S04 c. Becomes first sulphurous acid, then sulphur. Cl d. Becomes'hydrochloric acid, and phosphoric acid is formed ; with alkalies a chloride is formed. HCIO e. Becomes hydrochloric acid. With alkalies a chloride is formed. HCIO. /. Becomes hydrochloric acid. Prove same as above. Br g. Forms hydrobromic acid. The action takes place also in alkaline mixture, forming a bro- mide. HBr03 h. Forms hydrobromic acid. I i. Forms hydriodic acid ; and in alkaline mixture an iodide. HI03 /. Forms hydriodic acid. h. Forms ferrocyanic acid. In this case the formation of the same cannot be proven by addition of ferric chloride, because an excess of hypophosphorous acid changes ferric chloride to ferrous chloride, which then gives a precipitate with ferricyanide of potassium. A good method is to add a slight excess of fixed alkali, and then an excess of alcohol which will precipitate the ferrocyanide of potassium, which may, after washing with alcohol, be dissolved in water and tested in the usual way. Pb"" I. Becomes Pb" both in presence of fixed alkali and acids. Ag' m. Becomes Ag° “ “ “ “ Hg' n. Becomes Hg° “ “ “ “ Hg" 0 Becomes Hg° “ “ “ “ Asv and As'" p. Become metallic arsenic (in presence of hydrochloric acid). Bi'" q. Becomes Bi". The action takes place both in presence of alkalies and acetic acid. Cu" r. Becomes (Cu2)"or Cu2H2, and, on boiling, metallic copper is formed. See 754. Mn"+W 8. In presence of acids gives Mn". (No action in presence of alka- lies.) MnVll t. With alkalies gives Mniv. With jueids gives Mn". Phosphorous Acid—Phosphoric Acid. 221 Co'" u. Becomes Co". Ho action in alkaline mixture. Hi"' v. Becomes Hi", “ “ (Pe2)VI w. Becomes Pe". “ “ CrVI x. Becomes Cr'". “ “ PHOSPHOROUS ACID. H3POr Oxidation valence .H'3P "0~"3 H I 0 O 1 1 Structural valence H-O-P-O-H, or H-O-P-O-H I H 757. Phosphorus anhydride, P203, is formed by the combustion of phosphorus at a low temperature or in a limited supply of air. It is a volatile, crystalline solid, inflammable, and having a garlic-like odor. It can- not be formed by heating H„POa, which would produce H3PC4 and PH,, Heated in a sealed tube it decomposes thus : SP„03 = 3P2Ob + P4. Phos- phorous acid is formed by the action of water upon the anhydride or upon the trichloride. PC13 + 3H,0 = HsPOa + 3HCI Two atoms of its hydrogen, rarely only one, are displaced by metals—e.g., K2HP03, PbHP03, BaHPO,,. It is a strong reducing agent, oxidizing to phosphoric acid when exposed to the air. Reduces salts of silver and gold to tlie metallic state. It is changed to phosphoric acid by most of the strong oxidizing acids, and by many of the higher metallic oxides. PHOSPHORIC ACID H3PQ4 Oxidation valence H'SPV O"4 0 ii Structural valence H-O-P-O-H 1 0 1 H 758. Phosphoric anhydride, P2Os, is prepared by burning phosphorus in oxygen ov air It cannot be formed by heating H3P04, which would form volatile metaphosphoric acid, HPOs. It is a white, flaky, deliquescent solid, fusible, subliming unchanged at a red heat. It dissolves in water, forming three varieties of phosphoric acid : (I) Metaphosphoric acid, HPOs. (2) Pyrophosphoric acid, H4P„07. (3) Orthophosphoric acid, H3P04. 759. Metaphosphoric acid, HPOa,* is made by the action of cold water * For structural formula see Remsen's Theoretical Chemistry, 3d ed.,p. 185. 222 Phosphoric Acid. on P206, or by the action of a red heat upon HBP04 or H4PaOT. Also by action of H2S on the lead salt, or of H2S04 on the barium salt. It is slowly changed to HsP04 by boiling with water. Metaphosphates are especially distinguished, by the means mentioned in 769 and 771. Also, they are not precipitated by solutions of mag- nesium salts with ammonium hydroxide, unless very concentrated, or by the molybdate solution. The silver precipitate, AgP03, ivhite, is soluble in alkali metaphosphate solutions, distinction from pyrophosphates. Free metaphosphoric acid precipitates solutions of silver nitrate, lead nitrate, and lead acetate, the precipitates being insoluble in excess of meta- phosphoric acid, and soluble in moderately dilute nitric acid. Barium, calcium, and ferrous chlorides, and magnesium, aluminium, and ferrous sul- phates, are not precipitated by free metaphosphoric acid. Ferric chloride is precipitated, a distinction from orthophosphoric acid. See 769. There are various polymeric modifications of metaphosphoric acid, distinguished from each other chiefly by physical differences of the acids and their salts. Pure meta- phosphoric acid is a white, viscid, or waxy solid. (Ordinary glacial phosphoric acid owes its hardness to the universal presence of sodium metaphosphate.) Fusion with excess of sodium carbonate converts both metaphosphates and pyrophos- phates to normal orthophosphate. 760. Pyropliosphoric acid, H4P„O.,* is formed by heating H3P04 to 215° 0., or by the action of H2S on the lead salt and by the action of H2S04 on the barium salt. Pyrophosphates are precipitated from their solutions by silver nitrate, as silver pyrophosphate, Ag4P207, white, soluble in ammonium hydrate and in nitric acid. The pyrophosphates of the alkaline earth metals are difficultly soluble in acetic acid. The most of the pyrophosphates of the heavy metals, except silver, are soluble in solutions of alkali pyrophosphates, as double pyrophos- phates soluble in water (distinctions from orthophosphates). Ammonium molybdate reacts but slowly with pyrophosphate solutions— and not until orthophosphate is formed hy digestion with the nitric acid of the reagent solution. Magnesium salts with ammonium hydroxide give a precipitate of double pyrophosphate, soluble in alkali pyrophosphate solution. Free pyropliosphoric acid gives precipitates with solutions of silver ni- trate, lead nitrate or acetate, and ferric chloride; no precipitates with barium or calcium chloride, or with magnesium or ferrous sulphate. (Fur- ther, see 769.) 761. Orthophosphoric acid is made : (1) By boiling P205, HPOs, or h4p2o7 in water. (2) By boiling phosphorus in dilute nitric acid. (3) By the combustion of PH3 in moist air. (4) By action of water on PCI,.. * For structural formula see Remsen's Theoretical Chemistry, 3d ed , p. IS'. Phosphoric Acid. 223 Orthophosphoric acid, H3P04, is a translucent, feebly erystallizable, and very deli- quescent soft solid ; reduced by heat first to pyrophosphoric acid, then to metaphos- phoric acid, which is volatile. Orthophosphoric acid is formed from phosphorus by oxidation in water; and from metaphosphoric acid or pyrophosphoric acid by digesting with dilute mineral acids, or even by long boiling in water, or, as sodium salt, by fusion with excess of sodium car- bonate. Phosphoric acid is formed from metallic phosphates by transposition with acids in cases where a precipitate results, as a lead or barium phosphate with sulphuric acid, or silver phosphate with hydrochloric acid. But when the products are all soluble, as cal- cium phosphate with acetic acid or sodium phosphate with sulphuric acid, the transposi- tion is only partial ; so that unmixed phosphoric acid is not obtained. A non-volatile acid, like phosphoric, is not separated from liquid mixtures, as the, volatile acids are, like hydrochloric. The change represented by equation (a) can be so verified that pure phos- phoric acid will be separated, but the changes shown in equations (b) and (c) do not com- prise the whole of the material taken. In the operation (b) some sodium phosphate and some nitric acid will be left, and in (c) some trihydrogen phosphate will no doubt be made. a. CaH4(PO4)2 + H2C204 = CaC204 + 2H3P04 b. Na2HP04 + 2HNO, = 2NaNOs + HsP04 And Na2HP04 + HN03 = NaNOs + NaH2P04 c. 2CaHP04 + 2HCI = CaCl2 + CaH4(P04)2 762. The Orthophosphates, dimetallic and tr{metallic, are insoluble in water—except those of the metals of the ordinary alkalies. They are all, however, more or less soluble in aqueous phosphoric acid by formation of monometallic salts, as CaH4(P04)2, having an acid reaction. Lithium phos- phate is nearly insoluble in water. Phosphates are insoluble in alcohol. In analysis, the molybdate precipitate (7GB) is most distinctive. Separa- tion by the ferric phosphate precipitate in presence of acetic acid (306) is employed. Separation from oxalate, as calcium precipitate, by acetic acid, is used in systematic qualitative work (764). Ignition test: see 771. 763. Soluble salts of all metals, except those of the alkalies, pre- cipitate solutions of ordinary phosphates (dimetallic and trimetallic ortho- phosphates). 764. Soluble salts of the alkaline earth metals, with dimetallic alkali phosphates, as NaJEEP04, form white precipitates oE phosphates, two-thirds metallic, as CaHP04; with trimetallic alkali phosphates, white precipitates of phosphates, normal or full metallic, as Cas(P04)2. The precipitates are soluble in acetic acid, and in the stronger acids (761 c). Concerning the ammonium magnesium phosphate, see 79. 765. Solutions of orthophosphates give, with soluble ferric, chromic, and aluminium salts, mostly the normal phosphates, Fe2(P04)2, etc. The ferric phosphate is but slightly soluble in acetic acid, and for this reason it is made the means of separating phosphoric acid from metals of the earths and alkaline earths (306). Solution of sodium or potassium acetate is Phosphoric Acid. added ; and if the reaction is not markedly acid, it is made so by addition of Acetic Acid. Ferric chloride (if not present) is now added, drop by drop, avoiding an excess. The precipitate, ferric phosphate, is brownish- Avhite. With zinc and manganous salts, the precipitate is dimetallic or normal— ZnHP04, or zn,(PO.) 2—according to the conditions of precipitation. Man- ganic salts give a colored solution, as explained in 352. With salts of nickel, a light green normal phosphate is formed ; with cobalt, a reddish normal phosphate. 766. Silver salts precipitate normal silver phosphate, Ag3P04, light yel- low, soluble in acetic and nitric acids and in ammonium hydroxide. The color of the silver precipitate distinguishes ortho- from pyro-phosphoric acid. Lead salts precipitate mostly Pb3(P04)2, but slightly soluble in acetic acid. Bismuth salts form BiP04, peculiar in being insoluble in dilute nitric acid. Copper forms a biuish-wbite precipitate, either normal or two- thirds metallic. Mercurous salts precipitate mercurous phosphate, Hg6- (P04)2, white. Mercuric nitrate (not the chloride) precipitates mercuric phosphate, Hg3(P04)2, white. 767. If a disodium or dipotassium orthophosphate is added to solution of silver nitrate, free acid is formed, and an acid reaction to test-paper is induced (a). But with a trisodium or tripotassium phosphate, the solution remains neutral {b)—a, means of dis- tinguishing the acid phosphates from the normal. a BNa2HPO4 + OAgNO3 = 2Ag3P04 + H3P04 + 6NaNO* b. Na3P04 + 3AgN03 = Ag3PQ4 + 3NaNO3 768. Ammonium Molybdate, in its nitric acid solution (604), fur- nishes an exceedingly delicate test for phosphoric acid, giving the pale yellow precipitate, termed ammonium phosphomolyhdate. The molybdate should be iu excess, therefore it is better to add a little of the solution tested (which must be neutral or acid) to the reagent, taking a half to one cub. cent, of the latter in a test-tube. For the full delicacy of the test, it should be set aside, at 30° to 40° C., for several hours. 769. Free orth©phosphoric acid is not precipitated by ordinary salts of third and fourth group metals (in instance of ferric chloride, a distinction from pyrophosphoric acid and metaphosphoric acid),* but is precipitated in part by silver nitrate, and lead nitrate and acetate. Ammoniacal solu- tion of calcium chloride or of barium chloride precipitates free ortho- phosphoric acid. * A solution containing 5 p. c. ferric chloride, mixed with one-fourth»its volume of a 10 p. c. solution of orthophosphoric acid, requires that near half of the latter be neutralized (so that phosphate is to phos- phoric acid as 1.114 is to 1.000) before precipitation occurs. On the other hand, 4c.c.ofa 5 p. c. solution of ferric chloride, mixed with 1 c. c of a 6 p. c. solution of metaphosphoric acid, form a precipitate, to dissolve which 30 c. c. of the same metaphosphoric acid solution or 5 c. c. of a 24 p. c. solution of hydrochloric acid are required. Four c. c. of asp. c. solution of silver nitrate with Ic.c.ofalo p. c. solution of orthophos- phoric acid give a precipitate, to dissolve which requires 7 c. c. of the same orthophosphoric acid solution. [The Author’s report of work by Mr. Morgan, Am. Jour. Phar., xlviii. (1876), 534.] Sulphur. 225 770. Orth ©phosphoric acid—or an orthophosphate with acetic acid— does not coagulate egg albumen or gelatine. This is a distinction of both orthophosphoric acid and pyrophosphoric acid from metaphosphoric acid. 771. Ignition with metallic magnesium (or sodium) reduces phosphorus from phos- phates to magnesium phosphide, P.,Mg3, recognized by odor of PH3. formed on contact of the phosphide with water. A bit of magnesium wire (or of sodium) is covered with the previously ignited and powdered substance in a glass tube of I he thickness of a straw,, and heated. If any combination of phosphoric acid is present, vivid incandescence will occur, and a black mass will be left. The latter, crushed and wet with water, gives the odor of phosphorus hydride. SULPHUR. S = 31.984 772, Specific gravity, 1.95 to 2.05. Melting point, 114.5° C. (238° F.) (Brodie). Vapor density (H =l) above 800° C. is 32, showing that the molecule is S2; but at 480° 0. it is 96, showing that the molecule at that temperature is S6. Occurrence.—(l) Found in a free state, and as S02 in volcanic dis- tricts. (2) As H2S in some mineral springs. (3) As a sulphide. Iron pyrites, FeS2; copper pyrites, CuFeS„; orpiment, As2S3; realgar, As2S2 ; blende, ZnS ; cinnabar, HgS ; galena, PbS. (4) As a sulphate. Gypsum, CaS04 + 2H„O ; heavy spar. BaS04; kieserite, MgS04 -f- H2O ; bitter spar (epsom salts), MgS04 -fi 7H20 ; glauber salt, Ha2S04 + 10H20, etc. Preparation.—(l) The native sulphur is separated from the clay and rock in which it is embedded, partly by melting and partly by distillation. (2) From PeS2 by heating in close cylinders 3FeS2 =Fe3S4 -f- S2 ;orat a higher temperature FeS2 = FeS -f- S. Much of the sulphur contained in pyrites is converted into and utilized as sulphuric acid. 773. Properties.—Sulphur is a solid—in yellow, brittle, friable masses (from melting); or in yellowish, gritty powder (from sublimation); or in nearly white, slightly cohering, finely crystalline powder (by precipitation from its compounds). It melts at about 114,5° C. (238° F.); at higher temperatures, it suffers peculiar physical modifications of consistence, etc.; and distils at 448.4° 0. It is not sensibly volatile at ordinary temperatures, hut has a slight, characteristic odor. In chemical activity, volatility, and other properties, sulphur stands as the second member of the Oxygen Series: O, 15.9633 ; S, 31.984 ; Se, 78.797; Te, 127.96. Sulphur is insoluble in water; slightly soluble in alcohol and in ether, freely soluble in carbon disulphide; but with physical solvents other than Water, its different modifications have different solubilities. On being heated it melts at 114 5° C. to a pale yellow liquid ; as the temperature rises it grows darker and thicker, until at about 180° C. it is dearly solid, so that the dish may be inverted without spilling. At 260° C. 226 Hydr os ulpu uric A cm. it again becomes a liquid as at first; and at 448.4° C. it boils and is con- verted into a brownish-red vapor. If it is slowly cooled, exactly the same physical changes take place in the reverse order, becoming thick at 180° C. and thin again at 114.5° Cand at lower temperatures solid. If, at a temperature near its boiling point, it is poured into cold water, it forms a soft, ductile, elastic string, resembling india-rubber. In a few hours this ductile sulphur changes back to the ordinary form, the change evolving heat. But if poured into water from the other liquid form—that is, at 114.5° o.—it forms only ordinary, brittle sulphur. Ordinary and precipi- tated sulphur is soluble in CS2; the ductile variety is insoluble. Fre- quently commercial samples are found to contain a definite per cent, of one of the insoluble varieties. 774. Oxidation.—Sulphur, when fused with the following elements, combines with them to form sulphides: Pb, Ag, Hg, Sn, As, Sb, Bi, Cu, Cd, Zn, Co, Ni, Pe, Sr, Ca, Mg, K, Na, In, Tl, Pt, Pd, Rh, Ir, In, Ce, La, Di. Sulphur dissolves readily in hot solutions of hydroxides of potassium, sodium, cal- cium, or barium, forming supersulphides and thiosulphates {a). These can be separated by alcohol, in which the sulphides dissolve. Sulphur is acted upon slowly by active oxi- dizing agents, as hot concentrated nitric acid (b), or chlorine generated in presence of water (c), with formation of sulphuric acid. Hot concentrated sulphuric acid very slowly oxidizes sulphur to sulphurous anhydride, by its own reduction to the same com- pound {d}: a. 3Ca(OH)2 + GS.2 = 2CaS5 + CaS203 + 8H20 h. S2 + 4HN03 = 2H2504 + 4NO c. S2 + 6Clo + BHOO = 2H0504 + 12HC1 d. S2 + 4H2504 + 2H20 = CHoSOs In the air, at ordinary temperatures, finely divided sulphur is very slightly oxidized, by ozone, to sulphuric acid ; at about 260° C. (500° F.) it begins to oxidize rapidly to sulphurous anhydride, burning with a blue flame. gvi-« becomes SVI when fused with alkaline carbonate and nitrate or chlorate. That is, free sulphur, S°, or any compound containing sulphur with less than six bonds, is oxidized to a sulphate if fused with an alkaline nitrate or chlorate, nitric oxide or a chloride being formed and carbon di- oxide escaping. HYDROSULPHURIC ACID. H2S. Oxidation valence Structural valence H-S-H 775. Vapor density (H = l), 17. Melting point (freezing point),-85.5° C. (-122° F.) (Faraday). Boiling point, -61.8° C. (-79.3° F.) (Reg- istault). Becomes a liquid at 10° C. (50° F.) under a pressure of 17 at- mospheres. Hydrosulphuric Acid. 227 Occurrence.—Pound in volcanic gases and in many mineral springs. Its presence in spring-water is accounted for by the reduction of sulphates to sulphides by organic matter and subsequent liberation of H2S by car- bonic acid. 776. Preparation.—For laboratory use it is prepared by action of dilute acids (usually H2S04) upon ferrous sulphide. The ferrous sulphide is pre- pared either by fusion of the iron with the sulphur, or by bringing red hot iron rods in contact with sticks of sulphur. The ferrous sulphide melts as fast as formed, and is made to drop into a tub of cold water. Dilute H2S04 should be used («). Concentrated H2S04 has no action on FeS, unless heated, and then SOa is evolved (£)'; and frequently free sulphur is formed (if only a little water is present) by the action of the H2S upon the H2S03 first formed. Hydrosulphuric acid is slowly formed if hydrogen be passed through melted sulphur ; also if sulphur be fused with tallow or paraffine, and when organic substances (e.g., eggs) containing sulphur are allowed to decay. a. FeS -f H2S04 = FeSQ4 + H2S b. 2FeS + 10H2SC>4 = Fe2(504)3 + 0SO2 + 10H20 777. Properties.—lt is a colorless, poisonous gas, having the odor of rotten eggs. By pressure and cold it may be condensed, first to a liquid, then to a solid. At 0° C. water dissolves 4.37, and at 15° C. 2.66, volumes of the gas. The solution in open vessels vaporizes gas constantly, at ordinary tem- peratures, until exhausted ; more rapidly when boiled. Both gas and solu- tion feebly redden moist litmus paper ; and have a very strong, character- istic odor. (The concentrated gas is a quick poison, by inhalation.) Absolute hydrosulphuric acid is combustible in the air—burning with a blue flame to sulphurous anhydride and water. 778. The solubility of the Metallic Sulpirides in water, dilute acids, hot nitric acid, and in alkali sulphides, is shown in the grouping of the bases, and the sub-grouping of the second-group precipitates. In analysis, sulphides are known by generation of H2S (779), or sepa- ration of S by oxidizing solvents, and by the color test with nitroferricya- nide (782). 779. Sulphuric acid, dilute, transposes the metallic sulphides ; except those of arsenic, tin, mercury, silver (and lead), which are decomposed with difficulty, or not at all : PeS + H2S04 = PeS04 + H2S The gaseous hydrosulphuric acid, when liberated, is recognized by its odor, by blackening paper moistened with lead acetate, or with a solution of a lead salt with excess of potassium or sodium hydroxide (388 and 387). In the detection of traces of the gas, a slip of bibulous paper, so moistened, may be inserted into a slit in the smaller end of a cork, which Hyd rosulphuric Acid. 228 is fitted to the test-tube, wherein the material to be tested is treated with sulphuric acid ; the tube being set aside in a warm place for several hours. A very delicate test is made by conducting the gas into ammoniacal solu- tion of nitroferricyanide (782). If any oxidizing agents are present—as chromates, ferric salts, manganic salts, chlorates, etc.—hydrosulphuric acid is not generated, but instead sulphur is separated, or sulphates are formed (780). The sulphides not transposed with hydrochloric or sulphuric acid, are recognized by the separation of sulphur on treatment with nitric acid, or with nitrohydrochloric acid. Also, these sulphides and certain super- sulphides, attacked with difficulty by acids, as iron pyrites and copper py- rites, are reduced and dissolved, with evolution of hydrosulphuric acid, by dilute sulphuric acid with zinc. The gas, with its excess of hydrogen, may be tested by 782. 780. Hydrosulphuric acid is a strong reducing agent, and the metallic sulphides act in the same capacity with a greater or less degree of force. The reactions with oxidizing agents are given at length in 785. The hydrogen of H2S takes oxygen readily ; the sulphur more slowly. In the oxidation of metallic sulphides, generally, less sulphur is left unoxi- dized than occurs in the oxidation of hydrosulphuric acid—owing to the stronger tendency to form sulphates. 781. Solutions of metallic sulphides give precipitates with soluble salts of second and third group metals ; hydrogen sulphide, with salts of second-group metals only. The precipitates are sulphides, except with chromium and aluminium ; reduction occurring with ferric and arsenic salts, which form ferrous and arsenious sulphides. The precipitates have strongly marked colors—that of zinc being white; those of iron, copper, and lead, Hack; arsenic, yellow ; antimony, orange-red; mercury, succes- sively white, yellow, orange, and Hack. 782. Solutions of nitroferricyanides (693) give, with soluble metallic sulphides (or with hydrosulphuric acid after addition of an alkali, or with free sulphur after digesting with an alkali), an intense, rich purple color, disappearing after some time. Add a drop of the reagent, to a few drops of the solution, on a white porcelain surface. Vapors are tested for hydrosulphuric acid by conducting them into ammoniacal solution of sodium nitroferrieyanide. (Vapors are tested for ammonia by passing them into solution of nitroferricyanide with hydrosulphuric acid.) 783. By ignition in the air, sulphur gives its characteristic odor of sulphurous anhy- dride. Many of the sulphides yield more or less sulphurous anhydride ; most of them are also, partly or wholly, converted to sulphates. When ignited on charcoal with sodium carbonate—or (distinction from sulphates) if ignited in a porcelain crucible with sodium carbonate—soluble sodium sulphides are obtained. The production of the sodium sulphide is proved by the black stain of AgoS, formed on metallic silver by a moistened portion of the fused mass. (Compounds of selenium and tellurium, 612.) Hydro sulphuric Acid. 784. Formation of Sulphides.—(l) By fusion of the metals with sul- phur, see 774. (2) By action of H2S upon the free metals, hydrogen being evolved. With Hg and Ag this occurs at ordinary temperature, but with most metals a higher temperature is needed. (3) Action of H2S on me- tallic oxides or hydroxides. Those sulphides which are decomposed by water {e.g., A12S3, Cr2S3) are not formed in its presence, but by action of H2S upon the oxide at a red heat. (4) By action of soluble sulphides upon metallic solutions. The ordinary sulphides of the first three groups are formed thus, except ferric salts, which are precipitated as FeS, and alumi- nium and chromic salts as hydroxides. (5) By action of CS2 upon oxides at a red heat. (6) By action of free sulphur upon oxides at a red heat. (7) By the action of charcoal upon the oxyacids of sulphur at a red heat in presence of an alkaline carbonate. 1. 2Fe + S2 = 2FeS 2. 2Ag + H2S = Ag2S +H2 3. Pb(OH)2 + H2S =r PbS + 2H20 2Fe2(OH)6 + 6H2S = 4FeS +S2 4- 12H20 4. 2Fe2C16 + 6(NH4)2S = 4FeS +S2 + 12NH4C1 5. 2CaO + OS2 = 2CaS + C02 6. 4CaO + JIP., = 4CaS + 250., 7. K2SO4 + 20 = K2S + 2C02 2Bi2(504)3 + GNa2C03 + 150 = 4Bi + 6Na2S + 21C02 Estimation.—The sulphur in sulphides is oxidized by nitric acid, or chlorine, or by fusion with sodium carbonate and nitrate, and, after precipi- tation with barium chloride, weighed as barium sulphate. 785. Oxidation.—a. Free sulphur liberated from hydrosulphuric acid may sometimes be recognized simply by its appearance. But when white precipitates are formed at the same time, the whole should be allowed to settle, then the sulphur dissolved in carbon disulphide, and again sepa- rated by evaporation, or precipitated from the carbon disulphide solution by addition of alcohol, and then further tested by 773. For action upon HNO,, and HN05, see 713 and 733. H2SO, h. Forms water and sulphur [Watts’ Dictionary, 3, 203]. Sometimes, especially if the moist gases are used, pentathionic acid, H2S506, is formed. 10SO2 + 10H2S 2H2S6Os + 8H20 + 5S2. H2S04 c. No action if the sulphuric acid is dilute. With strong acid, sulphur and sulphurous anhydride are formed. To prove the latter, add sulphuric acid to dry ferrous sulphide and boil, or pass hydrosulphuric acid gas into hot sulphuric acid, and S02 will be evolved. Cl d. Forms first sulphur, and finally sulphuric and hydrochloric acids. This takes place in alkaline mixture also, forming a sulphate. HCIO e. Same as above. 230 Thiosulphuric A cid. HCI03 /. With excess of hydrosulphuric acid, free sulphur and hydro- chloric acid are formed. With excess ofHHOls,Os, sulphuric acid is formed. Br g. Forms hydrobrornic acid and sulphur. In alkaline mixture a sulphate is formed. HBr03 h. Forms sulphur and hydrobrornic acid; with excess of HBrOs, sul- phuric acid is formed. I i. Forms sulphur and hydriodic acid. HI03 j. With excess of hydrosulphuric acid, hydriodic acid and sulphur are formed. H.Fe,(CU)„ h. Forms potassium ferrocyanide and sulphur. Proof: Boil to expel excess of hydrosulphuric acid, then add ferric chloride. Pb02 I. Forms PbS and sulphur. Asv. As2S3 and free sulphur are formed. Mn" +'» m. Forms In'' and (S2)°. That is, all compounds of manganese having more than two bonds are reduced to the dyad, and free sulphur is formed. K2Mn2Oa n. With potassium sulphide, potassium sulphate is formed [TV/. Schlagdenhafen, Bulletin de la Societe Chimique (2), 22, 16 ; and Jour. Chern. Society, 28, 912]. 4K2Mn2Os + 3K2S = 3K2S04 +• 4K20 + BMn02. This method he uses quantita- tively for the estimation of hydrosulphuric acid. With some dilute free acid, such as sulphuric, hydrosulphuric added in excess to potassium permanganate gives manganous sulphate and free sulphur. H2Cr04 0. Forms chromic oxide and sulphur. Ni2Q3 p. Becomes nickelous sulphide and sulphur. 0°2O3 q. Forms cobaltous sulphide and sulphur. (Fe2)YI r. Forms Fe” and sulphur. The action takes place in either alkaline or acid mixture. THIOSXJLPHURIC ACID. H2S203. Oxidation valence H'2S"sO"8 O S „ il il Structural valence H-O-S-S-H, or H-O-S-O-H* il n O O 786. Thiosulphuric acid (formerly called hyposulphurous acid) has not been isolated. Thiosulphates are made by boiling sulphur in a soluble sul- * In union with oxygen the chemical polarity of sulphur is positive; in union with hydrogen, sulphur is negative; in union of sulphur with sulphur, the polarities of the one side are neutralized by those of the other side, so that in this union the total polarity of the sulphur stands at 0. See paragraph 614. Svi + S " = (Sa)iv. A. B. P. Thiosulphuric Acid. 231 phite (a), or in a soluble hydroxide (b) ; also by exposure of the persul- phides of the alkalies and alkaline earths to the air (c). The carbon diox- ide and oxygen of the air act upon the normal sulphides of the fixed alka- lies to produce a thiosulphate (d). а, 2Na2S03 S3 2NA25203 б. 3Ca(OH)2 + 652 = 3CaSS + CaS203 -I- 3H20 c. 2CaSS + 302 = 2CaS2O3 + 3S2 d, 2Na2S -(- 203 d- CO2 Na2S203 -f- Na2CO3 When thiosulphates are decomposed by acids, the constituents of thiosulphuric acid are dissociated as sulphurous acid and sulphur. Nearly all acids in this way decompose thiosulphates: The larger number of the thiosulphates are soluble in water; those of barium, lead, and silver being only very sparingly soluble. The thiosulphates are insoluble in alcohol. They are decomposed, but not fully dissolved, by acids, the decomposition leaving a resi- due of sulphur. In analysis, thiosulphates are distinguished by giving a precipitate of sulphur with evolution of sulphurous anhydride when their solutions are treated with hydrochloric acid (789); by their intense reducing power (790), shown in the blackening of the silver preci- pitate (788); and by non precipitation of calcium salts. 787. Alkali thiosulphate solutions dissolve the thiosulphates of lead, silver, and mer- cury ; also, the chloride, bromide, and iodide of silver, and mercurous chloride ; the iodide and sulphate of lead; the sulphate of calcium, and some other precipitates—by formation of soluble double thiosulphates: 3Na2S203 + 4HCI = 4NaCI + 2H2503 + S2 AgoS203 + 2NaAgS‘j03 Or .* AgnSaOs -4- SNaoSoOs Na4Ag2(S203)3 AgOl + Na25203 = NaAgS.Os + NaCl PbSO* + 8Na25203 = Na4Pb(S203)3 + Na2SO« 788. Barium chloride forms, in solutions of thiosulphates, a white pre- cipitate of barium thiosulphate, BaS203, nearly insoluble in water ; dissolv- ing in acids, except the sulphur residue.—Calcium chloride forms no pre- cipitate (distinction from sulphite).—Solutions of silver nitrate (420), lead acetate, and mercurous nitrate form at first white precipitates of thiosul- phates, soluble in excess of alkali thiosulphates, as stated in the preceding paragraph. These white precipitates, by standing, or quickly by warming, turn darker and finally black, by formation of sulphides, with sulphuric acid. Ag2S2Oj -f HoO = Ag2S + H2SO4 PbS2O3 + HoO = PbS -f HOSO4 Solution of copper salts, with thiosulphates, on long standing, precipi- tate cuprous salt, changed by boiling to cuprous sulphide and sulphuric acid, as above. For the precipitation of sulphides of arsenic, antimony, and, in the cold, tin, see 563. 789. The precipitation of sulphur with evolution of sulphurous anhy- dride, by addition of dilute acids—as hydrochloric or acetic—is charac- teristic of thiosulphates. It will be understood, however, that in pres- 232 Dithionous Acid. ence of oxidizing agents, which can be brought into action by the acid, sul- phides will likewise give a precipitate of sulphur (780). 790. Thiosulphates are Reducing Agents—even stronger and more ac- tive than the sulphites, to which they are so easily converted. This reduc- ing power is exemplified by the conversion of ferric salts into ferrous salts (a), and by the bleaching of iodine solutions (b), both of which changes are so sharply defined that they are useful in volumetric analysis. If the ferric solution be made red by addition of a few drops of thiocyanate, the exact point of complete reduction is made obvious: while the inevitable color of free iodine is nearly sufficient to mark the point when loss of color shows that all the iodine has entered into combination, but the addition of starch-paste renders the indication more exact. In both these reactions, the oxidation of the thiosulphate changes it into a tetra- thionate, as Na25406 (SO4 + 6KCI + 5H20 HYDROCHLORIC ACID. HCI. 835. Oxidation valence.... H'd-' Structural valence.... H-Cl The vapor density (H •= 1) is 18.185, showing that the molecule is HCI. 248 Hydrochloric Acid. HCI gas under ordinary atmospheric pressure liquefies at -102° 0., and solidifies at -115.7° C., melts again at -112.5° C. (Olszewski, 1884). Under a pressure of 40 atmospheres it liquefies at 10.6° C. (Mitchell). Occurrence.—Not found native except in volcanic gases and in springs and rivers in the vicinity of volcanoes. Found as a chloride in many minerals, its most abundant source being sodium chloride. 836. Preparation.—(l) For commercial purposes always made by action of H.2S04 on NaCl. Sulphuric displaces hydrochloric acid from all chlo- rides. Exception : It has no action on Hg2Cl2 and HgCl„, and the chlorides of silver, lead, and tin are but imperfectly transposed. (2) By direct combination of H and Cl by means of heat, light, or electricity. (3) The chlorides of the first two groups and of Zn, Pe, Co, and Ni are re- duced to the metallic state when heated in hydrogen gas. HCI being evolved. 1. 2NaCI + H2SO, = Na2SQ4 + 3HCI 2. H2 + CI2 = 2HCI 3. CuCl2 + H2 = Cu + 2HCI 837. Absolute hydrochloric acid is a colorless, caustic, suffocating gas. It dissolves in about two parts by weight of water. At 15° C. a saturated solution having a specific gravity of 1.2124 contains 43.09 per cent, of HCI (J. Kolb, 1872). Its strongest permanent solution contains about 33 per cent, of ncid (HC1); but this solution rapidly evolves acid in the air, more rapidly on warming, less rapidly as the solution loses strength. When of 20 per cent, acid, the liquid boils at 112° C. (233° F.), vaporizing with the water of its solution. The U. S. P. acid contains 31,9 per cent, of HCI, and its sp. gr. is 1.16. 838. Gaseous hydrochloric acid escapes with slow effervescence when liberated from compounds in concentrated solution ; reddens litmus; with vapor of ammonia, gives a white cloud (NH4CI as a solid), somewhat more dense than the fumes caused by the other volatile acids ; and, like aqueous hydrochloric acid (842), precipitates chlorides from salts of the first-group metals, when brought in contact with their solutions—a drop adhering to a glass rod being held in the gas. 839. Hydrochloric acid is formed from metallic chlorides by transposi- tion with sulphuric acid; except that silver, lead, and tin chlorides are transposed with difficulty, and mercurous and mercuric chlorides not at all, by sulphuric acid : 3NaCI + H2SO4 = Na2S04 + 3HCI 840. The normal Chlorides are all soluble in water, except those of the metals of the first group; silver, AgCl; mercurous, Hg2Cl2 ; and lead, PbCl2—the last named being sparingly soluble. In analysis, the silver ]ire- cipitate is most used (843). If bromides are present, the chlorochromic test is most conclusive (845). A large number of the metallic chlorides are soluble in alcohol, and .several are soluble in ether. JIYDR 0 CHL ORIV A CID. 841. The chlorides of metals are, generally, more volatile than other compounds of the same metals ; example, ferric chloride. 842. Solutions of chlorides and hydrochloric acid are precipitated, by solutions of silver salts, as silver chloride, AgCl, white (413) ; by solutions of mercurous salts, as mercurous chloride, Hg2Cl2, white (430); and by solutions of lead salts, when not very dilute, as lead chloride, PbCl2, white (390). Silver nitrate solution is the most complete and convenient pre- cipitant for chlorides. Exceptions: HgCl„ does not precipitate lead salts. The chlorine in green Cr2Cl6 is incompletely precipitated by AgW03 (Peli- GOt) ; whilst from a solution of oxychloride of molybdenum in H2S04 it is not precipitated at all (Blomstrand). 843. The properties of the precipitate of silver chloride are given in 413 and in 409. It is of analytical interest that it is freely soluble in am- monium hydroxide (considerably more freely than the bromide, and far more freely than the iodide of silver); soluble in hot, concentrated solution of ammonium carbonate (which dissolves traces of bromide, and no iodide of silver); insoluble in nitric acid, or but temporarily soluble in strong nitric (as in hydrochloric) acid, and precipitating again on dilution. It should be observed, that it is appreciably soluble in solutions of chlorides, and in ammonium nitrate ; hence, in reprecipitating traces of it, by nitric acid, from the ammonia solution, if there is excess of ammonium hydroxide, this should first be expelled : 2AgCI + 8NHtOH = (NH3)3(AgCI)2 + 3H.0 (NH3)3(AgCI)2 + 3HNO3 = 2AgCI + 3NH4NOs 844. Oxidizing agents (with heat) decompose hydrochloric acid. The action of manganese dioxide is formulated as follows : 4HCI + MnOo = MnCh + 2H,0 + Cl 2 Or: 2NaCI + 3H2504 + MnO, = Na2S04 + MnS04 + 2ERO + Cl2 For more complete statement of the oxidation of HCI sec 851. This reaction is applied in the manufacture of chlorine, which is dis- tilled from the mixture, and can be used in analysis for evidence of chlo- rides. A test for traces of free hydrochloric acid, in distinction from metallic chlorides, is made by heating the solution with MnG2, without adding an acid, and distilling into a solution of potassium iodide and starch. Larger proportions of HCI are more frequently separated by distilling it intact, 845. The reaction with chromic anhydride is in use as a test for hydro- chloric acid, more especially in presence of bromides : a. 2HCI + Cro3 = CrO2Cl2 (clilorochromic anhydride) + H2O b. 4NaCI + K2Cr207 + 3H2S04 = 2Cr02C12 + 2Na2SC>4 + K2SQ4 + 3H20 To obtain a rapid production of the gas, so that it may be recognized by it; color, the operation may be made as follows : Boil a mixture of solid potassium dichromata and sulphuric acid, in an evaporating-dish until bright red, and then add the substance 250 H YDR 0 CIIL DRW A CID. to be tested, in powder—obtained, if necessary, by evaporation of the solution. If chlo- rides are present, the chromium dioxy-dichloride rises instantly as a bright brownish-red gas. The distinction from bromine requires, however, that the material, which may be in solution, should be distilled, by means of a tubulated flask or small retort, the vapors being condensed in a receiver, and neutralized with an alkali (c and d). The chro- mate formed makes a yellow solution (bromine, a colorless solution). As conclusive evi- dence of chlorine, the chromate (acidified with acetic acid), with lead acetate, forms a yellow precipitate (bromide, a white precipitate, if any): c. CrO,Cl2 + 2H20 = H2Cr04 + 2HCI d. CrO2Cl2 + 2(NH,)OH = iNH4)sOrO4 + 2HCI 846. The action of nitric acid with hydrochloric acid results from the mixture of the two acids, well known as nitro-hydrochloric acid, or “aqua regia,” and used for its free chlorine. Both nitrogen oxychloride and ni- trogen oxydichloride are formed ; their relative proportion varying with different conditions : ( NOCI2 + Ol ) 3HCI -f HNOs = < V + 2H20 ( or NOCI + 2CI ) The reaction occurs quite promptly in the concentrated acids without heat, but more rapidly with heat ; very slowly in moderately dilute acids, and only to a slight extent if the acids are very dilute. 847. The three chlorides insoluble in water (840) are not transposed or dissolved by acids, except that mercurous chloride is dissolved, by nitric acid and by chlorine, as mer- curic salt. They are dissolved for analysis by decomposition with alkali hydroxides. AgOl or PbCl2 is dissolved as sodium chloride, after fusion with sodium carbonate on charcoal (a), or in a porcelain crucible (b): a. 4AgCI + 2Na.CC>, + C = 4Ag + 4NaCI + 3CC2 h. PbCl2 + Na2C03 = PbO + SNaCI + C02 848. Heated in a bead of microcosmic salt, previously saturated with copper oxide in the inner blow-pipe flame, chlorides impart a blue color to the outer flame, due to copper chloride. 849. Estimation.—(l) It is precipitated by AgNOa, washed, and, after ignition, weighed as AgCl. (2) By a standard solution of AgN03. A lit- tle Na2HP04, or, better, K2Cr„O., is added to the chloride to show the end of the reaction. When enough AgW03 has been added to combine with the chlorine the next addition gives a yellow precipitate with the phosphate, or a red with the chromate. 850. Chlorides may be made, a, by direct union of the elements, mostly without heat. Whether an ous or ic salt is formed depends upon the amount of chlorine used. i. By the action of hydrochloric acid upon the corresponding oxides, hydroxides, carbonates, or sulphites. The solutions formed may be evaporated to expel excess of acid. If the chlorides thus formed contain water of crystallization it cannot be removed by heat alone, for part of the acid is by this means driven off, and a basic salt remains. If the anhydrous chloride is desired, it may always be made by a, and when Hypochlorous Acid. 251 thus formed may be sublimed without decomposition, c. Chlorides of the first group are best made by precipitation. Exception : Mercuric chloride does not precipitate lead salts, d. Metals soluble in hydrochloric acid evolve hydrogen and form chlorides. In these cases ous, and not ic, salts are formed. e. Many chlorides may be formed by bringing HgCl2 in contact with the hot metal. 851, Oxidation.—For some properties of hydrochloric acid consult 834 i. For action on HNOs see 846. HCIO a. Forms chlorine and water. HCIO + HCI = Cl2 -f- H2O. HCI03 h. Gives either free chlorine or a varying amount of the several ox- ides of chlorine [Storer’s Quantitative Analysis, 119]. HBr03 c. Forms free bromine and chlorine. HI03 d. In dilute solution, no action. If concentrated, yellow chloride of iodine, IC13, is formed, not taken up by carbon disulphide. On boiling, some free iodine is separated, which colors the carbon disulphide. e. In all cases where metals are dissolved in hydrochloric acid, hy- drogen is evolved. In such cases it is an oxidizing agent, the chlorine of the acid not changing its bonds, but the hydrogen is changed from combined, in which it is plus, to free hydro- gen, where it has no bonds, thereby losing one bond. Or, as explained in 614, (H2)° =(H + /H~')°. For its action on oxides see 834 i. Pb304 and Pb02 /. Form plumbic chloride and chlorine. Mn" +n g. Becomes In" and (Cl2)° is formed. That is, all compounds of manganese having more than two bonds are reduced to the dyad, with evolution of chlorine. In case of dilute K2Mn208 this change is preceded by formation of manganese peroxide [S. U. Pickering, Jour. Chem. Society, 35, 654]. Coand Hi''' h. Are changed to cobalfous and nickelous chlorides, with evolution of chlorine. PeVI i. With the exception of ferrates, as K2Fe04, which forms ferric chlo- ride, the compounds of iron are not reduced. CrVI. Forms chromic chloride, and chlorine is evolved. HYPOCHLOROUS ACID. HCIO. 852. Oxidation valence H'Cl'O-" Structural valence H-O-Cl Vapor density of the anhydride, C120, 43.5. Hypochlorous anhydride. Cl2O, can be obtained at -20° C. (-4° P.), by the reaction, HgO + 2C12 = C120 + HgCl2, as an orange-colored, explosive liquid, gaseous at ordi- Hvp o cjil op ous A cm. nary temperatures, and decomposing, spontaneously and sometimes violently, into chlo- rine and oxygen. It dissolves in water, forming hypochlorous acid, HCIO. Hypochlorous acid, in aqueous solution, is a yellow liquid; when strong, decompos- ing rapidly at 0° C. (32° F.); when dilute, decomposing gradually by boiling (a); decom- posed by hydrochloric acid (&), and by metals (r)—its decompositions furnishing chlorine or oxygen, or both chlorine and oxygen. a. 4HOIO = 2H20 + 2C12 + 02 b. HCIO + HCI HoO + Cl 2 c. 2HCIO + 2Zn = ZnO + ZnCl2 + H2O 4HCXO + 4Ag = 4AgCI + 2H20 +O2 853. The Hypochlorites at e formed by treating bases with chlorine (short of satura- tion), as shown in b, 833. The calcium hypochlorite and chloride, mixed or combined together as formed by action of chlorine upon calcium hydroxide, is in very extensive use —as chlorinated lime, or “ chloride of lime ” [CaCI2 Ca(ClO)2, or 2Ca(OCI)CI]. The sodium hypochlorite-and chloride—mixed as formed by chlorine in solution of sodium hydrate or sodium carbonate, or by double decomposition between solution of the calcium Jiypochlorite-and-chloride and solution of sodium carbonate—is pharmacopoeial, under the name of solution of chlorinated soda (NaCI.NaCIO). The chemical structure of these important chlorinated compounds has been difficult to ascertain. 854. Hypochlorites are very instable, whether solid or in solution, de- composing by the weakest acids, by the carbonic acid of the air (a), and by heat (b), also to some extent at ordinary temperatures. In this manner, they act as powerful oxidizing agents. The deportment of hypochlorites is represented by the action of chlorine in alkali solutions (833 Z>); a con- venient agent of especial force, as for the decomposition of ammonia (45 b). a. A. 2Ca(dO)2 + 2C02 = 2CaCO3 + 2C12 + 02 B. CaCl2.Ca(CIO)2 + 2C02 = 2CaCO, + 2C12 c. Ca(CIO)2.CaCI2 + 2H2S04 = 2CaS04 + 2HsO + 2C12 D. 4NaCIO + 4HCI = 4NaCI 4- 2H20 + 2C12 +O2 b. In boiling solutions: 3Ca(CIO)2 = Ca(ClO3)2 + 2CaCl2 At a higher temperature: Ca(ClO3)2 = CaCl2 + 302 All hypochlorites are soluble in water; those of the first-group metals being especially instable. Their solutions bleach litmus and indigo. Silver nitrate, added to the solutions of hypochlorites with chlorides, precipitates the chloride, AgCl, at first leaving hypochlorites in solution ; while the soluble silver hypochlorite quite rapidly decomposes with the pre- cipitation of chloride and formation of chlorate of silver (a), the latter slowly changing to chloride. a. SAgCIO = 3AgCI + Agol03 (corresponding to b, 854) Oxidation.—The oxidizing power of this acid is, in general, the ssme ns that of free chlorine. Chlorous Acid—Chlorine Peroxide—Chloric Acid. 253 CHLOROUS ACID. HCI02. 855. Oxidation valence H'Cl/"0~"2 Structural valence H-O -Cl=o Vapor density of the anhydride, C1203, experimental (H = l), 58.4. The anhydride is prepared, according to Millon, by mixing 15 parts As.,03, 20 parts KCI03, 60 parts HN03, sp, gr. 1.33, and 20 parts water; the temperature is cautiously kept at 25° to 45° C. ; the C1203 gently dis- tills over, and is received in a flask partly filled with water. It should be observed that the amount of As2Oa taken is limited ; a larger amount would reduce the HCI0S to HCI. Dilute H„S04 may be used instead of HNOs, and instead of As2Q3 De Vry uses tartaric acid, Schiel uses cane sugar, and Carius uses benzene ; in which cases C02 is also formed. The greenish- yellow anhydride passed into water forms the reddish-brown acid, and by treating this with hydroxides a few salts have been prepared—e.g., KC10„, NaCl02, AgCl02, Pb(ClO2)2, Ba(01O2)2, Sr(C102)?. The anhydride ex- plodes at 57° C., forming oxygen and chlorine. Most chlorites are soluble; silver chlorite sparingly. The acid decolors K2Mn2Og, and bleaches indigo solution. If a slightly acidulated dilute solution of a ferrous salt is mixed with a dilute solution of chlorous acid, the liquid transiently acquires an amethyst tint, and does not assume the yellowish color of ferric salts until the lapse of a few seconds (Leatsse^). CHLORINE PEROXIDE. C 102. 856. Oxidation valence C1ivO~"2 Structural valence... .not satisfactorily determined. Vapor density at 30° C. (H =l) is 33.65. Made by the action of H2S04 (a) or H2C204 (£) upon KCI03. It is a dark yellow gas, which condenses by cold to an orange-colored solid, which melts at -76° 0, (Faraday), and under 730.9 m.m. pressure boils at 9.9° 0. (Schacherl, 1881). It decom- poses explosively by heat, light, or by contact with reducing agents, such as P, S, sugar, ether, and turpentine. Passed into KOH forms KCI02 and KCI03 (c) (Millon). a. SKCIOa + 2H2504 = 2KHSO, + ECI04 + H2O + 2C102 b. 2KCI03 + 2H20204 = K20204 + 2H20 + 2C102 + 2COa c. 2CI02 + 2KOH = KOlO2 + KCIO3 + H2O CHLORIC ACID. HCI03. Oxidation valence H'C1V 0~"a O II Structural valence H-0-Cl=o 857. Preparation.—From KCI03, by adding H2SiF6 (a). Or from Ba(C103)2, by adding exactly the required quantity of H2S04 (d). The Chloric Acid. chlorates of the alkalies and alkaline earths are made by action of excess of chlorine upon a hot solution of the hydroxides (c). Other chlorates are formed by action of Ba(C103)2 upon the respective sulphates (d), or of AgCl03 upon the resj)ective chlorides (e). And all may be formed by action of the free acid upon the corresponding hydroxides or carbonates. The mercurous and ferrous salts are very instable, and those of antimony, tin, bismuth, and manganese are not with certainty known. The anhy- dride, C120,., has not been isolated. a. 2KC103 + HoSiPc = K2SiF6 + 2HCI03 b. Ba(C103)2 + H2SG4 = BaSO4 + 2H0103 c. 6KOH + 3C12 = KOl03 + SKCI + 3H20 d. CdSOi + Ba(C103)2 = BaSQ4 + Cd(0103)2 e. 2AgCI03 + ZnCl2 = 2AgCI + Zn(ClO3)2 858. Properties.—The aqueous solution may be concentrated in a va- cuum until its sp. gr. is 1.282 and contains 40.1 per cent, of HCI03. It is a colorless, syrupy liquid, having a slight odor, resembling nitric acid. It first reddens litmus, and then blenches it. Chloric acid is somewhat in- stable at ordinary temperatures ; when heated, it rapidly decomposes with formation of yellow products, including perchloric acid (HCIOJ, chlorine and water. It oxidizes organic and other combustible substances with violent rapidity. All chlorates are resolved by heat into chlorides and oxygen (2KC103 = 2KCI BOa). In presence of various metallic oxides, etc., the oxygen is separated more easily, the metallic oxides remaining unchanged. With manganese dioxide, the oxygen of potassium chlorate is obtained at about 200° C.; ferric oxide, at 120° C.; platinum, black, at 270°. Copper oxide and lead dioxide may be used. When triturated or heated with combusti- ble substances, charcoal, organic substances, sulphur, sulphites, cyanides, thiosulphates, hypophosphites, reduced iron, etc., etc.—chlorates violently explode, owing to their sudden decomposition, and the simultaneous oxida- tion of the combustible material. This explosion is more violent than with corresponding mixtures of nitrates (as in gunpowder, 730). 859. All the chlorates are soluble in water ; those of the first-group bases being somewhat instable in solution. Like nitric and acetic acids, chloric acid is not precipitated. Except the mercurous, the least soluble of the metallic chlorates is that of potassium, which requires 12 to 16 parts of cold water for its solution. Potassium chlorate is only slightly soluble in alcohol. Chlorates are usually identified by the gaseous products of decom- position (860). 860. Sulphuric acid causes dissociation of chlorates—if in the solid state, with detonation, and, unless in small quantities, with violent explo- sion ; and with formation of greenish-yellow gas, chlorine peroxide. Per- chlorate is likewise formed. The gas has the odor, oxidizing and bleaching Chloric Acid. 255 power of chlorine, and it imparts its color to solutions in which it is formed. The products vary with conditions, but are chiefly formed as follows : 3KCI03 + 2H2S04 = 2KHS04 + KCI04 + 2C102 + H2O 4C102 = C 1'"203 + C1v205 If a dilute solution of a chlorate is colored light blue with the solution of indigo in sulphuric acid, and the solution kept cold, no bleaching occurs, even with the further addition of dilute sulphuric acid. But, on addition of solution of sodium sulphite, the color soon disappears, by formation of chlorine or its oxides. 861. Hydrochloric acid decomposes chlorates, rapidly when heated, with the forma- tion of free chlorine and chlorine peroxide—the mixture called euchlorine. The gas and solution have the color, odor, and bleaching effect of chlorine, intensified. This is a ready and effective means of generating chlorine for analytical purposes. The propor- tion of free chlorine to oxidized chlorine is variable, the subjoined equations showing the character of the results. According to Bunsen, sometimes only free chlorine is pro- duced; at others, only C 120; at others, both gases are evolved, and occasionally Cl2Os is formed. 2KC103 + 4HCI = Cl2O + Cl 2Oa + 2H20 + 2KCI 2KCI03 + 4HCI = 2KCI + 2CI02 + Cl 2 + 2H.0 BKCIO3 + 10HC1 ~ 3KCI 4- 2C102 + 4C12 + 5H20 4KCI03 + 16HC1 = 4KCI + 2C102 + 7C12 + 8H20 KClOa + 6HCI = KCI + 3C12 + 3H20 2KC103 -t- GHCI = 2KCI + 3C120 + 3H20 2KC103 + BHCI = 2KCI + 2C120 + 2C12 + 4H20 If a solution of a chlorate be acidified with dilute sulphuric acid, and zinc added, the liberated chloric acid is reduced to hydrochloric acid by the nascent hydrogen. If a dilute solution of an alkaline chlorate is boiled with a copper-zino couple,* it is completely reduced to chloride with separation of oxide of zinc (Thoepe and Eccles). Chloric acid behaves like nitric acid with brucine (Luck), diphenyla- mine, paratoluidine, and phenol, dissolved in concentrated sulphuric acid, or at all events the reactions are so similar that the two acids cannot be distinguished with certainty by these reagents. Chloric may, however, be distinguished from nitric acid by its action on phenol in hydrochloric acid solution (compare 738), inasmuch as the chloric acid produces an orange-red turb i d i ty. 862. Estimation.—(l) Reduce the chlorate to a chloride by dilute H„S04 and Zn, or if it is an alkaline chlorate, by fusion. It may then be estimated as a chloride. (3) Add HCI and KI, and determine the liberated iodine by standard solution of Na„S203. Which of the products of decom- position given above may actually be formed, whether all or only certain of them, cannot be foreseen. But no matter which of them may be formed, they all of them agree in this, that, in contact with solution of potassium iodide, they liberate 6 atoms of iodine for every molecule of HCI03. * Gladstone and Tribe’s copper-zinc couple is prepared by treating thin zinc foil with a 1 per cent, solu- tion of copper sulphate until the zinc is covered with a black deposit of reduced copper. When washed and dried it is ready for use. 256 Chloric Acid. 863. Oxidation.—Alkali chlorates when fused with, an alkali, or an alkali carbonate, and a free metal or a lower oxide, or salt of the metal, generally oxidizes it to a higher oxide, or to a salt having an increased num- ber of bonds ; and the chlorate is reduced to a chloride—e.g., MnVI - n be- comes MnVI. That is, any compound of manganese having less than six bonds is oxidized to the hexad («). In equations (b) and (c) the sulphur is also oxidized. Cr”' becomes CrVI (d). Sbv ~ n becomes Sbv (e). Asv_'1 becomes Asy (/). PbIV - n becomes PbIV (g). Pe” ~ n becomes Pe”' (A). Co"'~n becomes Co'” (i). Also CIY ~n becomes JIV (y). Pv~w becomes Pv (k). Iv - n becomes Iv {I). SVI —w becomes SVI (m). a. 8M11304 + 18KOH + OKCIO3 = OK2MnO4 + SKCI + 9H20 b. MnS + 3K2C03 + 2KCI03 = K,MnQ4 + K2S04 + 3KCI + 3C02 c. MnS2Oe + 6KOH + KCIO3 = K2MnQ4 + 3K2S04 + KCI + 3H2Q d. Cr2Cl6 + IONaOH + NaCl03 = 2Na2CrQ4 + 7NaCI + 5H20 e. 6SbCI3 + 18KOH + 3KCI03 = 3Sb203 + 30KCI + 9H20 /. 3As4 + 36KOH + lOKCIOs = 12K3AsO4 + 10KC1 + 18H20 g. 3Pb304 + Na2003 + 2NaCI03 = 9PbO2 + 2NaCI + Na2C03 h. 6FeSO3 + 13KOH + 3KC103 = 3Pe203 + GK2SO4 + 3KCI + GH2Q i. 6C0012 + 12KOH + KCI03 = 30o203 + 13KC1 + 6H2Q j. 3MnC4H406 + 3K2C03 + 7KCI03 = 3K2MnG4 + 7KCI + 15C02 + 6H2Q k. 3Pb(H2PO2)2 + 18KOH + SKCI03 = 3PbQ2 + 6K3P04 + SKCI + 15H2Q l. Znl2 + K2COs + 2KCI03 = ZnO + 2KIOs + 2KCI + C02 m. 3K25506 + 13K2003 + 10K0103 = 15K2S04 + 10KC1 + 12C02 Free chloric acid is a strong oxidizing agent, and if an excess of the reducing agent is used, it is converted into HCI, or a chloride, but if heat be employed the chloric acid splits up and some free chlorine, or its oxides, may be formed. For its action on acids already given, see H202O4, 659 ; H3P02, 756 ; HONS, 702 ; H4Fe(CN)6, 691 ; H6Fe2(CISr)12, 694; H2S, 785 ; H2S03, 813 ; HCI, 861. HBr a. Forms bromine. I b. Forms HIO,, HI c. Forms first 1° then Iv (free iodine and iodic acid). Hg' d. Forms Hg". Sn" e. Forms SnIv. Sb"' /. Forms Sbv. As”' g. Forms Asv. Pe" h. Forms Pe'". Cu' i. Forms Cu", Per chloric A cid— Bromine. 257 PERCHLORIC ACID. HCI04. 4 Oxidation valence H'C1viiO~"4 O II Structural valence H-0-Cl=o / ii O 864. Potassium perchlorate is first formed by heating KCI03 to about 350° C. (a), and, after separating the chloride by washing, it is treated with concentrated H2S04 and distilled, the HCI04 passing into the receiver (b). Or it may be made by evaporating a solution of chloric acid (c). a. 2KC103 = KCIO4 + KCI + 02 b. 3KCI04 + H2S04 = K2SO* + 3HC104 c. BHCIO3 = 4HC104 + 3H20 + SOL. + 802 The pure perchloric acid is a colorless, very heavy liquid (sp. gr. 1,782), which soon becomes yellow from decomposition. It cannot be kept for any length of time. When heated, it undergoes decomposition, often with ex- plosion. In its oxidizing properties it is more powerful than chloric acid. It burns the skin in a very serious manner, and sets fire to paper, charcoal, etc., with explosive violence. This want of stability, however, belongs only to the pure acid. If water be added to it heat is evolved, and a dilute acid of a far greater permanence is obtained. Diluted perchloric acid does not even bleach, but reddens litmus in the ordinary way. Hydrochloric acid, nitric acid, and sulphurous acid do not decompose aqueous solutions of perchloric acid or perchlorates ; solution of indigo, therefore, previously added to it, is not decolorized (difference from all other acids of chlorine). The alkaline perchlorates are not reduced by the cop- per-zinc couple (distinction from chloric acid, 861). All the metallic perchlorates are soluble in water ; potassium perchlo- rate requiring 58 parts (rubidium perchlorate, 92 parts) of water at 21° C., the other metals forming salts more freely soluble. The potassium salt is insoluble in alcohol. In ignition, perchlorates act very much like chlorates, but more explosively. BROMINE. Br = 79.768. Oxidation valence of free bromine (Br2)° Structural valence of free bromine Br-Br 865. The vapor density at 100° C. (H =l) is 80, showing that the molecule is Br2, but at 1570° C. (Meyer and Zublust, Berichte, 13, 405) it is 53.33, which would make the molecule about 2, the molecules undergoing partial dissociation into atoms. Solidifies at -7.2° 0. (Philipp, 1879). Boiling point, 59.27° C. (Thorpe, 1880). 258 Bromine. Occurrence.—lt is never found free, but usually as a bromide of K, Ha, or Mg. Found in sea-water, sea-weeds, and salt springs. Preparation.—(l) The mother-liquor from the salt wells is condensed by evaporation, and the greater portion of the HaCl, Ha2S04, MgS04 is separated by crystallization. By the old method the Naßr is then treated with chlorine («), an excess of chlorine being avoided, which would form bromine chloride. The solution is then shaken up with ether, which ab- sorbs the bromine. Then the ether solution is shaken with solution of KOH, which forms KBr and KBrC)3 (h). The ether is separated and is used again for the same purpose. The mixed KBr and KBr03 is then ignited (r;). It is then distilled with Mn03 and H2S04 (d). if any iodine is pres- ent it is removed by precipitation as cuprous iodide, and the last traces of chlorine are removed by adding KBr and redistilling (compare a). If the anhydrous bromide is required it is distilled from fused CaCl2. (2) A later method is to add the necessary Mn02 and H2S04 directly to the mother- liquor and distill off the bromine. The use of ether is omitted, KOH or NaOH is added to the distillate, which is then treated as in (1). a. 2Naßr + Cl2 = 2NaCI + Br2 b. 6KOH + 38r2 = KBrQ3 + SKBr + 3H20 ic. 2K8r03 (ignition) = 2KBr + 302 d. 2KBr + 2H2S04 + MnQ2 = K2SQ4 + MnS04 + 2H20 + Br2 866. BROMINE, at ordinary temperatures, is a brown-red, intensely 'caustic liquid, freely evolving brown-red, corrosive vapors of a suffocating, chlorine-like odor, and boiling at 47° 0. (117° F.) It is soluble in 33 parts ■of water, with an orange-yellow color ; and freely soluble in alcohol, ether, ■chloroform, and in carbon disulphide—with the same or a deeper color. Carbon disulphide, and chloroform, after agitation with its dilute water solutions, remove the bromine as a subsiding liquid in a yellow to red-brown layer ; ether, less perfectly, in a supernatant layer of the same color. It dissolves colorless in alkali hydroxides, with combination, forming bro- mides and bromates. The change corresponds exactly to that of chlorine, as shown in equation b, 833, and 865 b. It dissolves, without combination or loss of color, in solutions of hydrobromic acid, and of bromides. Its water solution is permanent, but sunlight decomposes it thus : 3Br2 + 3H.0 = 4HBr + 02 In vapor or solution, bromine bleaches litmus and indigo; colors starch- paste yellow ; precipitates silver salts, yellow-white, bromide and bromate (as by equation c, 833); and lead salts, white. Bromine decomposes hydro- sulphuric acid with separation of sulphur, and subsequent production of sulphuric acid ; changes ferrous to ferric salts, and (in presence of water) acts as a strong oxidizing agent. It displaces iodine from iodides, and is displaced from bromides by chlorine ; its character being intermediate between that of chlorine and that of iodine. Hydrobromic Acid. 259 The Acids of Bromine are : Hydrobromic acid, HBr~'. Hypobromous acid, HBr'O. Bromic acid, HBrv 03. 867. Estimation.—(1) The bromine is made to act upon Kl, and the iodine which is liberated is estimated by standard solution of Na2S2Os. (2) It is estimated by the amount of As2C3 which it oxidizes in alkaline solu- tion. (3) It is converted into KBr by H2S or H2S03, and then precipi- tated by AgNC3, and weighed as Agßr. 868. Oxidation.—Bromine as an oxidizing agent is intermediate in strength between chlorine and iodine. In presence of acids it never acts as a reducing agent. In oxidizing it becomes Br-'; that is, hydrobromic acid or a bromide. For its action upon H2C204, see 659 ; H3P02, 756 ; H2S, 785 ; H2S03, 813 ; Cl, 834 ; HCNS, 702 ; H4Pe(C]Sr)6, 691. SVI~n a. Becomes SVI in alkaline and acid mixture. HI i. Forms 1° and HBr, In presence of alkali hydroxides., forms a bromide and iodate Pb" and Pb304 c. In alkali mixture become Pb02. Hg"“» d. Forms Hg" in alkaline and acid mixture. That is, both metal- lic mercury and mercurous compounds are oxidized to mer- curic compounds. AsT-w e. Forms Asv in alkaline and acid mixture. That is, free arsenic [As° or (As4)°] and the triad arsenic are changed to the pentad. Sbv~ra f. Forms Sbv in alkaline and acid mixture. SnIT - n g Becomes SnIV in alkaline and acid mixture. MnIV ~n h. Becomes MnIV with alkalies, not with strong acids. Cr'" i. Becomes CrVI “ “ “ Ni" j. Becomes Ni'" Co" and Co304 Jc. Becomes Co'" with alkalies, not with strong acids. Pe'" ~n I. In acid mixture becomes TV". Pevi-n, m jn aikaiine mixture becomes PeVI (a ferrate). HYDROBROMIC ACID. HBr. Oxidation valence H'Br-' Structural valence H-Br 869. Yapor density (H —1), 39.1. Condenses to a liquid at-69° 0., and solidifies at -73° 0. (Faraday). Preparation.—Bromine does not combine with hydrogen, even in the sunlight ; but if the mixed gases are passed over hot platinum sponge, com- bination results. It is prepared : (1) By action of Br on P and subsequent Hydrobromic Acid. distillation. Either variety of P will answer, but the amorphous is safer. (2) By adding H3P04 to KBr and distilling, precautions being taken to prevent bumping. (3) By distilling two parts of Caßr2, two parts H2S04, and one part H2O (Bertrand, 1876). (4) By action of Br on H3P02 or on Ca(H2PO2)2 and distilling. (5) By action of Br on Ha2S03. (6) By heat- ing Br with paraffine, HBr distills over and the residue is chiefly carbon. (7) By adding to Baßr2 exactly enough H2S04 and decanting. (8) If made by action of H2S04 on KBr it generally contains free bromine. (9) By add- ing H„S to the bromides of the first and second group metals. (10) By adding Br to an excess of H2S solution, and removing H2S04, if any be pro- duced, by adding Ba(OH)2 or BaG03 and distilling. 1. P4 10Br„ + I6HOO = 4H3P0., + 20HBr 4. Ca(H2PC>2)2 + 4Br2 + 4H20 = CaH4(PO4)2 + BHBr 5. NaaSOs + Bra + H2O = Na2S04 + 2HBr 870 Properties.—HBr is a colorless gas, fuming in moist air and hav- ing a very irritating odor. Its aqueous solution is colorless and not decom- posed by exposure to the air. Water saturated at 0° C. contains 83.03 per cent, of HBr, and its sp. gr. is 1.78, very nearly HBr.H,O (Bureau). If a saturated solution is boiled, chiefly HBr is given off, and if a dilute solu- tion is boiled, chiefly H2O is given off. until in both cases the remaining liquid contains 47.38 to 47.86 per cent, of HBr, its sp. gr. 1.485, its boiling point constant at 136° 0., and its composition almost exactly HBr.sH20, which distills over unchanged. Its vapor density of 14.1 agrees with the calculated vapor density of HBr.sH20. 871. The solubilities of the bromides lie intermediate between those of the chlorides and those of the iodides, not differing much from the former. In general terms, all bromides are soluble in water, except those of the first- group bases. Further, mercuric bromide is only sparingly soluble in water. Lead bromide is less soluble than lead chloride. Bismuth bromide is de- composed by water, to a greater extent than bismuth chloride (364), and antimonious bromide is decomposed by water. Cuprous bromide is formed as a precipitate by reduction from the soluble cupric bromide. The double bromides of lead and potassium or sodium, and of silver and potassium or sodium, are soluble in a little water containing alkali bro- mides as concentrated solutions, but are decomposed by much water; the potassio and the sodio mercurous and mercuric bromides are soluble in water. In alcohol, the alkali bromides are sparingly or slightly soluble ; calcium bromide, soluble ; mercuric bromide, soluble ; mercurous bromide, insolu- ble. Silver bromide is sparingly soluble in ammonium hydroxide ; nearly in- soluble in ammonium carbonate solution. In analysis, bromides are usually identified by the color of a carbon- Hydrobromic Acid. 261 disulphide solution of free bromine, iodine, if present, being first removed (873). 872. Silver nitrate solution precipitates, from solutions of bromides, silver bromide, Agßr, yellowish-white in the light, slowly becoming gray to black. The precipitate is insoluble in, and not decomposed by, nitric acid, sparingly soluble in concentrated aqueous ammonia, nearly insoluble in concentrated solution of ammonium carbonate, slightly soluble in excess of alkali bromides, soluble in solutions of alkali cyanides and thiosulphates. It is decomposed by chlorine. Solution of mercurous nitrate precipitates mercurous bromide, Hg2Br2, yellowish-white, soluble in excess of alkali bromides. Solutions of lead salts precipitate, from solutions not very dilute, lead bromide, Pbßr2, white (391 and 871). 873. Sulphuric acid decomposes all bromides, except those of silver and mercury: when dilute, mostly with production of hydrobromic acid; when concentrated, chiefly with formation of bromine. The vapor from the hot mixture reddens or bleaches litmus ; has the yellowish-brown color and suffocating odor of bromine, and when cooled colors starch-paste yellow. Chlorine-Water separates the bromine much more quickly and com- pletely, giving better results in dilute solutions, but in excess it decolors the bromized starch. The more delicate test is made by adding carbon disulphide,* then dilute chlorine-water, drop by drop, in the cold solution; then agitating, and allowing the heavier liquid to subside (866). The presence of bromine is indicated by a yellow color, or if there is much bromine a yellowish-brown to brownish-red color. lodine colors violet If free iodine is' present, bro- mine cannot be identified, by its vapor, its color with starch, or its color in solution with carbon disulphide. All the iodine of iodides will be liberated before any of the bromine can be : therefore, before these tests can be made for bromine, the iodine must either be oxidized to iodic acid, or wholly ex- pelled. Dilute hydrochloric acid will not color dilute solutions of bromides, or, in absence of oxidizing agents, yield color to disulphide of carbon. Bromides of potassium, sodium, and of most other metals, are not decomposed by ignition Silver bromide melts undeeomposed ; but is slowly reduced, and blackened, in the air and by light. Tested in the cupric bead, according to 848, bromides give a greenish-blue color to the outer flame—not very marked. * Carbon disulphide is a better color solvent, for bromine or iodine, than chloroform, and far better than ether. It must be free from sulphurous or sulphuric acids. Saturated chlorine-water is liable to act on carbon disulphide, giving it a yellow color, simulating bromine. On adding alcohol to this yellow liqu'd sulphur precipitates. Hence the direction to use dilute chlorine-water, and avoid excess. Hypobromous Acid. 874. Metallic bromides are formed : (1) By direct union of the elements, but in a few cases heat is required to effect the combination. (3) By action of HBr upon the metallic oxides, hydroxides, and carbonates. (3) Many bromides are formed by action of HBr on the free metal, ous salts and not ic being formed. (4) Bromides of the first group are best made by precipitation. (5) Bromides of K, Na, Ba, Si, and Ca are made by the action of bromine on their hydroxides and subsequent fusion. 6KOH + 38r2 = KBrOs + SKBr + 3H2Q 2KBrOs (ignited) = 2KBr + 302 875. Estimation.—Gravimetrically, it is precipitated by AgH03, and, after ignition, weighed as Agßr. (3) Volumetrically (in same manner as HCI) by standard solution of AgWC3, using K2Cr2C7 to show when the re- action is complete. (3) It may be oxidized to free bromine and estimated as in 867. 876. Oxidation.—a. The bromine of HBr can never act as an oxidizing agent, for the reason that it is already reduced to its very lowest state of oxidation. When zinc is oxidized by HBr the change in bonds may be ex- pressed thus : Zn° + 3H'Br~' = Zn"Br‘2 + (H2)°. That is, the zinc gains the two bonds lost by the two atoms of hydrogen, and the bonds of the bro- mine remain unchanged. When HBr acts as a reducing agent free bro- mine is formed, except that in alkaline mixture a bromate is produced. H2S04 b. Dilute, no action ; concentrated becomes S02. Cl c. In acid mixture forms HCI and Br°, but in alkaline mixture forms a chloride and bromate. HCIO e. Same as free chlorine. HCI03 f. Forms chiefly HCI. HBrOa g. Both acids are changed to free bromine. HIOs h. Forms 1° and Br° (free iodine and free bromine). Pb" + n i. Becomes Pb Sbv j. Becomes Sb"'. Biv ic. Becomes Bi'". Mn"+re /, Becomes Mn". Co'" m. Becomes Co". Hi'" n. Becomes Hi". CrVI o. Becomes Cr'". FeVI p. Becomes Fe'". HYPOBROMOUS ACID. HBrO. Oxidation valence H'Br'O-" Structural valence H-O-Br ACID. HBrO. 877. The anhydride, Br20, has not been isolated. A solution of the acid is obtained by action of Br on HgO, in presence of H2O. HgO + 38r3 -f HaO = Hgßr? -f 2HBrO Bromic Acid. 263 On warming, the HBrO is decomposed at 60° 0. But it may be distilled unchanged in a vacuum at 40° C. HBrO is also formed by action of Br on AgaO or on AgH03. The hypobromites of K, Ha, Ba, Ca, Sr may be formed by adding Br to an excess of the cold hydroxides. Aqueous HBrO is a yellow liquid, very instable, and a strong oxidizing and bleaching- agent. The hypobromites are more instable than the corresponding hypo- chlorites. BROMIC ACID. HBrOa. Oxidation valence H'BrvO~"3 O 11 Structural valence H-0-Br=o 878. The anhydride, Br206, has not been isolated, and the acid, HBrOs, is known only in solution. The bromates of K, Na, Ba, Sr, and Ca are made by action of Br on the hot hydroxides (a)\ also by treating Br with Cl in presence of the hydroxides (£), and separating it from the bromide either by alcohol or recrystallizing. The remaining bromates may be made by adding bromic acid to the hydroxides or carbonates. Those of the first group and a few others are best made by precipitation. Free bromic acid may be made by action of Br on AgßrOa (c), or by adding exactly enough H2S04 to Ba(BrC3)2. 879. Properties.—lt may be concentrated in a vacuum until it contains 50.59 per cent. HBrC3. Under ordinary atmospheric pressure decomposi- tion begins when the liquid contains 4.27 per cent. HBrO,; Br, O, and H2O being formed (cl). The bromates are all decomposed by heat, some forming a bromide and O—e.g., the bromates of the alkalies, Hg. and Ag (e). Some forming an oxide and Br—e.g., bromates of Mg, Al, and Zn (/). Others a mixed oxide and bromide—e.g., bromates of Pb, Cu, etc. When fused with re- ducing agents bromates explode like chlorates. a. 6KOH + 38r2 = KBrO, + SKBr + 3H2Q b. Br2 + 12KOH + 5C12 = 2KBr03 + 10KC1 + 6H2Q c. + 3Br2 + 3H20 = SAgßr + 6HBrOs d. 4HBrOa = 2H .O 4- 3Br2 + 502 e. 3K8r03 = 2KBr + 30, /. 2Zn(8r03)2 (at 200° C.) = 2ZnO + 2Br2 + 5Q2 880. All the bromates are soluble in water; those of the first-group bases but spar- ingly soluble. Silver nitrate precipitates, in solutions not very dilute, silver hromate, Agßr03, white, very sparingly soluble in water, soluble in ammonium hydroxide, not eas- ily soluble by nitric acid, its color and solubility in ammonium hydroxide differing a little from the bromide (872). It is decomposed by hydrochloric acid with evolution of bro- mine—a distinction from bromides and from other argentic precipitates. 264 lodine. 881. Sulphuric, hydrochloric, and nitric acids liberate bromic acid from metallic bromates. With very dilute sulphuric acid, in cold dilute solution of pure bromate, very little bromine is set free—the HBrCs mostly remaining for some time intact, and the solution colorless, so that carbon disulphide will not extract much color. The gra- dual decomposition of the HBrOs is first a resolution into HBr and O, and as fast as HBr is formed it acts with HBr03, so as to liberate the bromine of both acids. Now, if the solution contained bromide as well as bromate, an abundance of free bromine is ob- tained immediately upon the addition of dilute sulphuric acid in the cold. 882. Hence, if dilute sulphuric acid in the dilute cold solution does not color the car- bon disulphide, and if the addition of solution of pure potassium bromide immediately develops the yellow color, while it is found that no other oxidizing agent is present, we have corroborative evidence of the presence of a bromate. And, if we treat a solution known to contain bromide with dilute sulphuric acid and carbon disulphide, and obtain no color, we have conclusive evidence of the absence of bromates. A mixture of bromate and iodate, treated with hydrochloric acid, furnishes bromine without iodine, coloring carbon disulphide yellow. The ignited residue of bromates (879), in ail eases if the ignition be done with sodium carbonate, will give the tests for bromides. 883. Oxidation.—a. When bromic acid is reduced an excess of the reducing agent usually produces HBr or a bromide. But with a few reduc- ing agents only free bromine is formed, b. For action upon H2C„04 see 659 ; HONS, 702 ; H4Fe(CN)6, 691 ; H3PQ2, 756; H2S, 785 ; H2S03, 813 ; HCI, 851 ; HBr, 876. HI c. Forms Br° and I°. Hg' d. Becomes Hg'' and Br~' (a bromide). Sn" e. Becomes SnIV and Br~'. Sb'" /. Becomes Sbv and Br~'. Pe" g. Becomes Pe'" and Br'. As'" h. Becomes Asv and Br~' (HBr). lODINE. 1 = 126.5-57. Oxidation valence of free iodine (I )° Structural valence of free iodine I-I 884. The vapor density (H = 1) up to 700° C, is 126.5, showing that the molecule is 12.I2. But at 1400° C. it is about two-thirds as much, and on in- creasing the temperature a point is reached at which the density is about one-half, 63.3, and a farther increase of temperature no longer decreases the density—a very significant fact, showing that the molecule dissociates into atoms, and that the original molecule contains two atoms, and not more than two. The sp, gr. of solid iodine (water =1) at 17° C. is 4.948 (Gay-Lussac). It melts at 114.15° C., and boils at 184.35° C. (Ramsay .and Toung, 1886, Jour. Chem. Soc., 49, 453). lodine. Occurrence.—It is found in sea-water, in some salt springs, and in some vegetable growths. The quantity of iodine in sea-water is very small, but certain sea-weeds have the power of storing it up. 885. Preparation.—(l) Sea-weed is carbonized by heat in a retort. The residue, called help, is lixiviated with water, which dissolves the sodium iodide. It now contains carbonates, sulphates, sulphides, etc., from which it is partially freed by crystallizing. It is then mixed with one-eighth its bulk of H2S04, and allowed to stand for 24 hours. The C02 and H2S pass off, and much of the newly formed Na2SQ4 crystallizes out. Finally Mn02 is added, heat applied, and iodine sublimes (a). After all the iodine has sublimed the bromine begins to be set free, and is condensed. The amount of bromine thus secured is about one-tenth that of the iodine. (2) Another method sometimes used is to precipitate the iodine by CuS04 and FeSQ4 (£), and then distill the Cu2I2 with MnO„ and H2S04. (3) When the solution contains only traces of iodine, it is freed by nitro-hydrochloric acid, and filtered through lamp-black, which absorbs it. KOH solution is then added, which dissolves it (c). The residue after evaporation is then sublimed with MnOs and H2S04. a. 2NaI + MnO, + 2H>S04 = Na2S04 + MnSQ4 + 2H20 + I, h. 2NaI + 2CuS04 + 2FeS04 = Cu2I2 + Na2504 + Fe2(S04)3 c. 312 + CKOH = KI03 + SKI + 3H20 886. IODINE is solid; in soft scales or hexagonal prisms, with a dark iron-gray color and graphitoidal lustre. It is precipitated as a brownish- black powder. It vaporizes very slightly at ordinary temperatures—with a characteristic odor, resembling chlorine, but more offensive. It melts at 114.15° C., and boils at 184.35° 0.; the vapor having an intense, bright violet color. 887. It is slightly soluble in water, dissolving in 7,000 parts ; freely soluble in alcohol, ether, chloroform, carbon disulphide, petroleum naphtha, glycerine, and in solutions of iodides (including HI). All solutions of un- combined iodine have red-brown, brownish-yellow, or violet tints. The carhon disulphide solution is violet (marked distinction from bromine), the other solutions brownish-yellow (but little darker than those of bromine). Solutions by chemical combination are referred to in 889. 888. Starch-paste is colored blue by a little iodine, violet by a further addition of iodine ; and by still greater excess a blue-green (or, in presence of bromine, a brown) color is produced. This test is exceedingly delicate for iodine.* The iodized starch is decolored by heating in solution to 70° or 80° C. * The union of iodine and starch is probably an example of molecular adhesion rather than of union within molecules. When dry starch is saturated with ether solution of odine, and exposed for some time to the heat, of the water-bath, about 4 per cent, of iodine is retained. This corresponds nearly with the formula 0Os'2„I, Prepared under other conditions, it holds 7 to 8 per cent, of iodine (CBHjOO6),01. 266 lodine. (158° to 176° F.), but regains its color on cooling. Its color is destroyed by strong chlorine, and by alkalies. No compound of iodine colors starch. 889. Though expelled from combination with bases by chlorine, bromine, nascent oxygen, and other strong electro-negatives, iodine acts in many relations as an oxidizing agent, readily entering into combination, as iodides, when acted on by reducing agents. On the other hand, in relation to a limited number of active electro-negatives, it may act as a reducing agent, becoming the subject of oxidation, in the formation of iodates. lodine chlorides also are formed, IC15, ICI3, and ICI, of yellow to brown colors. lodine slowly bleaches litmus and other vegetable colors, and stains the skin yellow- brown. Colorless solutions are formed by all the alkali hydroxides with iodine ; the fixed alkali hydroxides forming iodides and iodates (a). With ammonia in water solution, it dissolves more slowly, becoming colorless; the solution contains the most of the iodine as ammonium iodide, and liable to deposit a dark-brown powder, termed “ iodide of nitro- gen.” very easily and violently explosive when dry. This substance is a variable substi- tution of one, two, or three atoms of I for Hin NH3 (45 c). Among reducing agents, solutions of thiosulphates quickly dissolve and decolor iodine, forming iodides and a more highly oxidized acid of sulphur, tetrathionic acid (b). Solutions of sulphites and of sulphurous acid convert iodine into colorless hydriodic acid (c).—Arsenious acid in alkaline mixture is oxidized to an arsenate, and a colorless iodide is formed {d). Hydrosul- phuric acid dissolves iodine as hydriodic acid, the solution of which is so prepared (e). The altali hydroxides, and reducing agents, decolor iodized starch, by taking its free iodine into combination. a. 3I2 + 6KOH = SKI + KI03 + 3H20 (corresponding with 833 b). b. I 2 + 2Na2S203 = 2NaI + Na2S406 (compare c). c. I 2 + H2O + Na2S03 = 2HI + Na2S04 d. H3AsOs + SNaHCOs +I2 = Na3As04 + 2NaI + SCOa + 4H20 e. 21-2 + 2H2S = 4HI + S2 The chief Acids of lodine are : Hydriodic acid, Hl"'. lodic acid, HP 03. Periodic acid, HIVII04. 890. Estimation.—The estimation of free iodine is important, for upon it depends the determination of a large class of substances which liberate iodine from KI (e.g., chlorine, bromine, etc.), or which, when boiled with HCI, yield Cl {e.g., the higher oxides and metallic acids). The usual method is to add a standard solution of Na„S,03, in presence of a little starch-paste, to show when the action is complete. 2Na2S203 + I* = 2NaI + Na2S406 891. Oxidation.—When iodine oxidizes it becomes I-' ; that is, either hydriodic acid or an iodide. When it reduces it becomes Iv or Ivn ; that is, iodic acid or an iodate or a periodate, b. For its action on HNOs see 733 ; H4Fe(CN)6, 601; H3POa, 75(5; H3S, 785 ; H3SO;„ 813; Cl, 834 ; HCI03, 863 ; Br, 868. Hydeiodic Acid. 267 Hg' c. Becomes Hg" in presence of alkalies and acids. Sn" d. Becomes SnIV “ “ “ Sb'" e. Becomes Sbv in presence of alkalies only. As'" /. Becomes Asv “ “ “ MnIV ng. Becomes MnIV “ “ “ Co" h. Becomes Co'" “ “ “ Fe" i. Becomes Fe'" “ “ “ Cr'" j. Becomes CrVI “ “ “ Hi" k. Is not changed by iodine. HYDRIODIC ACID. HI. Oxidation valence HI-' Structural valence H-I 892. Vapor density (H =1) is 63.7. At 0° C. a pressure of four at- mospheres liquefies it. Under ordinary atmospheric pressure it solidifies at -51° 0. (Faraday). Preparation.—(l) By action of P on iodine and separation by distilla- tion. (3) By adding H3P04 to KI and distilling. (3) By action of H2S on iodine and distilling. 1. P4 + 10I2 + 16HnO = 4H3P04 + 20HI 2. SKI + H3PO4 = KsP04 + SHI 3. 2H2S + 2I2 = 4HI +S2 893. Absolute bydriodic acid is gaseous at ordinary temperatures, but freely soluble in water; being easily obtained in solution containing 50 to 60 per cent, of acid, and having a boiling point above that of water, but giving off some vapor at common temperatures. Both the gas and the solution are colorless, and redden litmus. Hydriodic acid decomposes gradually in the air with separation of iodine—more rapidly at higher temperatures; so that the evolved gas is always strongly colored with iodine, and the exposed solution commences at once to turn brownish- yellow with tiie free iodine dissolved by the acid. The liberated gas has a slight chlorine-like odor, and a stronger offensive odor (due to both the iodine and hydriodic acid). Upon brief exposure, both the gas and the solution give abundantly the reactions of free iodine (with starch, carbon disulphide, etc.) 894. Like hydrochloric and hydrobromic acids, hydriodic acid is pro- duced by transposition from the metallic iodides, by the action of dilute sulphuric acid; but an attempt to separate the HI by distillation would result in its decomposition, as shown in equation/. Also, by large excess of hydrochloric and hydrobromic acids. The iodides of silver, lead, mer- cury, and tin are transposed with difficulty by sulphuric acid, more readily by hydrochloric acid. Hydriodic Acid. 895. The iodides (including hydrogen iodide) are decomposed by oxidizing agents more readily than the bromides. Ozone promptly decomposes all iodides, not excepting those of the alkali metals ; while atmospheric oxygen decomposes hydriodic acid and iron and calcium iodides but slowly, and alkali iodides not at all. lodine is liberated from iodid at once by chlo- rine, bromine (a), iodic acid (b), and bromic acid. lodine is first set free and then oxi- dized to iodic acid, by acidulated chlorate, by hypochlorites' (with occurrence of iodine chlorides and final formation of periodates), and by concentrated nitric acid with heat (c); dilute nitric acid slowly separating iodine {d), and scarcely decomposing lead, silver, and mercury iodides. Acidulated potassium nitrite acts more promptly than nitric acid. Manganese dioxide with sulphuric acid is employed in the manufacture of iodine (e). Permanganate solution, added in excess, produces iodates, iodine being first separated and at last all oxidized ; in neutral or alkaline dilute solutions (1 part salt to '240 parts water), a distinction from bromides, which do not decolor the permanganate. Chromates, acidulated, cause immediate separation of iodine. Concentrated sulphuric acid (/) and ferric chloride (g) are reduced by iodides. Further, see 907. a. 2HI + Br- = 2HBr + I 2 b. SKI + KIO3 + 3H2S04 = 3K.504 + 3H20 + 312 c. KI + 2HN03 = KI03 + 2NO + H2O d. 6KI + BHNO3 = GKNO3 + 2NO + 4H20 + 8I2 e. 2KI + 2H2504 + MnO,. = K2S04 + MnSQ4 + 2H20 + Ia /. 2HI + H2S04 = 2H..0 + S02 + I 2 g. 2KI + Fe2Clo = 2FeCl2 + 2KCI + I 2 896. The metallic iodides are all soluble in water; except those of Ag, Pb, and Hg, except pulladous iodide, cuprous iodide, bismuth iodide de- composed by water, and stannous iodide sparingly soluble in water. Lead iodide is sparingly soluble, and mercuric iodide very sparingly soluble, in water. The double iodides of lead, silver, and mercury with alkali metals—as KI.AgI and (KI) 2HgI,—are soluble in water; i.e , the iodides of first- group metals are soluble in solutions of alkali iodides, by combination ; mer- curous iodide in part only, as explained in 438. Alcohol dissolves many of the iodides soluble in water—including the alkali iodides, and those of barium and calcium—and dissolves mercuric, but not mercurous or argentic, iodide. Silver iodide is but very sparingly soluble in concentrated solution of ammonium hydroxide, and insoluble in hot solution of ammonium acid car- bonate (distinctions from the chloride). It dissolves in solution of potas- sium cyanide. The iodides of silver and of lead are soluble by decomposition in solu- tion of alkali thiosulphates (a); lead iodide in fixed alkalies (387). The iodides of silver and mercury are not decomposed, the iodide of lead slowly decomposed, by dilute nitric acid. a. Agl -(- Na3S303 = Nal -J- NaAgS203 Hydriodic Acid. 269 In analysis, iodides are most easily identified by the color of the carbon disulphide solution of liberated iodine (902). The silver precipitate of iodide is separable from chloride by solution of the latter in ammonium hy- droxide (921). 897. Silver nitrate solution in excess precipitates, from solutions of iodides, silver iodide, Agl, yellow-white, blackening in the light (without notable separation of iodine). For the solubilities of the precipitate, see 896, and compare 414. For its separation from chloride and bromide, see further 921 Solution of mercuric chloride precipitates the bright, yellowish-red to red, mercuric iodide, Hgl2. The precipitate redissolves on stirring, after slight additions of the mercuric salt, until equivalent proportions are reached, when its color deepens. For the solubilities of the precipitate see 446.—Solution of mercurous nitrate precipitates mercurous iodide, Hg2I2, yellow to green (see 438). 898. Solution of plumbic nitrate or acetate precipitates, from solutions of iodides not very dilute, lead iodide, Pbl2, bright-yellow—soluble, as stated in full in 392. 899. Palladous chloride, PdCl2, precipitates, from solutions of iodides, palladous iodide, Pdl2, black, insoluble in Avater, alcohol, or dilute acids, and visible in 500,000 parts of solution. The reagent does not precipitate bromine at all in moderately dilute solutions, slightly acidulated with HCL Palladous iodide is slightly soluble in excess of the alkali iodides, and is soluble in ammonium hydroxide (595). 900. Copper sulphate, with sulphurous acid or other reducing agent, precipitates from solutions of iodides the cuprous 'iodide, Cu2I2, which is ivhite, if there is sufficient reducing agent to prevent the precipitation of iodine, broAvn. The precipitate is not altogether insoluble in water ; there- fore the filtrate responds to the delicate tests for iodine (equation in 345 b). Bromine is not precipitated Avith copper. 901. Concentrated sulphuric acid decomposes iodides, solid or in con- centrated solution, with the reaction stated at 895/. The evolved gas has the violet color of iodine, and the offensive odor of mingled iodine and hy- driodic and sulphurous acids. When cooled and somewhat diluted, the liquid gives the iodine color with starch (888); or, on agitating gently Avith carbon disulphide, and permitting the latter to subside, the beautiful violet tint of iodine in this solvent. 902. Chlorine-water separates iodine more satisfactorily, in this test Avith carbon disulphide, especially from dilute solutions. The chlorine- water should be dilute and added (after the starch-paste or carbon disul- phide) drop by drop ; as an excess will destroy all characteristics of free iodine by formation of iodine chlorides and iodic acid (920). Nitrous acid-—as from zinc and nitric acid or from acid illation of ni- 270 Hydriodic Acid. trite's—is a good agent to displace iodine. It should be very sparingly used. (709). Bromine water is also employed for the same purpose. Bromides do not interfere with the easy recognition of free iodine ; un- less an excess of chlorine is added no bromine will be liberated, and, if liberated, it does not modify the color of iodine, in starch or in carbon di- sulphide, unless the bromine is in much greater quantity, and even then the color represents iodine. 903. Solution of ferric chloride, added in the proportion of 6or 8 drops to 3 or 4 cub. cent, (a fluid drachm or a little less) of the solution tested, together with carbon disulphide, slowly develops the violet tint in the sub- siding liquid, if iodine is present (895 g)—a distinction from bromine. For Separation of iodides from chlorides and bromides, 920 ; from iodates, 913. 904. The iodides of the alkali metals and oC the first-group metals fuse without de- composition; those of mercury sublime undecomposed; but other non-alkali iodides are mostly decomposed by ignition. Treated in the cupric bead of microcosmic salt, as directed for chlorine in 848, iodides give an emerald green glass. 905. Estimation.—(l) G-ravimetrically. It is precipitated by AgHQ3, and after gentle ignition weighed as Agl. (2) Yolumetrically, by stand- ard solution of AgHO,, using a little K2Cr2CX to show the completion of the reaction. (3) It is oxidized to free iodine by chlorine or other convenient oxidizing agent, and then determined by standard solution of 3NTa2S203. 906. Preparation of lodides.—(1) By direct union of the elements, but in some cases heat must be employed, and the metal must be in a fine state of division. (2) By action of HI on the hydroxides, oxides, and carbonates of the metals. (3) lodides of the first group are best made by precipita- tion. (4) Many metals, especially when finely divided, dissolve in HI with evolution of hydrogen. 907. Oxidation.—a. For the action of HI on H6Fes(CH)ia see 694 ; HN02, 713 ; HNOa, 733 ; H2S04, 829 ; Cl, 834 ; HCI03, 863 ; Br, 868 ; HBr03, 883. HI03 h. Free iodine is liberated from botli acids. Pb" f n c. Becomes Pb" and 1° (free iodine). Asv d. Becomes As'" “ I°. Sbv e. Becomes Sb'" “ I°. Cu" /. Becomes Cu' “ I°. Co"' g. Becomes Co' “ I°. Hi"' h. Becomes Hi" “ I°. Fo'" i. Becomes Fe" “ I°. i k; CrVI j. Becomes Cr'" “ I°. lodic Acid. 271 Mn" + w Becomes Mn", and if HI is in excess 1° is formed. If dilute hydriodic acid with K„Mn2Oa is used manganese peroxide is first formed, and if the permanganate is in great excess potas- sium iodate is formed. lODIC ACID. HIOa. Oxidation valence H'PO“"3 O 11 Structural valence 11-0-I=o 908. Preparation.—(l) By boiling I witli HN03, sp gr. 1.42, or, better, the fuming, sp. gr. 1.48. (2) By treating I with Cl or HCIO in presence of H2O. (3) By transposing iodates with stronger acids, as Ba(I03)2 with H2S04, using the latter slightly in excess, and after filtering from the insoluble BaS04 the HIQ3 is separated from the remaining H2S04 by crystallization. (4) Formed by action of auric oxide on I. (5) By action of AgN03 on I. 1. 3I2 + lOHNO3 = OHIO 3 + lONO + 2H20 2. I 2 + 5C12 + 6HaO = 2HIOs + 10HC1 4. 3I2 + 5Au203 + 3H20 = OHIO3 + IOAu 5. SAgN03 + 3I2 + 3H20 = SAgI + SHNOs + HIO3 909. Absolute iodic acid. HIO3 (or H>I206), is a white, crystallizable, odorless solid, permanent in the air ; at. 170° C. (388° F.; resolved into water and iodic anhydride (H2O and I 205). lodic anhydride is a crystallizable solid, at 300° C. resolved into iodine and oxygen. Bromic anhydride is not known. lodic acid is freely soluble in water and in alcohol ; the solutions reddening litmus, and afterwards bleaching it.* 910. lodic acid is formed by prolonged action of nitric acid and other oxidizing agents upon iodine. Its salts, the iodates, are formed together with iodides in dissolving iodine in aqueous alkalies (889 a), as well as by oxidation of iodides (895). lodic acid is easily obtained by transposing metallic iodates with sulphuric acid (a); its radical not easily breaking up when separated from metals, as chloric and bromic acids do. 911. The lodates—including hydrogen iodate—are decomposed by reducing agents, with the formation of iodides (of metals or of hydrogen) and with other results. Sulphurous acid is oxidized by iodic acid, first with separation of iodine («); then, by excess of the sulphurous acid, with formation of hydriodic acid (5). Hence, sulphur- ous acid, added short of saturation, with starch, forms a delicate test for iodates, and a distinction from iodides ; but excess of the reagent destroys the color. Thiosulphates produce iodine, or hydriodic acid. Hydrosulphuric acid also reduces iodates, pre- cipitating at first iodine and sulphur (c). With excess of the reducing agent, the final products are hydriodic acid and sulphur (d) ; with excess of the iodate, iodine and sulphate (e). Hydriodic acid instantly separates, from iodic acid, all the iodine of both acids (/); hence, an intermixture of a metallic iodate with an iodide is revealed at once by adding a dilute or weak acid that will not itself liberate iodine, but will produce both the acids of iodine, so that they can decompose each other. In solution of potassium iodide, for example, a slight addition of tartaric acid shows the presence of iodate by the immediate, not progressive, appearance of the iodine color, the test being more deli- cate by use of carbon disulphide. In solutions not of iodides, an iodide may be added, a. 2K103 + H,S04 = K2S04 + 2HIOs. lodic Acid. with tartaric or acetic acid, in search for iodates. But it must be remembered that pure iodides, so treated, form hydriodic acid, which, by atmospheric oxidation, progressively liberates iodine, and will soon give a deep color to starch or carbon disulphide. Hydro- chloric acid forms with iodates mostly iodine chlorides (g), iodine not being liberated (distinction from broraates, 883). Morphia reduces iodic acid, with separation of iodine as a final product. Further, see 916. a. 2K103 + 5H2SQ3 = I* + 2KHSO., + 3H2504 + H2Q Or: 2HIOs + 5H2503 = I 2 + 5H2SQ4 + H2O b. The iodine is taken up by excess of H2S03, as in 889 c. c. 4H103 + 10H..S = 2I2 + 12H..0 + 5S2 d. 2HIOs + 6H2S = 2HI + 6H20 + 3S2 e. BKIO3 + SH2S = 4I2 + 3K2S04 + 2KHSO, + 4H20 /. HI03 + SHI = 31* + 3H20 g. KIO3 + OHCI = (Id* + 2CI) + KCI + 3H2C lodates in dry mixture with combustible bodies are reduced, on heating or con- cussion, with detonation, but much less violently than chlorates or nitrates. Heated alone, iodates are either reduced to iodides with liberation of oxygen (iodates of potas- sium, sodium); or to oxides with liberation of iodine and oxygen (iodate of barium). Compare Bromates, 879. 912. The iodates are either insoluble or sparingly soluble in water, except those of the alkali bases, a marked difference from bromates and chlorates. Barium, silver, and lead iodates are insoluble in water. The alkali metals form acid iodates. In alcohol most of the iodates are insoluble ; barium iodate, insoluble ; calcium and potassium iodates, scarcely at all soluble (distinctions from iodides). Silver iodate is readily soluble in ammonium hydroxide (distinction from iodide); it is slightly soluble in dilute nitric acid (more so than the iodide). lodates are identified by separation of free iodine, known by its color in carbon disul- phide solution or in mixture with starch (911); and by precipitation of barium salt (912). 913. Solution of silver nitrate precipitates, from even very dilute solutions of iodates and from solutions of iodic acid if not very dilute, silver iodate, AgIO„ white, crystalline, soluble in ammonium hydroxide, soluble in an excess of hot HN03. In the ammonia solution, hydrosulphuric acid precipitates silver iodide. Barium chloride precipitates barium iodate, Ba(I03)3, nearly insoluble in cold and little soluble in hot water, insoluble in alcohol, scarcely soluble in dilute nitric acid, readily soluble in dilute hydrochloric acid. Hence, dilute solutions of free iodic acid should either be neutralized or tested with barium nitrate. This precipitate, by addition of alcohol, is a complete separation from iodides, and, when well washed, decomposed with a very little sulphurous acid (911 a), and found to color carbon disulphide violet, its evidence for iodic acid is conclusive. Barium iodate is transposed with ammonium car- bonate, on digestion in solution and with ammonium hydroxide (separation from perio- date). Salts of lead give a white precipitate of lead iodate, Pb(103)2. Ferric chloride gives, in solutions not dilute, a yellowish-white precipitate of ferric iodate, Fe2(103)6, spar- ingly soluble in water, and freely soluble in excess of the reagent. Boiling decomposes it. Alcohol precipitates potassium iodate from water solution, an approximate separa- tion from iodide. 914. lodates of the alkalies and alkaline earths are easily made by the action of iodine on the hydroxides, and separation by alcohol or by crystallization from the iodides which are formed in the reaction. All iodates may be made by action of the acid on the Periodic Acid. hydroxides or carbonates. A few are best made by precipitation. The precise composi- tion of many of the iodates has not been determined. 915. Estimation.—(1) By precipitation with AgNOa, and, after drying at 100° C., weighing as Agl03. (3) By reducing to HI or an iodide and then proceeding as directed in 90). (8) Yolumetrically, by treating with HI and estimating the free iodine liberated. 916. Oxidation.—a. For action on H2C204 see 659; HfiFe2(CN)12, 694; HCNSr 703; HN02, 713; H3P02, 756; H2S, 785; H2S03, 813; C 1,834; HBr, 876; HI, 907. Sri'' 6. Becomes Snlv and I-' Sb c. Becomes Sbv and I°. As «. Becomes Asv and I®. Fe" v Becomes Pe'" and I°. PERIODIC ACID. niO, or HSI06. Oxidation valence HTviiO~"4, or H'5IviiO~"s H O H-O O O-H ii M/ Structural valence H-0-I=o, or H-O-I^O-H n n O O 917. Preparation.—(l) By action of HCI04 on I (a). (3) By action of H2S04 on Pb(103)2 (b). (3) Periodate of sodium is made by action of Cl upon I, Nal, or NalOs in. solution of NaOH (c). The potassium sait is prepared in a similar manner. a. I 2 -f 2HC104 = 2HI04 + Cl2 b. Pb(104)2 + H2S04 = PbS04 + 3HI04 c. I 2 + 16KOH + 7C12 = 2KIO, + 14KC1 + BH2O Neither the anhydride nor the acid HIC4 has been isolated. Only HI04, with two molecules of water, has been crystallized. One view represents the normal acid as HIO,, which is supported by the actual preparation of the corresponding salts, such as KIOi, Agl04, Pb(IC4)2, etc. Another view represents the normal acid as H6I06 (H104.3H20 = HslOs). This view is supported by the actual production of Ag6I06, NasIO„, Ba5(I06)2, etc. The aqueous solution may be boiled, but at 140° C. it begins to decom- pose, forming first HIO3, oxygen, and water; then 12C6,I2C6, and finally free iodine and oxygen. Ignition reduces alkali periodates to iodides, evolving oxygen. Ignition also reduces other periodates in the same manner that it does iodates. They are all reduced by HCI, evolving Cl; by H2SC3, giving first free iodine and then HI, and forming H2S04; and HI liberates free iodine, but only one-fourth of the total quantity in the salt (Rammelsbukg). According to Lautsch, its behavior with Hg2(NOs)2 is characteristic. The pentasodic periodate, NaslOe, gives a light yellow precipitate consisting of deca- mercurous periodate, Hg10I2O12. 5Hg3(N03)3 + SNa&IOe = Hg10I3Ol 3 + IONaNOs Reactions of the Acids of Chlorine, Etc. Chlorides. Bromides. Iodides. Chlorates. lodates. Hypochlorites. AgNOs, in excess — Ag-Cl, white (842). Ag-Br, white (872). Agl, yellow-white (897). No pre. (859). AgI03, white (912). No pre. (853). Ag-N03, with, excess of the solution Pre.? (409). Pre. (414). Sol. (414). Pre. tested. NH4OH, to the Ag pre Dissolved. In part dissolved. No pre. Not dissolved. Dissolved. TTcrrm^_ Hg-I2, yellow-red. No pre. No pre. Hg-Cl,, with excess of the solution No pre. No pre. (446). Hg-„I2 (887). tested. Hg-2(N03)2,in excess Hg-2C12, white. Hg-2Br2, white. No pre. Pre. No pre. BaCl2 Ba(I03)2 (913). CuS04. with H2SOs Cu2I2 (900). HI (894) (911 a) HIOs (910). H2S04, dilute— HC1 (soluble gas). HBr HC103, C102, Cl Cl (bleaches). HC1 HBr HI C102 and Cl Cl Cl Br (873). Br I (902). I Chromate, with HoSC4 Cr02Cl2 (845). KjMnjOg, dilute, neutral I (895). NaoSOi, with H2S04 I and HI (911). I, HI, S (911). HC1 s, h2so4. H2S S, H2S04 (785). 918. Comparison of Certain Reactions of the Acids of Chlorine, Bromine, and lodine, Taken in Water Solution, as Potassium Salts, or other Soluble Compounds. Separation op the Acids of Chlorine, Etc. 275 919. THE SEPARATION of the acids of chlorine, bromine, and iodine is-eflfected by oxidations, reductions, color solutions, precipitations, separative solutions, and vapo- rizations. In many cases of separation, the acids to be sepai/ated will act upon each other. 920. The Recognition of chlorides, bromides, and iodides—by evolving their chlo- rine, bromine, and iodine, in presence of each other—can be accomplished as follows— for the iodine the test being very easy ; for chlorine, indirect but unmistakable ; for bro- mine, dependent upon much care and discretion.* The lodine is liberated with dilute chlorine-water, added drop by drop, and is readily detected by starch, or carbon disulphide, according to 903. (As to interference of thiocyanates, see 701.) The Chlorine is vaporized (from another portion) as chloro- chromic anhydride, and the latter identified by its color and its various products, as described in 845. Before the Bromine is identified the iodine is to be either removed as free iodine, or oxidized to iodate (873). The oxidation to iodic acid is effected as fol- lows: Treat with good chlorine-water till free iodine no longer shows its color ; add a drop or two more of the chlorine-water, and dilute with water, keeping cool ; then add the carbon disulphide, agitate, and leave the solvent to settle, for the yellow color of bro- mine. The removal of free iodine may be done as follows: Add chlorine-water, drop by drop, as long as the iodine tint seems to deepen by the addition ; add the carbon disul- phide, agitate, leave to subside, and remove the lower layer, either by taking it out with a pipette, or by filtration through a wet filter. Repeat, if need be, till iodine color is no longer obtained ; then continue, with dilute chlorine-water, in test for bromine. If iodide in large proportion is to be removed, it is well, first, to precipitate it out, as far as possible, by copper sulphate and a reducing agent, as directed in 933. The filtrate is then to be treated by either method above given. 921. The Separation by ammonium hydroxide, as a solvent of the silver precipi- tates— AgCl, Agßr. Agl —when conducted with dilute ammonium hydroxide, may be made nearly complete between the chloride and the iodide, but it is very imperfect be- tween the bromide and either of the others. The hot and strong solution of ammonium acid carbonate separates the chloride from the bromide (compare 843, 873, 896). 922. The direct removal of iodides hy precipitation, leaving bromides and chlorides in solution, can be effected (approximately) by copper sulphate with sulphurous acid (900), or quite completely by palladous chloride (899). With the copper sulphate, the reduction ought to be thorough ; and this result is better secured by sulphurous acid than by ferrous sulphate, and without loading the solution with another metallic salt. The action of palladium chloride is subject to no objection, except the scarcity and ex- pensiveness of the reagent. 923. Chloric acid is separated from hydrochloric and all other acids of chlorine, bro- mine, and iodine (except from hypochlorous acid, and from traces of bromic acid), by remaining in solution during the precipitation by silver nitrate (859). 924. Chloric acid is separated from nitric acid—after finding that silver nitrate gives no precipitate in another portion of the solution, acidulated—by evaporating and igniting the residue, then dissolving, and testing one portion of the solution by silver nitrate for the chloride formed from chlorate during ignition (858). The other portion of the solution is tested for nitric or nitrous acid. * In consequence of the relative commercial values of bromine and iodine, and the medicinal relations of bromides and iodides, it is of great importance to search commercial iodides for intentional and consider- able mixtures of bromides—an impurity likely to escape cursory chemical examination. There are, how- ever, very slight and usually unobjectionable proportions of bromides generally to be found in the iodides of commerce, and occurring from the difficulty of exact separation in the manufacture of iodine from kelp. 276 Equations. 925. If we have to separate chloric acid both from nitric and hydrochloric acids, a solution of silver sulphate must be used instead of the nitrate, to precipitate out all the hydrochloric acid. The filtrate from this is evaporated, ignited, dissolved and tested, as in 924, for chloride, indicating chlorate in the original solution, and another portion is tested for nitric acid. Also, chlorates are distinguished (not separated) from nitrates, by oxidation of ferrous sulphate in solution with acetic acid on heating, and the conse- quent formation of the red solution of ferric acetate (218, 635). The solution tested must contain no free acids, and no nitrites or other oxidizing agents beside the two in question, but may contain chlorides ; and, of course, the ferrous sulphate must be pure enough not to color when heated alone with the acetic acid. Mix the ferrous sulphate solution with the acetic acid, boil, then add the solution to be tested, and heat nearly to boiling, for some minutes. If no red color appears, chlorates are absent, and nitrates may be present. 926. Hypochlorites are separated with chlorates from chlorides (bromides), etc., by silver nitrate; and distinguished from chlorates (in the filtrate from AgCl, etc.) by bleaching litmus, and by their much more rapid decomposition and consequent precipita- tion of any silver in solution. They are also more active than chlorates, as oxidizing agents. 927. The identification of iodic acid is simple and certain, by use of reducing agents (.911), or precipitants (913). The identification of bromic acid, in presence of other acids, is indicated in 880 to 882. 928. G-. Yortman’s method of detecting chlorine in presence of bromine and iodine is as follows : The solution containing the halogens combined with the alkali or alkali earth metals is heated with acetic acid and peroxide of lead until the supernatant liquid is colorless and has no longer the slightest odor of iodine or bromine ; in this way the whole of the bromine and part of the iodine are driven off, the remainder of the latter remaining as iodate of lead along with the excess of lead peroxide. This is filtered off, the precipitate washed with boiling water, and the chlorine precipitated from the filtrate by addition of silver nitrate. M. Dechan’s method {Jour. Gnem. Soc., 1886, 49, 682) consists (1) in boiling the mix- ture with a solution of 40 grammes of K2Cr.j07, dissolved in 100 c.c. of water, which liberates and expels all of the iodine without disturbing the bromine and chlorine. (2) 8 c.c. of a dilute solution of sulphuric acid (consisting of equal volumes of H2S04 sp. gr. 1.84, and water) are added to 100 c.c. of the dichromate solution, and on boiling the bromine is distilled off without disturbing the chlorine ; after which the chlorine is de- tected in the usual manner. For A. Longi’s process for the analysis of a mixture of chlorides, bromides, iodides, chlorates, bromates, iodates, ferrocyanides, and ferricyanides, see Ghem. News, 47, 209. 5K2Cr207 + GKI = Cr203 + 8K2Cr04 + 8Ia EQUATIONS. 929. It is recommended that the student write all the equations representing the analysis, oxidation, and properties of all compounds, and balance them according to the rule given in 615. It will be seen that in most cases an oxidizing agent is made to act on two reducing agents, usually both in the same salt; or double reduction may occur—thus in No. 25 Pb02, K2C03, KOI, and H..0 are formed; in No. 66, K2MnO4, K2SO4, NO, and CO2; in No. 57, K2Cr04, K2Mn04, KOI, and C02; in No. 80, the Mn of both com- pounds is oxidized ;in No. 17, Hg, As, SnCI4, and H2O are formed ;in No. 21, Pb, Na2S, C02, and H2O ;in Nos. 15 and 61 there is no action. In balancing equations for Equations. 277 recitation, the student should in all oases give his authority for each equation. It is to be understood that the second substances in each equation are to be used in excess, or in as great quantities as may be necessary to fully oxidize or reduce the substance placed first. Unless otherwise expressed, H2S04 means the dilute acid. It is advised that the teacher extend this list to several hundred for class use. The frequent introduction of blank equations like Nos. 15 and 61 has proven instructive. 1. Mn(OH)2 + Pb02 + HN03 39. KC104 (ignition) 2. Pb3(As03)2 + KOH + Cl2 40. Fe20(Cr04)2 + SnCl2 + HC1 3. Sn + HN03 (sp. gr. 1.40) 41. FeBr2 4- KOH + Cl2 4. Pb3(As04)2 + A1 + KOH 42. Cu(N0312 + A1 + KOH 5. Mn(H2P02)2 + KC103 + K2COs 43. MnS203 + KN03 + K2C03 (igni- (fusion) tion) 6. K2S04 + Na2C03 + O'(fusion; 44. Na2S506 + NaHCOs + C (fusion) 7. MnS + Pb02 + H2S04 (hot dilute) 45. FeBr2 + HN03 8. SnS + KOH + Cl2 46. Cr2I6 + KOH + Br2 9. Fe(OH)2 + H.,SO, (sp. gr. 1.83 hot) 47. Mn.Oj + Pb304 + HNOs 10. Pb3(As04)2 + Zn + H2S04 48. Pb(H2P02)2 + KOH + Cl2 11. Cr2(S04)3 + Mn(NOs)2 + K2CG3 49. P4 + KOH (fusion) 50. Hg3(As04)2 + Na (amalgam) 12. MnSOs + KOH + I2 51. Cr2(S04)3 + KC103 + K2C03 (fu- 13. Zn403Cr04 + SnCl2 -f KOH sion) 14. NiC204 + KOH + Cl2 52. Bic0,(0r04)2 + SnCl2 + KOH 15. Fe + HN03(sp. gr. 1.43 cold; see 939) 53. MnSOs + Pb02 + H2S04 (hot di- 16. K0103 (ignition) lute) 17. Hg3(As04)2 + SnCl2 + HC1 (sp. gr. 54. MnS2O0 + KOH + Cl2 130) 55. K2Mn208 + H2S04 (sp. gr. 1.83 hot) 18. Na2S4Of, + b aOH + Cl2 56. AsH3 (ignition in the air) 19. A1 + KOH 57. Cr.Clo + Mn(C103)2 + K2C03 (fu- 20. MnS + KN03 + K2C03 (ignition) sion) 21. PbS04 + NaHC03 + C (fusion) 58. Cr203 + Or2(N03)6 + K2C03 (fu- 22. Cr03 + H2S04 (sp. gr. 1.83 hot) sion) 23. BiBr3 + KOH + Cl2 59. Bi20(Cr04)2 + SnCI2 + KOH 24. Mn304 + Pb304 + HN03 60. CoS03 + KOH + Cl2 25. PbC204 + KOH + Cl2 61. Cu + HOI (see 939) 26. Fe + H2S04 62. Fe203 (white heat) 27. EiAs04 + Na (amalgam) 63. Ag3As04 + SnCl2 -l- HC1 (sp. gr. 28. Or2Cl6 + KC103 + . K2C03 (fusion) 1.30) 29. Mn(OH)2 + K2JYrn2Og + H2S04 64. Na2S506 + NaOH + Cl2 30. MnC204 Pb304 + H2S04 (hot di- 65. Pb(NO?)2 + A1 + KOH lute) 66. MuS2Og + KN03 + K2C03 (igni- 31. FeS + KOH + Cl2 tion) 32. FeC204 + H2S04 (sp. gr 1.83 hot) 67. Na2S203 + NaHCOs + C (fusion) 33. K2Cr207 (fusion) 68. PbMn208 -f- Zn -t- H2S04 34. Mn304 + Mn(N03)2 + K2C03 (fu- 69. Fel2 + KOH + Br2 sion) 70. Mn304 + Pb3Ot -f- H2S04 35. Mn304 + Cr2(N03)6 + K2C03 (fu- 71. HgC204 + KOH + Cl2 sion) 72. Fe + H2S04 (sp. gr. 1.83 hot) 36. C + H2S04 (sp. gr. 1.83 hot) 73. SbCl3 + A1 + HOI 37. Fe3(As03)2 + KOH + Cl2 74. (NH4)20r207 (ignition) 38. As4 + HN03 75. Hg2C204 + K2Mn2Os + H2S04 Problems in Synthesis. 76. Mn(H2P02)2 + Pb02 + H2S04 84. Pb02 + H0SO4 (sp. gr. 1.83 hot) 77. FeSOs + KOH + Cl2 85. Sb203 (ignition) 78. Fe(H2P02)2 + H2S04 (sp. gr. 1.88 86. Cu5H2(As04)4 + HI hot) 87. Sbl3 + KOH + Cl2 79. KN03 + NH401 (fusion) 88. HgCr04 + A1 + KOH 80, Mn203 1 lV[n(N03)2 + K2003 (fu- 89. Mn304 t KC103 -f K2C03 (igni- sion) tion) 81. H3As04 + SnCl2 + HC1 (sp. gr. 90. K2Cr207 + NH..C1 (fusion) 120) 91. KOH + I2 82. S + H2S04 (sp. gr. 1.83 hot) 92. Fe3(As03)2 + KOH + I2 83. Cr2I6 + KOH + Cl2 PROBLEMS IN SYNTHESIS. 930. For the sake of more thorough drill in the principles of oxidation, a few prob- lems are here given; a part of them the student should practically work at his table, but they are chiefly designed for class exercises. Special care should be taken that a pure product be formed, and that the ingredients be taken from the sources indicated. Thus, in the 31st, the chlorine for the ammonium chloride must be obtained from silver chlo- ride, and the nitrogen of potassium nitrate must be converted into ammonia, and then united with the chlorine, and the product purified. The student is not to suppose that these problems represent operations that are finan- cially profitable, but merely chemical possibilities, and their solution will compel an ac- curate comprehension of a great variety of important principles. It is recommended that the teacher increase the number of these; an ordinary class may with profit discuss from three to five hundred. In each case the authority for each step in the process should be stated. 1. Make pure mercuric bromide, from mercurous chloride and aluminie bromide. 2. “ chromic chloride, ‘ potassium chromate “ hydrochloric acid. 3. “ arsenic acid, ‘ potassium arsenite. 4. “ potassium arsenate, ‘ potassium arsenite “ potassium hydroxide. 5. “ plumbic nitrate, ‘ plumbic chloride “ zinc nitrate. 6. “ mercurous nitrate, ‘ mercuric chloride “ bismuth nitrate. 7. “ mercurous oxide, ‘ mercuric oxide. 8. “ mercuric bromide, ‘ metallic mercury “ potassium bromide. 9. “ mercuric bromide, ‘ metallic mercury “ silver bromide. 10. “ lead nitrate. ‘ lead dioxide “ potassium nitrate. 11. “ mercurous phosphate, ‘ phosphoric acid “ mercuric chloride. 12. “ barium sulphate, * lead sulphide “ barium hydroxide. 13. “ bar’m hypophosphite, ‘ calc’m hypophosphite “ barium chloride. 14. “ leael chromate, ‘ chromic chloride “ lead sulphate. 15. “ chromic chloride, ‘ potas. acid chromate “ silver chloride. 16. “ barium chromate, ‘ chromic chloride “ barium sulphate. 17. “ mercuric chromate, ‘ mercuric sulphide “ chromium nitrate. 18. “ chromium sulphate, ‘ potas. acid chromate “ bismuth sulphite. 19. “ phosphoric acid, ‘ sodium phosphate. 20. “ phosphorus, ‘ calcium phosphate. 21. “ lead iodate, ‘ lead sulphate “ potassium iodide. 22. “ silver iodate. ‘ silver bromide “ potassium iodide. 23. “ ferric arsenate. ‘ ferrous sulphide “ arsenious acid. Problems in Synthesis. 279 24. Make pure mercuric bromide, from mercuric sulphide and lead bromide. 25. “ ammonium sulphate, i 6 potassium nitrate “ sulphur. 26. * ‘ ammonium chloride, i 6 lead nitrate “ silver chloride. 27. “ sodium chloride, i 6 sodium sulphate “ silver chloride. 28. “ phosphorus, sodium phosphate. 29. “ lead sulphide, i t triplumbic tetroxide ‘ ‘ bismuth thiocyanate. 60. “ ferrous sulphite, 66 ferrous chloride “ barium sulphate. 31. “ ammonium chloride, a potassium nitrate “ silver chloride. 32. * ‘ mercurous nitrate, i 6 mercuric chloride “ potassium nitrate. 33. “ potassium sulphate, i ( sodium sulphite “ potassium nitrate. 34. “ mercurous chloride. 66 mercurous sulphide “ ferric chlorate. 35. “ potassium iodide, ( 6 potassium chloride “ sodium iodate. 36. “ sodium iodate, sodium chloride “ potassium iodide. 37. “ sodium phosphate, 6 6 potassium phosphate “ sodium chloride. 38. “ potassium bromide, i I silver bromide “ potassium chloride. 39. “ potassium chloride, i i silver chloride “ potassium bromide. 40. “ strontium nitrate, il strontium sulphate “ mercurous nitrate. 41. “ mercurous sulphide, 66 mercuric bromide “ potassium sulphite. 42. “ potassium sulphate, 66 sodium sulphate * ‘ potassium hydroxide. 43. “ sodium sulphate, ( i potassium sulphate “ sodium hydroxide. 44. “ potassium chromate, 66 chromic chloride “ potassium hydroxide. 45. “ potassium iodide, I i sodium iodate “ potassium nitrate. 46. “ sodium iodate, 66 potassium iodide “ sodium nitrate. 47. “ potassium chloride, 66 sodium chloride “ potassium nitrate. 48. “ potassium carbonate, i ( oxalic acid “ potassium chloride. 49. “ ammonium sulphate. i 6 potassium nitrate “ sodium sulphide. 50. “ manganese peroxide, 6 6 lead permanganate. 51. “ arsenious sulphide, 6 6 lead arsenate “ potas’m thiocyanate 52. “ arsenious sulphide, 6 6 lead arsenate “ potassium sulphite. 53. “ arsenious sulphide. 6 6 silver arsenate “ barium sulphate. 54. “ potassium nitrite, 66 sodium nitrite “ potassium chloride. 55. “ lead ferrocyanide, 66 cupric ferrocyanide “ metallic lead. PART 111. SYSTEMATIC EXAMINATIONS. SEPARATION OP THE ACIDS PROM THE BASES. 931. The preliminary examination of the Solid Material in the dry way will give indications drawing attention to certain acids. Solutions can be evaporated to obtain a residue for this examination. Thus, detonation (not the decrepitation caused by water in crystals) indicates chlorates, nitrates, bromates, iodates. Explosion or deflagration will occur if these, or other oxygen-furnishing salts—as permanganates, chromates—are in mixture with easily combustible matter (858). Hypophosphites, heated alone, defla- grate intensely. A brownish-yellow vapor indicates nitrates or nitrites (730); a green flame, borates (636).—The odor of burning sulphur : sulphides, sulphites, thiosulphates, or free sulphur. The separation of carbon black :an organic acid. The formation of a silver stain : a sulphur compound (836). 932. When dissolving a solid by acids for work in the wet way, indications of the more volatile acids will be obtained. Sudden effervescence: a carbonate (oxalate or cya- nate) (664, 665). Greenish-yellow vapors: a chlorate (880), Brownish-yellow, chlor- nitrous vapors on addition of hydrochloric acid : a nitrate.—The characteristic odors : salts of hydrosulphuric acid, sulphurous acid, hydrobromie acid, hydriodic acid, hydro- cyanic acid, acetic acid.—The separation of sulphur : a higher sulphide, etc. It will be remembered that chlorine results from action of manganese dioxide, and numerous oxi- dizing agents, upon hydrochloric acid. 933. If the Material is in Solution, the bases will be first determined. (Certain volatile acids will be detected in the first-group aeidulation—by indications mentioned in the preceding paragraph.) Now, it should first be considered, what acids can be present in solution with the bases found? Thus, if barium be among the bases, we need not look for sulphuric acid, nor, in a solution not acid, for phosphoric acid. 934. As a general rule, the non-alkali metals must be removed from a solution before testing it for acids, unless it can be clearly seen that they will not interfere with the tests to be made. Metals need t.o be removed : because, firstly, in the testing for acids by precipitation, a precipitate may be obtained from the action of the reagent on the base of the solution tested, thus : if the solution contain silver, we cannot test it for sulphuric acid by use of barium chloride (and we are restricted to use of barium nitrate). And, secondly, in testing for acids by transposition with a stronger acid—the preliminary examination for acids—certain bases do not permit transposition. Thus,chlorides, etc., of lead, silver, mercury, tin, and antimony, and sulphide of arsenic, are not transposed by sul- phuric acid, or not promptly. 935. If neither arsenic nor antimony is among the bases, they may all be removed (a) by boiling with slight excess of sodium or potassium carbonate, and filtering. Arsenic and 280 Conversion of Solids into Liquids. 281 antimony, and all other bases of the second group, may be removed (6) by boiling with hydrosulphuric acid, and filtering. When the bases are removed by sodium or potas- sium carbonate, the filtrate must be exactly neutralized by nitric acid, with the expulsion of all carbonic acid by boiling. Then, for nitric acid, the original substance may be tested. The filtrate, from the third or fourth group, though free from all bases which need to be removed, is not suitable material for general tests for acids ; because it is loaded with ammonium salts, which act as solvents on many of the precipitates to be obtained. 936. The separation of Phosphoric acid from bases is a part of the work of the third group of metals, and is explained in 305 and 306. The removal of Boracic acid, Oxalic acid, Silicic acid, is described in 308. The non-volatile Cyanogen acids can be separated from bases by digesting with potassium or sodium hydrate (not too strong, 684), adding potassium or sodium carbo- nate and digesting, and then filtering. The residue is examined for bases, by the usual systematic process. The solution (677 c) will contain the alkali salts of the cyanogen acids, and may contain metals whose hydroxides or carbonates are soluble in fixed alkali hydroxides. CONVERSION OP SOLIDS INTO LIQUIDS. 937. Before the fluid reagents can be applied, solids must be reduced to liquids. To obtain a complete solution, the following steps must be observed : First. The solid, reduced to a fine powder, is boiled in ten times its quantity of water. Should a residue remain, it is allowed to'subside, and the clear liquid poured off or separated by filtration. A drop or tico evaporated on glass, or clean and bright platinum foil, will give a residue, if any portion has dissolved. If a solution is obtained, the residue, if any, is exhausted, and well washed with hot water. Second. The residue, insoluble in water, is digested some time with hot hydrochlo- ric acid. (Observe 983.) The solid, if any remain, is separated by filtration, and washed, first with a little of this acid, then with water. The solution, with the wash- ings, is reserved. Third. The well-washed residue is next digested with hot nitric acid. Observe if there are vapors of nitrogen oxides, indicating that a metal or other body is being oxi- dized (718). Observe if sulphur separates (564). If any residue remains it is separated by filtration and washing, first with a little acid, then with water, and the solution re- served. Sometimes it does not matter which acid is used first. But if a first-group base be present HN03 should be added first, for HCI would form an insoluble chloride. If the substance contain tin (especially an alloy of tin) HNOs would form insoluble raeta- stannic acid, HIOSn50i5, in which ease HCI should be used first. Fourth. Should a residue remain it is to be digested with nitro-hydrochloric acid, as directed for the other solvents. The acid solutions are to be evaporated nearly to dryness, and then redissolved in water, acidulating, if necessary, to keep the substance in solution. Fifth. Should the substance under examination prove insoluble in acids, it is likely to be either a sulphate (of barium, strontium, or lead); a chloride, or bromide, of silver or lead; a silicate or fluoride—perhaps decomposed by sulphuric acid—(745); and it must he fused with a fixed alkali carbonate, when the constituents are transposed in such man- ner as to render them soluble. The watery solution of the fused mass will be found to contain the acid; the residue, insoluble in water, the metal, now soluble in hydrochloric or nitric acids (compare 823). 282 Removal of Organic Substances. If more than one solution is obtained, by the several trials with solvents, the mate- rial contains more than one compound, and the solutions, as separated by filtration, should be preserved separately, as above directed, and analyzed separately. The sepa- rate results, in many cases, indicate the original combination of each metal. CONVERSION OF SOLUTIONS INTO SOLIDS: Before solids in solution can be subjected to preliminary examination—either for metals or for acids—they must be obtained in the solid state. This is done by evaporation. REMOVAL OF ORGANIC SUBSTANCES. 938. The methods of inorganic analysis do not provide against interference by or- ganic compounds; and, in general, it is impossible to conduct inorganic analysis in ma- terial containing organic bodies. The removal of the latter can be effected, Ist, by com- bustion at a red or white heat, with or without oxidizing reagents; 2d (in part), by oxi- dation with potassium chlorate and hydrochloric acid on the water-bath; Bd, by oxida- tion with nitric acid in presence of sulphuric acid, at a final temperature of the boiling point of the latter; 4th, by solvents of certain classes of organic substances; sth, by Dialysis. These operations are conducted as follows: 939. Combustion at a red or white of course, excludes analysis for mercury, arsenious and antimonious bodies (except as provided in 520), and ammonium. The last- named constituent can be identified from a portion of the material in presence of the organic matter (45 and 705). If chlorides are present iron will be lost at temperatures much above 100° C., and potassium and sodium waste notably at a white heat, and slightly at a full red heat. Certain acids will fe expelled, and oxidizing agents reduced. The material is thoroughly dried and then heated in a porcelain or platinum cruci- ble, at first gently. It will blacken, by separation of the carbon of the organic com- pounds. The ignition is continued until the black color of the carbon lias disappeared- In speciai eases of analysis, it is only necessary to char the material; then pulverize it, digest with the suitable solvents, and filter; but this method does not give assurance of full separation of all substances. Complete combustion, without use of oxidizing agents, is the way most secure against loss, and entailing least change of the material; it is, how- ever, sometimes very slow. The operation may be hastened, with oxidation of all mate- rials,by addition of nitric acid, or of ammonium nitrate. The material is first fully charred; then allowed to cool till the finger can be held on the crucible ; enough nitric acid to moisten the mass is dropped from a glass rod upon it, the lid put on, and the heat of the water-bath continued until the mass is dry, when it may be very gradually raised to full heat. This addition may be repeated as necessary. The ammonium nitrate may be added, as a solid, in the same way. 940. Oxidation with potassium chlorate and hydrochloric acid on the water-bath does not wholly remove organic matter, but so far disintegrates and changes it that the fil- trate will give the group precipitates, pure enough for most tests. It does not vaporize any bases but ammonium, but of course oxidizes or chlorinates all constituents. It is es- pecially applicable to viscid liquids ; it may be followed by evaporation to dryness and ignition, according to 989. The material with about an equal portion of hydrochloric acid is warmed on the water-bath, and a minute portion of potassium chlorate is added at short intervals, stir- ring with a glass rod. This is continued until the mixture is wholly decolored and dis- solved. It is then evaporated to remove chlorine, diluted and filtered. If potassium Preliminary Examination of Solids. and chlorine are to be tested for, another portion may be treated with nitric acid, on the water-bath. The organic matter left from the action of the chlorine or the nitric acid may be sufficient to prevent the precipitation of aluminium and chromium in the third group of bases; so that a portion must be ignited. As to arsenic and antimony, see 530. 941. The action of sulphuric with nitric acid at a, gradually increasing heat, leaves behind all the metals (not ammonium), with some loss of mercury and arsenic (and iron?) if chlorides are present in considerable quantity. In this, as in the operations be- fore-mentioned, volatile acids are lost—sulphides partly oxidized to sulphates, etc. The substance is placed in a tubulated retort, with about four parts of concentrated sulphuric acid, and gently heated until dissolved or mixed. A funnel is now placed in the tubule, and nitric acid added in small portions, gradually raising the heat, for about half an hour—so as to expel the chlorine, and not vaporize chlorides. The material is now transferred to a platinum dish, ami heated until the sulphuric acid begins to vapor- ize. Then add small portions of nitric acid, at intervals, until the liquid ceases lo darken by digestion, after a portion of nitric acid is expelled. Finally, evaporate off the sul- phuric acid, using the lowest possible heat at the close. 942. The solvents used are chiefly ether for fatty matter, and alcohol or ether, or both successively, for resins. Instead of either of these, benzene may be used; and many fats and some resins may be dissolved in petroleum naphtha. It will be observed, that ether dissolves some metallic chlorides, and that alcohol dissolves various metallic salts. Be- fore the use of either of these solvents upon solid material, it should be thoroughly dried and pulverized. Fatty matter suspended in water solutions may be approximately re- moved by filtering through wet, close filters; also, by shaking with ether or benzene, and decanting the solvent after its separation. 943. By Dialysis, the larger part of any ordinary inorganic substance can be ex- tracted in approximate purity from the greater number of organic substances in water solution. The degree of purity of the separated substance depends upon the kind of organic material. Thus, albuminoid compounds are almost fully rejected ; but saccha- rine compounds pass through the membrane quite as freely as some metallic salts. (Con- sult Watts’ Dictionary, 11. 316 ; 111. 715.) PRELIMINARY EXAMINATION OE SOLIDS. 944. Before proceeding to the analysis of a substance in the wet way, a careful study should usually be made of the reactions which the substance undergoes in the solid state, when subjected to a high heat, either alone or in the presence of certain reagents, before the blow-pipe, or in the flame of the Bunsen Burner. This examination in the dry way precedes that in the wet, and should be carried on systematically, following the plan laid down in the Tables, and noting carefully every change which the substance under in- vestigation undergoes, and if necessary making reference to some of the standard works- on Blow-pipe Analysis. In order to understand fully the nature of these reactions, the student should first acquaint himself with the character of the different parts of the- flame, and the use of the blow-pipe in producing the reducing and oxidizing flames. 945. The flame of the candle, or of the gas-jet, burning under ordinary circumstances, consists of three distinct parts : a dark nucleus or zone in the centre, surrounding the ■wick, consisting of unburnt gas—a luminous cone surrounding this nucleus, consisting of the gases in a state of incomplete combustion. Exterior to this is a thin, non-lumin- °us envelope, where, with a full supply of oxygen, complete combustion is taking place : here we find the hottest part of the flame. The non-luminous or outer part is called the- oxidizing flame; the luminous part, consisting of carbon and unconsumed hydrocar- bons, is called the reducing flame. 284 Preliminary Examination of Solids. 946. The flame produced hy the Blow-pipe is divided into two parts—the oxidizing flame, where there is an excess of oxygen, corresponding to the outer zone of the candle- flame, and the reducing flame, where there is an»excess of carbon, corresponding to the inner zone of the candie-flaine. Upon the student’s skill in producing these flames, de- pend very largely the results in the use of the blow-pipe. In order to produce a good oxidizing flame, the jet of the blow-pipe is placed just within the flame, and a moderate blast applied—the air being thoroughly mixed with the gas, the inner blue flame, corresponding to the exterior part of the candle-flame, is pro- duced : the hottest and most effective part is just before the apex of the blue cone, where combustion is most complete. The reducing flame is produced by placing the blow-pipe just at the edge of the flame, a little above the slit, and directing the blast of air a little higher than for the oxidizing flame. The flame assumes the shape of a non-luminous cone, surrounded by a pale-blue mantle ; the most active part of the flame is somewhat beyond the apex of the luminous cone. 947. The blast with the blow-pipe is not produced by the lungs, but by the action of the muscles of the cheek alone. In order to obtain a better knowledge of the manage- ment of the flame, and to practise in producing a good reducing flame, it is well to fuse a small grain of metallic tin upon charcoal, and raising to a high heat endeavor to pre- vent its oxidation, and keep its surface bright ; or better, perhaps, to dissolve a speck of manganese dioxide in the borax bead on platinum wire—the bead becoming amethyst- red in the outer flame and colorless in the reducing flame. The beginner should, work only with substances of a known composition, and not attempt the analysis of unknown complex substances, until he has made himself perfectly familiar with the reactions of at least the mo'i frequently occurring elements. The amount of substance taken for analysis should not be too large ; a quantity of about the bn Ik of a mustard-seed being, in most cases, quite sufficient. The physical properties of the substance under examination are to be first noted- such as color, structure, odor, lustre, density, etc. Heating in Glass Tube Closed at One End. 948. The substance, in fragments or in the form of a powder, is introduced into a small glass tube, sealed at one end, or into a small matrass, and heat applied gently, gradually raising it to redness, if necessary with the aid of the blow-pipe. When the substance is in the form of a powder it is more easily introduced into the tube by placing the powder in a narrow strip of paper, folded lengthwise in the shape of a trough ; the paper is now inserted into the tube held horizontally, the whole brought to a vertical posi- tion, and the paper withdrawn ; in this way the powder is all deposited at the bottom of the tube. By this treatment in the glass tube, we are first to notice whether the sub- stance undergoes a change, and whether this change occurs with or without decomposi- tion. The sublimates, which may be formed in the upper part of the tube, are especially to be noted. Escaping gases or vapors should be tested as to their alkalinity or acidity, by small strips of red and blue litmus inserted in the neck of the tube. Heat in Glass Tube Open at Both Ends. 949. The substance is inserted into a glass tube from two to three inches long, about one inch from the end—at which point a bend is sometimes made ; heat is applied gently at first, the force of the air-current passing through the tube being regulated by inclin- Preliminary Examination of Solids. 285 ing the tube at different angles. Many substances undergoing no change in the closed tube, absorb oxygen, and yield volatile aclns or metallic oxides. As in the previous case, the nature of the sublimate and the odor of the escaping gas are particularly to be noted. The reactions of sulphur, arsenic, antimony, and selenium, are very character- istic ; these metals, if present, are generally easily detected in this way. Heat in the Blow-pipe Flame on Charcoal. 950. For this test, a well-burned piece of charcoal is selected, and a small cavity made in that side of the coal showing the annular rings ; a small fragment of the sub- stance is placed in the cavity, and, if the substance be a powder, it may be moistened with a drop of water. The coal is held horizontally, and the flame made to play upon the assay at an angle of about twenty-five degrees. The substance is brought to a mode- rate heat, and finally to intense ignition. Any escaping gases are to be tested for their odor ; the change of color which the substance undergoes, and the nature and color of the coating which may form near the assay, are also to be carefully noted. Some sub- stances, as lead, may be detected at once by the nature of the coating. Ignition of the Substance previously Moistened with a Drop of Cobalt Nitrate. 951. This test may be effected either by heating on charcoal, in the loop of platinum wire, or in the platinum-pointed forceps. A portion of the substance is moistened with a,-drop of the reagent, and exposed to the action of the outer flame. When the substance is in fragments, and porous enough to absorb the cobalt solution, it may be held in the platinum- pointed forceps, and ignited. The color is to be noted after fusion. This test is rather limited; Aluminium, Zinc, and Magnesium giving the most characteristic re- actions. Fusion with Sodium Carbonate on Charcoal. P52. The powdered substance to be tested is mixed with the Soda, moistened, and placed in the cavity of the coal. Some substances form, with soda at a high heat, fusible compounds—others infusible. Many bodies, as silicates, require fusion with alkali car- bonate before they can be tested in the wet way. Many metallic oxides are reduced to metal, forming globules, which may be easily detected. When this test is applied for the detection of sulphates and sulphides, the flame of the alcohol-lamp is to be substituted for that of the gas-flame, as the latter generally contains sulphur compounds. Examination of the Color which may be imparted to the Outer Flame. 953. In this way many substances may be definitely detected. The test may be ap- plied either on charcoal or on the loop of platinum wire—preferably in the latter wTay. When the substance will admit a small fragment is placed in the loop of the platinum wire, or held in the platinum-pointed forceps, and the point of the blue flame directed upon it. If the substance is a powder it may be made into a paste with a drop of water, and placed in the cavity of the charcoal, the flame being directed horizontally across the coal. The color which the substance imparts to the outer flame in either case is noted. In most cases the flame of the Bunsen Burner alone will suffice; the substance being Preliminary Examination of Solids. heated in the loop of platinum wire—which, in all cases, should be first dipped in hydro- chloric acid, and ignited, in order to secure against the presence of foreign substances. Those salts which are more volatile at the temperature of the flame, as a rule give the most intense coloration. When two or more substances are found together, it is some- times the case that one of them masks the color of all the others—the bright yellow flame of Sodium, when present in excess, generally veiling the flame of the other elements. In order to obviate this, Bunsen has furnished us a method,* by the use of colored media (stained glasses, indigo solution, etc.) The appearance of the flame of various bodies, when viewed through these media, enables us often to detect very small quantities of them in the presence of large quantities of other substances. Treatment of the Substance with Borax and Microcosmic Salt. 954. This is best effected in the loop of platinum wire. This is heated and dipped into the borax or salt of phosphorus, and heated to a colorless bead: a small quantity of the substance under examination is now brought in contact with the hot bead, and heated, in both the oxidizing and reducing flames. Any reaction which tabes place during the heating must be noticed; most of the metallic oxides are dissolved in the bead, and form a colored glass, the color of which is to be observed, both while hot and cold. The color of the bead varies in intensity, according to the amount of the substance used ; a very small quantity will, in most cases, suffice. Certain bodies, as the alkaline earths, dis- solve in borax, forming beads which, up to a certain degree of saturation, are clear. When these beads are brought into the reducing flame, and an intermittent blast used, they become opaque. This operation is called flaming. As reducing agents, certain metals are employed in the bead of borax or salt of phos- phorus. For this purpose Tin is generally chosen—Lead and Silver being taken in some cases. These metals cannot be used in the loop of platinum wire, as they will alloy the platinum. The beads are first formed in the loop of wire ; then, while hot, shaken off into a porcelain dish, several being so obtained. A number of these are now taken on charcoal and fused into a large bead, which is charged with the substance to be tested, and then with the tin or other metal. For this purpose tin foil (or lead foil) is pre- viously cut in strips half an inch wide, and the strips rolled into rods. The end of the rod is touched to the hot bead to obtain as much of the metal as required. Lead may be added as precipitated lead (“proof-lead ”), and silver as precipitated silver. By aid of tin in the bead, cuprous oxide, ferrous oxide, and metallic antimony are obtained and other reductions effected, as directed in 348 and elsewhere. * For a full account of the method of analysis by flame reactions and colored media, suggested by Cabtmell ; and by films on porcelain, as developed by Bunsen ; consult Watts' Diet., Ist Supplement, p. 125; also Planner's Manual, Blow-pipe Anal., Richter, N. Y. Solids: Preliminary. I. Heat a portion, finely pul- 1. The Substance suffers no change: verized, in a Dry Glass Absence of volatile bodies (including combined water), of organic compounds, and of those which Tube closed at one end (948). change color on heating. 2. The Substance changes color: Organic compounds blacken from separation of carbon, which burns away. Cu and Co salts blacken at high heat. ZnO and most Zn salts, yellow while hot, white when cold. PbO and Pb salts, yellow while hot, yellow when cold. Bi2Q3 (white) and many Bi salts, orange to red-brown while hot, pale-yelloiv when cold. Fe2Q3, and salts, red to black while hot, reddish-brown when cold. Cd(OH)2 and many Cd salts, brown while hot, brown when cold. Sn02, brown while hot, yellow when cold. 3. The substance fuses: Most alkali salts and numerous other salts. Many salts dissolve in their water of crystallization when heated, becoming solid again by vaporization. 4. The substance subtimes, partially or wholly : H20 of crystallization, combination, or absorption. Sublimate condensing in cold part of the tube. Hg (449), gray, easily rubbed to globules. HgCl2 first melts, then forms white crystalline sublimate. Hg2Cl2, without melting, forms a sublimate, ydlow while hot, white when cold. HgS, a black sublimate, turning red on trituration. 955. PRELIMINARY EXAMINATION OF SOLIDS. Note physical properties, such as Structure, Gravity, Color, Odor, etc. TABLE I. 288 Solids: Preliminary. As, steel gray sublimate ; garlic odor. As203 sublimes in white octahedral crystals, does not fuse (472). As2S3, sublimate nearly black while hot, reddish-yellow when cold. Sb2S3 fuses yellow ; forms white, amorphous sublimates. NH4 salts, those not decomposing, white sublimate (46). Fe2Cl6 slowly sublimes as a reddish-yellow stain (220). S, free or by reduction of sulphide, gives reddish-brown drops, yellow when solidified. H2C204, a heavy white vapor and crystalline sublimate, I, a violet vapor and blue-black sublimate. 5. The substance evolves a gas or vapor : O indicates the presence of a nitrate, chlorate, bromate, iodate, or peroxide. A small piece of coal placed upon the assay glows upon being heated. H2S, from hydrated sulphides, some sulphites, blackens lead-paper. Recognized by its odor. S02, from sulphites, thiosulphates, certain sulphates, etc. Recognized by its odor and bleaching effect. NH3, from its compounds which decompose (46), odor, and alkaline reaction on litmus. ON, recognized by characteristic odor and violet flame. Oxides of Nitrogen, from nitrates or nitrites, reddish-brown, acrid vapor. Acetone, from acetates, characteristic fragrant odor. II. Heat in a Glass Tube open at both ends (949). Certain of the changes stated above as occurring in operation I. are modified by oxidation. Oxides are obtained from metals. S and Sulphides yield S02. Recognized by its odor and action on litmus-paper. As yields a sublimate of As203. Sb yields a sublimate (white), of Sb203 and Sb205. Bi, a sublimate, dark-brown while hot, lemon-yellow when cold (Bi203). Te, gray sublimate of tellurous acid (Te02). Se and Selenides evolve Se02, odor resembling that of rotten horse-radish (612). Hg, sublimate of metallic mercury. T.A BLE I.—Continued. Solids: Preliminary. III. Heat in the Blow-pipe Flame on Charcoal (950). 11. The substance decrepitates: Crystals containing water, as NaCl. (If finely pulverized, the decrepitation is avoided.) 2. The substance deflagrates : Nitrates, Chlorates, lodates, Hypophosphites, Permanganates, etc. 3. The substance fuses, and is absorbed by the charcoal: Salts of alkalies and some salts of alkaline earths. 4. The substance is infusible and phosphorescent: Ba, Sr, Ca, Mg—the residue is alkaline to test-paper. AI2O3, MgO, ZnO (yellow while hot), not alkaline to test-paper. 5. The substance forms an incrustation on charcoal: Pb, lemon-yellow while hot, sulphur-yellow when cold. In thin layers, bluish-white, volatile with bluish flame. Bi, dark orange-yellow while hot, lemon-yellow when cold. Zn, yellow while hot, white when cold, greenish-white flame. Cd, red brown, volatile, dark-yellow flame. As, white. Readily volatilized, distant from the assay, faint blue flame, Sb, white, pale-green flame. Sn, faint-yellow while hot, white when cold. IV. The Substance (or incrusta- tion of Test III., 5) is moist- ened with solution Cobalt Nitrate and strongly ignited (951). 1. The mass or incrustation is colored: ZnO—yellowish-green. SnO—bluish-green. Sb205—dirty dark-green, A1203, Si02, phosphates—blue. MgO—flesh-color or pink. BaO—brick red. SrO, CaO—gray. 290 Solids: Preliminary. V. Heat with Na2OOs, on char- coal in the inner Blow-pipe Flame (952). 1. Metallic grains are obtained: Bi, Sh—brittle. Pb, Cu, Sn, Aa—malleable. 2. An infusible magnetic powder is obtained: Fe, Ni, Co. VI. Heated in the Blow-pipe Flame, or in the Bunsen Flame on Charcoal, or in a loop of platinum wire (953). 1. The substance colors the outer flame: (If Test V. does not reduce metal, heat on platinum wire for flame color, before the blow-pipe or in Bunsen’s flame.) Yellow : Na and its salts, even in small quantities, impart an intense reddish-yellow. Other salts, even in large quantities, do not interfere with this reaction; viewed through a green glass, appears orange-yellow ; moistened with sulphuric acid, the test is more delicate (30). Violet: K and most of its salts, except borates, phosphates, and silicates, give bluish-violet flame, distinguished in presence of very small quantities of sodium compounds. Excess of the latter prevents the reaction ; Li also masks the reaction. In presence of sodium, the potassium flame appears reddish-violet when viewed through a blue glass (20). Red : Ca and its compounds produce a yellowish-red flame (139). Sr and many of its salts yield a crimson flame, masked by much Ba (123). Li and its salts produce a carmine-red flame (55). Sodium interferes with the reaction ; Potas- sium does not. Green : Yellowish-green—Ba and most of its salts. Also, Mo and its compounds. Emerald-green—Cu and most of its compounds (349). Bluish-green—B203 and phosphates. Yellowish-green—B2C3, best obtained by the addition of sulphuric acid. Heat on platinum wire until the sulphuric acid is expelled, then moisten with glycerine and ignite (638). Whitish-green—Zn, Blue: Light blue—As and many arsenic compounds. Azure-blue—Pb, Se. Also, CuCl2. Greenish-blue—Sb, CuBr2. TABLE I.—Continued. Solids: Preliminary. H 1 3 «. . . 55 so O § « 3 EC a o W •§ § ° 0 I® 1 5 §- § t -12 ►1 e „• (Q *•* e s>i .„ S ||l Sill CQ 3 O -p § •SI * O g >‘ prf § e Tn O * B M >5 -S P S 111 g }| =*r M ~ «0 a « - W I|-; ft •S?>| Oh S 5 £ *f i a *tg W § So 3 B ! « 2 E* |*- g, W .§ S o 3-S §r 2 1 W § 00 n &l CQ r2 3 H 8 '- ° c? a ■» 9 s So 15 ° e - ?! I o Eh M <0 '5 “ CQ rtf 2 S -1 a ” .«“ | J. Q. oo c« n -s> 0 W ri1 I 1 - 3 £ I 8 ■§ s s m •?<> * M §. a . § I* M VO 'O m *> o> vs so 'I Color of the With Microcosmic Salt or Phosphate of Ammonium-aodium With Sodium Pyroborate (Borax). Bead. In outer or oxidizing Flame. In inner or reducing Flame. In outer or oxidizing Flame. In inner or reducing Flame. Colorless. Si (swims undissolved). Al, Mg, Ca, Sr, Ba, Sn, (s. s., opaque). Ti, Zn, Cd, Pb, Bi, Sb, (not sat.) Si (swims undissolved). Al, Mg, Ca, Sr, Ba, (mps. not clear). Ce, Mn, Sn. h. c.; Si, Al, Sn, (sups, opaque). Al, Mg, Sr, Ca, Ba, Ag, (not sat.) Zn, Cd, Pb, Bi, Sb, Ti, Mo. Si, Al, Sn, (s. s., opaque). A'kaline earths and earths. h. c.; Mn, Ce. h.: Cu. Yellow or Brownish. h. (s. s.): Pe, Ur, Ce. c.; Ni. h.; Fe, Ti. c.; Ni. h., not sat.; Fe, Ur. h., sups.; Pb, Bi, Sb. h.; Ti, Mo. Red. h. (s. s.): Fe, Ni, Cr, Ce. c.; Cu. h.; Ni, Ti with Fe. h.; Fe, Ce. c.; Ni. a.; Cu (sups, opaque). Violet or Amethyst. h., c.; Mn. c.: Ti. h. c.; Mn. h.; Ni with Co. c.; Ti. Blue. h., c.; Co, c.: Cu. h., c.; Co. c.; W. h. c.: Co. c.; Cu. h. c.; Co. Green. h.; Cu, Mo; Be with Co or Cu. c.: Cr. c.; Cr. Cr. c.: Cr. h.; Ur, Mo. h.; Cu, Fe with Co. sups.; Fe. Gray and Opaque. Ag, Pb, Sb, Cd, Bi, Zn, Ni. The same as with Microcosmic salt. 292 Treatment of a Metal or an Alloy. TREATMENT OF A METAL OR AN ALLOY.* 957. On account of the different effect that Nitric Acid has upon the uncombined metals, it is used as a solvent in their detection. Thus : Gold and Platinum are not attacked by Nitric Acid. Tin and Antimony are oxidized and converted into compounds that are insoluble both in water and an excess of the acid. 6Sb + lOHNO3 = BSb.2Os + lONO + 5HzO 15Sn + SOHNOa + 5H.0 = 3HIOSnr,Ois + SONO All the other metals are oxidized and converted into compounds that dissolve either in water or an excess of the acid: e.g.: 3Pb + BHNO3 = BPb(NO3)2 + 2NO + 4H20 Bi + 4HNOs = Bi(No3)a +NO + 2H20 Method of Procedure.f Place a small quantity of the metal or alloy, about equal in bulk to a pea, having previously obtained it in as finely divided a state as possible, in an evaporating-dish, or any suitable vessel, cover well with Nitric Acid,,pp. gr. 1.20, and apply heat. Continue the application of heat, replacing from time to time the acid lost by evaporation, until the metal or alloy is dissolved or wholly disintegrated. If complete solution takes place immediately, pass on to A. If a residue remains decant the liquid portion upon a filter; again add Nitric Acid to the residue, heat, and again decant upon the same filter. Then thoroughly wash with hot water, either by boiling with water and decanting, or by transferring the whole to and pouring hot water through the filter. Add the first portions of the hot-water filtrate to the Nitric Acid filtrate already obtained, and treat the mixture as directed in A, after having, first, evaporated a drop or two on platinum foil, to ascertain whether anything has really been dissolved. Treat the residue as directed in B. A.—The Nitric Acid Solution. This solution may contain any of the metals, except those mentioned under B. If the Nitric Acid has effected a whole or partial solution of the original metal or alloy, evaporate almost to dryness to remove excess of acid, add about ten times its bulk of water, and proceed with the separation and detection of the metals in the regular way. Should the concentrated liquid become turbid when diluted with water, the presence of bismuth is indicated. In this case enough acid must be added to clear up the so- lution. % * This section is furnished by J. W. Baird, A.M., Professor of Analytical Chemistry in the Massachu- setts College of Pharmacy, recently Instructor in Assaying and Qualitative Analysis in Lehigh University. t When gold or platinum constitutes more than one-quarter of the alloy, Nitric Acid fails to extract the whole of the base metals that otherwise are readily soluble. In such a case the amount of gold or platinum must be reduced to at least 35 per cent, by fusing the alloy with the requisite amount of that base metal whose absence is surely known. $ Arsenic, if present in the original alloy, now exists in the form of Arsenic Acid, the precipitation of which requires heat and long-continued passage of H2S. Treatment of a Metal or an Alloy. 293 B.—The Residue Insoluble in Nitric Acid. This may contain gold and platinum in their metallic forms, and tin * and anti- mony* in the form of metastannic and antimonie acids. The separation of the two former from the two latter depends upon the fact that the metastannic and antimonie acids are soluble in hydi'ochloric acid, forming SnCl4 and SbCls. Digest, therefore, the well-washed residue in concentrated hydrochloric acid at a boiling temperature for from 5 to 10 minutes; then add at once an equal volume of water (to take the SnCl4 into solution), and bring to the boiling point. If gold or platinum existed in the original metal or alloy it will now be found in the form of a dark-brown or black powder or mass, insoluble in the hydrochloric acid. If such a residue exists, decani while hot, again add hydrochloric acid, heat, and again de- cant. This solution may have a turbid appearance, especially when cold, due to the action of the water upon the SbCls; but without filtering proceed with the separation and de- tection of the tin and antimony by the usual process.f The Hydrochloric Acid Solution. The Dark-Colored Residue. Add, after washing, two volumes of hydrochloric and one of nitric acid; evaporate almost or quite to dryness, dissolve in a small quantity of water (to obtain a concentrated solution), and divide into two portions. The gold and platinum have been dissolved by the aqua-regia formed, and now exist as auric and platinie chlorides. First Portion—Test for Gold. Dilute with at least ten times its bulk of water; add a drop or two of a mixture of stannous and stannic chlorides ; a purple or brownish-red precipitate (or coloration), purple of cassius, constitutes the test for gold. A convenient way of preparing this mixture of stannous and stannic chlorides is to (a) Add a few drops of chlorine-water to a solution of stannous chloride; or (b) Add to a small quantity of stannous chloride enough fex-ric chloride to produce a faint coloration. Second Portion—Test for Platinum. Add, without dilution, an equal volume of a strong solution of ammonium chloride. The formation, either at first or on standing, of a lemon-yellow crystalline precipitate, consisting of the double chloride of platinum and ammonium—(NH4Cl)2Ptol4—consti- tutes the test for platinum. Addition of alcohol favors the precipitation. If the proportion of platinum is very small, the mixture, after ammonium chloride has been added, should be evaporated to dryness on a water-bath and the residue treated with dilute alcohol. The ammonium platinie chloride remains behind as a yellow crys- talline powder. * Traces may sometimes be dissolved. t Arsenic must be looked for in this as well as in the nitric acid solution. For when the alloy contains arsenic, part of it will combine with the antimony and tin, and be held in the residue. Group mo of Bases. I. Add Hydrochloric acid, a drop at a time, as long as a precipitate is produced: warm, agitate, and filter (546). Precipitate: GROUP I. Lead, PbCl2, white. Silver, Af?C1’ white. Mercurosum, Hg2Cl2, white. II. To the Filtrate, or, if Group I. is absent, to original solution + HCl, add Ilydrosulphuric Acid in excess; warm and filter (553). Precipitate : GROUP II. Arsenic, As2S3, yellow. Antimony, Sb2S3 or Sb2S5, orange. Tin, SnS, brown. SnS2, yelloiv. Lead, PbS, black. Bismuth, Si2S3, black. Copper, CuS, black. Cadmium, CdS, yelloiv. Mercury, HsS' (First, white to yellow.) III. If phosphates may be present (305), test, after expelling H2S by boiling, for P04, by 768. In absence of Phosphates, proceed by 301 A and B {or 963). In presence of Phosphates, proceed by 963 or by 307. Precipitates: GROUP III. In absence of Phosphates: §■» rS § • f Aluminium, Al2(OH)6, white, gelatinous. o '§ I A i Chromium, Cr2(OH)6, bluish-green. §5 §-2 I Ferricum, Fe2(OH)6, reddish-brown. o f Manganese, MnS, flesh-colored. Cobalt, CoS, black. B.®gS -! Is, Nickel, NiS, black. cS, i Pq L Zinc, ZnS, white. The entire group, by Ammonium Sulphide, Chloride, and Hy- droxide, the same as .above, except iron, PeS, black. In presence of Phosphates ; By ammonium sulphide, etc „ ., „ All as above (iron as FeS) and Ba, Sr, Ca, Mg, Al, Cr, as Phosphates, white. IV. To the Filtrate from Group III, containing Ammonium Chloride, add Ammonium Carbonate and Hydroxide ; digest with gentle heat {not boiling) for some time, and filter. Precipitate : GROUP IV. Barium, BaC03, white. Strontium, SrC03, white. Calcium, CaC03, white. V. The Filtrate con- tains : GROUP V. Magnesium salts. Ammonium salts. Potassium “ Sodium “ Lithium “ Proceed by 959. For Rare Metals, see 573 to 613. Proceed by 965. For Rare Metals, see 58 and 69. VUULCimu0 urnmnte i / tyucocc uij results, see 547. 'Rare Metals : 573 to 613. For Rare Metals, see 310 to 331. Proceed by 964. 958. GROUPING OP THE METALS. (Compare 1 and 2. Remove each group before testing for the next.) Bases: Group I. 295 Solution; PbCl2 (990). Remaining Precipitate : AgOl, Hg2Cl2. Test for Lead (549) by Digest with dilute, warm Ammonium Hydroxide, and filter. Sulphuric acid, giving PbS04, white. Solution: (NH3)3(Ag01)2. Residue : NH2Hg2Cl, black (433). Hydrosulphuric acid, PbS, black. Test for silver (550), after expelling any excess [Lead Oxychloride, white, 548.] Chromate . PbCr04, yellow. of ammonium hydroxide by boiling, by acid- The black color is the evidence of mercury. Iodides Pbl2 (392). uldting slightly with nitric acid. A preci- Additional proof raav be obtained by test- The precipitates may be tested by the pitate is AgOl. ing the residue, or a portion of it, for mer- blow-pipe (397 and 398). For other tests, the silver chloride may be re- cury by drying and heating with sodium duced to the metal by zinc (430), or on char- coal before the blow pipe (423), and the metal dissolved by nitric acid (407), the resulting solution giving all the reactions of silver. If the original solution does not contain non- alkali metals, it may be used to obtain the characteristics of silver. carbonate, in a glass tube (449). For reactions of mercury in solution, if the original solution contains non-alkali me- tals, dissolve the residue with nitro-hydro- chloric acid—as mercuric chloride. For the study of lead, see the Table at 453 and Follow the Table at 453, and refer to the text Follow 453, and refer to the text at 346. Also the text at 387. at 411, see 960 “ B.” The Precipitate is washed on the filter with one or two small portions of cold water ; then treated with much hot water, and filtered. -00 iO cS 0 a -5 t*) H M 01 t) o pg 0 U, o 03 M CO § 03 in 03 Bases: Group 11. Wash the precipitate thoroughly on the filter, and then digest (in the test-tube) with yellow ammonium sulphide,* using as little as possi- ble. Treat the Residue, well-washed, as directed in “ B.” Precipitate the Filtrate with dilute hydrochloric acid (diluting icith solution of hydrosulphuric acid), and filter; examining this Precipitate,! when washed, for As, Sb, Sn, according to “ A.” [For Gold, 580; Platinum, 589; Rare Metals, 593 to 613.] “ A.” Precipitate from the Ammonium Sulphide Solution : As2S3, Sb2S5, SnS2, (S2). Treat the washed precipitate with hydrochloric acid, and, if it does not dissolve, add a minute fragment of potassium chlorate, and digest in a test-tube, till the sulphides are dissolved {or decoloi ed). Filter out any sulphur, washing the filter with a few drops of hydrochlo- ric acid, and receiving the filtrate in an evaporating-dish. Remove all the chlorine by a very gentle heat, until wet litmus-paper isnot bleached by the vapor. Transfer to the Gen era,tor of a Marsh’s Apparatus {carefully rinsing in any white residue of antimony or tin oxides) (558). Solution : X H3As04, SbClB SnCl4. Place in the Generator sufficient zinc {a piece of platinum foil or wire, 288) and dilute sulphuric acid (or use Na-amalgam, Al, or Mg, 463 and 534), and receive the gas in solution of silver nitrate {as long as a precipitate is produced). Now filter, and wash the residue. In the Generator: Sn, (Sb, Zn salt). Gather and wash the solid contents of the gen- erator ; dissolve with moderately dilute hy- drochloric acid, as directed in 535. (An undissolved residue may be Sb, soluble in nitro-hydrochloric acid.) Test the solution (SnCl2)/c»r tin by HgCl2; a white or gray precipitate (537). Precipitate, from SbHs: Ag-Sb, (Ag). Digest with warm hydrochloric acid to dis- solve the antimony (506), dilute {not to pre- cipitation), and test the solution for anti- mony by— H2S: an orange precipitate. § Water: a white precipitate (497). Spots and Mirror, from SbH3 (466). Solution, from AsH3: H3As03 (AgNOa, hno3). Remove the silver by adding just enough hy- drochloric acid and filtering (409). Test the solution for arsenic by— H2S: a yellow precipitate. 11 AgNOs: (467). Spots from the flame (466). Sublimation from the sulphide (471). Follow 545, and the text at 529 and after. Follow 545, and consult the text at 497, Follow 545, and consult 458 and after. * If Copper is present, and Mercury absent, it is better to use sodium sulphide (557). t A white precipitate, of sulphur, will occur in any case. Colored precipitates indicate As, Sb, or Sn (556). t In many cases, 1 he original solution can be used for additional results through Marsh’s Test (463). § A black precipitate shows that too much HC1 was used, and some silver dissolved. 1 A black precipitate shows that too little or too much HC1 was used. 960. ANALYSIS OF GROUP 11. (Explanation at 555.) Bases : Group 11. If sulphides have been found in “ A,” the precipitate should be washed, first with (NH4)2Ss (or Na2S), then with water. Dissolve the pre- cipitate by action of hot, moderately dilute nitric acid (564). Evaporate the solution to expel excess of nitric acid. Residue: Solution; Pb(NC3)2, Bi(NC3)3, Cu(NQ3)2, Cd(NC3)2. HgS, S [PbSO.i]. Dissolve the (black) pre- cipitate by nitro-hy- drochloric acid, and To a portion of the solution, add a drop of very dilute sulphuric acid. If a precipitate appears (565), add the reagent to the whole solution, and filter. Precipitate; Filtrate; Bi, Cu, Cd, salts. Add ammonium hydroxide to a slight alkaline reaction. expel free chlorine by boiling (571). Solution: HgCl2. Test for mercury by Stannous chloride (448 a). Ammonia (443). Copper (448 d). Sublimation (449). PbSC4. Add Chromate, and heat (567). Or add KI (392). (Lead chromate dis- solves in potassium hydroxide, and repre- cipitates with acetic acid) Test the pre. of sul- phates, or sulphide, by reduction on char- coal (397). Precipitate: Bi(OH)3. Filter out and wash the precipitate and test it by stannite, according to 366. Or, test by dissolving with a very little Blue Solution : Copper (340), cadmium (374). Concentrate in an evaporating- dish, acidify with acetic acid, and confirm with ferrocyan- ide; a, broivn precipitate, Cu2Fe(CN)6. Test this, or the original solution, with metal- lic iron (346). Colorless Solution: Cadmium, absence of copper. To the concentrated (acidified or alkaline) solution, add H2S : a yellow precipitate, CdS. To the neutral solution add K4Fe(CN)6: Cd2Fe(CN)6, white. Cyanides (374). HC1 on a watch- glass and adding For copper, follow 876 and 339. For Cadmium, see the text at 376 water (358). If copper is present, treat the sul- phide obtained by adding H2S Follow 452, and the text at 442 and after. See 959. Follow 452, and the text at 387. Follow 376, and the text at 357. to the ammoniacal solution, by one of the methods given in 570, for separation of cad- mium. * If silver was not removed in the first group, it may be tested for, in the nitric solution of this precipitate, by HC1. 961. “ B.” SULPHIDES NOT SOLUBLE IN AMMONIUM SULPHIDE : PbS, Bi2S3, OuS, CdS, HgS * Bases: Group 111. To the filtrate from the second group (958), add NH4OH till alkaline, then add NH4CI and Ammonium Sulphide. Dissolve the well-washed Precipitate, on the Filler, by dilute, cold hydrochloric acid (303). Concentrate in an evaporating-dish. Solution : ZnOL, AlaCle, PeOl2 (314), IVInClj, OraOle, NiOla, COCI2. Residue : CoS, NiS, black. Test black residue by the blow-pipe Jor cobalt and nickel. Dissolve by nitric add, and test for Ni, by 280. Add potassium hydroxide solution to strong supersaturation and boil a short time (292). Filter, and wash. Separation of Zn and Al. Solution: K2ZnC2; K2A1204. Precipitate: Fe(OH)2 and Fe2(OH)6, Mn(OH)2, Cr2(OH)6, Co(OH)2, Ni(OH)„. Divide the solution into two portions. Examine five portions ; two in solution (for Fe and Ni), and three in precipitate (for Cr, Mn, Co), 1. For Zinc. 2. For Aluminium. 1 and 2. For Iron and for Nickel. 3. For Manganese. 4. For Chromium. 5. For Cobalt. Add ammonium sul- phide; a precipi- tate, ZnS (294). This 'precipitate may be dissolved by HC1, or, in absence of in- terfering bases, the original solution taken, for addition- al tests for zinc. Add Ammonium chloride in decided excess (393). A pre- cipitate : Al2(OH)6. Dissolve and test this solution, or test the original solution, if free from interfer- ing bases, for alu- minium. Dissolve with a drop or two of ni- tro-hydrochloric acid, and ex- pel all free chlorine. Test a portion for Iron ly potas- sium thiocyanate : a red solu- tion, Fe2(CNS)6. Test the original solution for Fer- rosum and Ferricum bi/ ferricy- anide and thiocyanate, 221, Boil with Pb02 and HNOs, as directed in 248 ; a red color, HoMnoOg. Or fuse with Na2C03 and KN03, as directed in 349 ; a green mass, manganate, as K2MnC4. {If Mn is present, proceed by 301, with the filtrate from 2d group.) Fuse with Na2CC3 and KN03 (187), dissolve in water, acidify by acetic acid and test for chromic acid, by 191, 192, 198 Test, also, the original solution for chro- mic acid (182). Test in the bo- rax bead (269). For Iron, follow the text at 204 and after. Follow 300, and the text at 392. Follow 300, and the text at 167. Test a portion for Nickel, by 380 or 281. Follow 300, and con- sult the text at 243 and after. Follow 300, and the text at 181 and after. Consult the text at 264. 962. ANALYSIS OF GROUP 111,, IN ABSENCE OF PHOSPHATES. The Scheme of Separations given in 301 is preferred to this, in most Cases. Bases: Group 111. To the Filtrate from Group 11., add ammonium chloride, and ammonium hydroxide to alkaline reaction, then ammonium sulphide to complete the precipitation. Reserve the Filtrate for Group IY. Precipitate: FeS, MnS, CoS, NiS, ZnS, S; Al.2(OH)6, Or2(OH)6; phosphates of Al, Or, and of Ba, Sr, Ca, Mg. Treat the well-waslied precipitate with cold dilute hydrochloric acid. If a black residue remains, test it for Co and Ni, as directed in the Table at 801 B. Also, this residue may be tested for Si (746). The filtered solution is boiled to expel H2S filtered if turbid, and a smaller portion reserved. Portion 1. Add a few drops of nitric acid and boil. Solution: Fe2Cl0, Al2Clr„ Cr2Cl6, MnCl2, [CoCl2, NiCl2], ZnCl2, H3P04, (H2SiQ3); also Ba, Sr, Oa, Mg, chlorides. (1) Test a small portion for iron, by thiocyanate. If iron is found, test the original solution by ferricyanide and thiocyanate as directed in 221. (2) To the remainder, add ferric chloride till a drop is precipitated yellow by ammonium hydroxide (shoeing that the P04 is all precipitated), concentrate to a small bulk, add water, nearly neutralize by K2C03, and add excess of barium carbonate. Let the mixture stand and filter. Portion 2. Solution: BaCl2, SrCl2, OaCl2, MgCl2, etc. Add sulphuric acid and filter. Precipitate: BaS04, SrS04 (CaS04). Test the precipitate for Sr, by the flame. [Test it for Ba by fusing with K2C03 (823), then dis- solving the well-washed residue by acetic acid and obtaining reactions with chromate, etc.] To the filtrate, add alcohol. A precipitate may be OaS04. Boil with water and add NH4OH and oxalate, for calcium. Precipitate: Al2(OH)6, Cr2(OH)6, [Fe.(P04)2, Fe2(OH)6, BaCOs]. Boil the precipitate for some time with sodium or potassium hydroxide. Solution: Mn012, ZnCl2, BaCl2, SrCl2, CaCl2, MgCl2. Add HOI and boil to expel C02. Add NH4OH to alkaline reaction, then ammonium sulphide, and warm and filter. Solution K2A1204. Acidify the solution with HC1 and add excess of ammonium hydroxide, and boil. Precipitate: Al2(OH)6. Precipitate: Cr2(OH)6, etc. Test for chromium as directed in 301 (p. 70). Precipitate: MnS, ZnS. Dissolve by HC1* and test the solution (MnCl2ZnCl2) by the directions in 301 (b), p. 70. [Or separate by Acetic Acid, 293.] Filtrate: BaCl2, MgCl2. Remove Ba, Sr, Ca, by adding H2S04, as directed in Portion 2, above, then add ammonium oxalate and hydroxide, filter, and tef-t the filtrate for Mg, by adding phosphate. Follow the text at 166, * If a black residue appear, test for cobalt and nickel, as directed in foot-note 301 b. 963. ANALYSIS OF GROUP 111., IN PRESENCE OF PHOSPHATES. (Explanation in 305.) [The Scheme at 307 may be used instead of this one ] Bases • Group IV. Dissolve the well-washed precipitate in dilute e.cetic acid. Solution: Ba(C2Hs02)2, Sr(C2H302)2, Ca(C2H302)2. To a sma’I portion of the solution, add potassium dichromate ; if a precipitate appears, add the reagent to the whole solution as long as a precipitate is produced (105), and filter. Precipitate: BaCr04, yellow. Solution; Sr(C2H302)2, Ca(C2H302j2, [KoCraOr]. The precipitate is soluble by hydrochloric acid, and this solution is precipitated by sul- phuric acid, as BaSC4, insoluble in acids. Precipitate by ammonium carbonate with ammonium hydroxide; filter, and wash the pre- cipitate, and dissolve it by acetic acid. Solution: Sr(C2H302)2, 0a(C2H302)2. Divide in two portions. See 100 and after. For Strontium : For Calcium : To a portion, add solution of calcium sul- phate ; boil, and leave for about ten min- utes. A precipitate indicates strontium, SrS04. Test another portion of the solution on a loop of platinum wire, by the flame (955 YI.) Add solution of potassium sulphate (and fil- ter), to insure the absence of strontium. To the filtrate {or solution not precipitated), add ammonium oxalate. A precipitate : CaC204, insoluble in acetic acid, soluble in hydrochloric. Test by the flame (955 VI.) See the text at 120. Follow the text from 132 to 140. 964. ANALYSIS OF GROUP IV. (Explanation in 142.) [The Method of Separation given in 146 may be used instead of this.] [Concerning the loss of traces of barium and calcium, see 144.] Basks: Group V. Evaporate a drop or two on clean platinum foil by a gentle heat (9). If a residue is obtained, ignite. If the residue or a part of it vaporizes immediately, ammonium compounds are indicated. (The melting affixed alkali salts is liable to be mistaken for vaporization.) For Ammonium, test a portion of the original material with calcium, or potassium, or so- dium hydroxide, at a gentle heat. The vapor of ammonia, NH3, is recognized hy its effect on moistened litmus-paper, its odor, etc. (89). Conduct the gas into water, and add potassium mercuric iodide witii KOH (40). Test another portion of the same solution by HgCl2 (41). If no inter- fering acids or bases are pre- sent, the K2HgI4 with KOH may be added to the original solution. For Magnesium, test the Fil- trate from Group IV. Add a little ammonium hydrox- ide; then enough ammonium chloride to dissolve any pre- cipitate which appears (79), and then sodium phosphate Precipitate : MgNH4P04. Concerning lithium phosphate, see 58. Calcium, in traces, may appear here, as phosphate (144). To guard against this fallacy, add oxalate to the filtrate from the fourth group, and filter out any precipitate so made. If magnesium is present, its compounds lessen the delicacy of the Flame Tests for the Metals of the Alkalies, and wholly prevent tests by precipitation. To remove Magnesium, precipitate it hy barium hydroxide and then remove the barium, as directed in 150. For Potassium, test by the flame (on platinum wire, with the Bun sen burner); violet, not obscured by blue glass (20). Test by Tartaric acid, using a concentrated solution. For Sodium, test {on platinum wire) for the flame color, yellow, and given by extrane- ous traces of sodium compounds (82). Ob- scured by the blue glass. For Lithium, test by the flame, crimson, and obscured only by very thick blue glass (55). Test by sodium phos- phate. The precipi- tate, Li3P04, is solu- ble in much ammo- nium chloride (dis- tinction from magne- sium). Consult 35 to 47. Folloiv the text at 76 and after. Follow the text at IQ to 21, and 958. (See 955.) Follow the text at 28 and after, and 955 VI. 965. THE ANALYSIS OF GROUP V. Acids: First Table. 966. PRELIMINARY EXAMINATION FOR ACIDS. Concerning the indications of acids in the Blow-pipe Examination, see 931 and 955 ; the indications in dissolving Solids, see 932 ; the considerations relative to any Bases already determined, see 933 ; the Removal of Bases, see 934-936 ; Reactions in the Wet Way, see 967, 968. For the Vaporous Acids : If the vapor reddens moist blue litmus-paper, some one or more of the following-named volatile acids is indicated: I. To a little of the solid 1. There is sudden effervescence : substance or residue by Carbonic anhydride, CO„. Colorless and odorless gas, feebly reddening litmus, and evaporation, or very con- making solution of calcium hydroxide turbid. See 665. If the original sub- centrated solution, in a stance contains non-alkali bases in solution, carbonates cannot be present. Solu- test-tube, add a little tions of carbonates give precipitates with salts of all non alkali metals, 667. concentrated Sulphuric Acid, and heat gently, Oxalates (651) and cyanates (696) also evolve C02. not enough to vaporize 3. Effervescence of gas having odor: the sulphuric acid. Hydrosulphuric acid, H2S. Blackens paper wet with lead acetate. Most sulphides For carbonates, sulphides, are transposed by hydrochloric acid with evolution of H2S ; decomposed by nitric and cyanides, it is better acid or chlorine with‘separation of S. Sulphides in solution precipitate salts of to add dilute sulphuric the first three groups of bases (781). See 779 and test by II. acid. Sulphurous anhydride, S02. Odor of burning sulphur. Bleaches litmus. Colors iodic acid and starch (809). Sulphites precipitate salts of all non-alkali bases. Follow the text at 803. Thiosulphates, also, evolve sulphurous anhydride, on decomposition—sulphur being separated, 786. Decomposed by all acids. Form precipitates (788), soluble in excess. Acids: First Table. 303 Hydrocyanic Acid, HCH. Peach-blossom odor. Gas precipitates silver nitrate, on a glass rod, as AgCN. Change to thiocyajiate, gives blood-red color with ferric salts (682); to ferrocyanide, gives blue color with iron salts (681). Alkali cyanides are decomposed by all acids. In solution, they precipitate most of the second and third group metals (679). Acetic acid, H(C2H302). Odor of vinegar. Digested with alcohol and sulphuric acid, gives odor of acetic ether (636). Acetates are transposed by hydrochloric and nitric acids. In solution, form no precipitates, but give red solution with ferric salts (635). Hydrochloric acid, HC1. Slight effervescence. Slight, irritating odor. The gas forms a white precipitate (AgCl) with solution of silver nitrate, on a glass rod ; the precipitate being insoluble in dilute nitric acid, but soluble in ammonium hy- droxide (843). The gas forms a white cloud with vapor of ammonia. Obtain Cr02Cl2 by 845. Chlorides precipitate salts of all the first-group bases (842). Compare the reactions by table at 918. For separations, see 920* and 928. 3. Appearance of gas having color : Hydriodic acid, HI. Odor is offensive, and slightly chlorine-like. Vapor colors violet in the air—both odor and color due to separation of free iodine. Vapor and mixture color starch blue. Carbon disulphide extracts iodine, violet. Soluble iodides are transposed by HC1 and decomposed by Cl and by HN03. In solution, give colored precipitates with salts of the first-group metals (see 897 to 903). Compare reactions by table, 918. For separations, see 920 and 928. Hydrobromic acid, HBr. Odor is acrid and chlorine-like. Color, slightly yellowish- brown—odor and color due to separation of free bromine. Starch colors yellow ; * For an application of the chlorochromic test, by Prof. Wiley, see Chem. News, xli. 176, April 16, 1880. 304 Acids: First Table. Carbon disulphide yellow ; iodine being absent in both tests. Soluble bromides transposed by HC1 and decomposed by HNO, and by Cl. In solution, give yellow- ish-white or white precipitates with salts of the first-group bases (872). Compare reactions, by aid of the table at 918. For separations, see 920 and 928. Nitrous acid, HNC2, and nitric peroxide, NC2, characteristic acrid odor. Color, red-brown. Color potassium iodide and starch, blue. Salts all soluble. See the text at 709. The brown gas is produced very sparingly from nitrates (unless reducing agents are present), but produced abundantly from nitrites. See special Tests (III.) Chloric acid, HC103. Detonation with sulphuric acid. Odor of chlorine. Color, greenish-yellow. Gas bleaches litmus. Chlorates decomposed by hydrochloric acid (861) and by nitric acid. By ignition, reduced to chlorides. Form no precipi- tates, except by reduction (compare reactions by the table at 918). Chlorine, Cl, is evolved in the decomposition of hypochlorites (853), and in that of chlorides with oxidizing agents (844). II. Fuse with pure Sodium Carbonate, on charcoal, and test the mass ob- tained. 1. When moistened, it blackens silver: Indications of sulphur in any combination or uncombined (783). Fusion with sodium carbonate on platinum foil or in a porcelain capside. The mass blackens silver : a sulphide, a thiosulphate, or one of the thionates. The mass does not blacken silver : a sulphate. 2. When acidified with hydrochloric acid, filtered, and the filtrate evaporated to dryness and treated with the same acid, an insoluble residue is obtained: Silicic anhy- dride, SiC2. Soluble in boiling solutions of fixed alkalies. See 746 and the bead-test (956). Acids: First Table. 305 III. Certain Tests not clas- sified. 1. For Nitric add and Nitrous acid: Trial to be made, in case of any material soluble in water (731 and 711). Only slight portions of nitrites could escape recognition by the brown gas evolved in Test I., as above mentioned. Nitrates, with reducing agents, evolve brown gas abundantly, in Test I. Test for Nitric and nitrous acids, by formation of the “ brown ring” (734). Tests for Nitrous acid by 710 and 709. 3. For Boric acid: Borates of non-alkali metals, insoluble (634). See Flame Test, 955 VI. Test by 636 and 638. 3. For Silicic acid: Compounds insoluble, except those giving a very strong alkaline reaction. See test in the bead, 956. See Test II., 3, above. 4. For Ferrocyanic, Ferricyanic, and TMocyanic acids: Heated with sulphuric acid, evolve HCN (687). Thiocyanates, fused with KNOs, form soluble sulphates. Insoluble salts of H4Fe(CN)6, and H(Fes(CN)ia, boiled with solution of KOH, form soluble potassium salts (677 c). When in solution, test with iron salts (331). For solubilities, ferrocyanides, see 689 ; ferricyanides, 693 ; thiocyanates, see 700. The CN of single cyanides is separated by slightly acidulating, adding calcium carbonate, and distilling—ferrocyanogen and ferricyanogen and thiocyanogen being retained as salts. Acids: Second Table. 967. EXAMINATION FOR THE MORE COMMON NON-VAPOROUS ACIDS, AS BARIUM AND CALCIUM SALTS. Concerning the Removal of Bases, before Testing for Acids, see 934. Heat a portion of the neutralized solution,* and add both barium chloride and calcium chloride. Precipitate: BaSC4, CaC204, CaHP04, or Ca3(P04)2, etc. Also, barium and calcium borates, silicates, fluorides (740), ferrocyanides, sulphites, chromates, arseniates, iodates—the acids of which will be identified by 966, or during the work for bases. Also, tar- trates (645), citrates (640), and salts of other organic acids not described in this work. Concerning Selenates, Tellurites, see 612; Molybdates, 605. Digest the precipitate icith dilute hydrochloric acid. If a precipitate remains, filter. Residue: BaS04 (white). Solution; H20204; H3P04 or 0aH4(P04)2, etc.; barium and calcium chlorides; other acids of See 822. less frequent occurrence, as indicated above. Sulphates precipitate lead and stron- Add ammonium hydroxide to a slight alkaline reaction. Precipitate: Ca0204, CaHP04, etc. tium salts. Add acetic acid to a distinctly acid reaction. For comparison of the chief reactions, see 823. Residue: Ca0204 (white). Solution: H3P04, CaH4(P04)2, etc. (761). See 655. Test by molybdate, as directed in 768. See 79. Preliminary examination for acids, The deportment of H.As04 in this table resembles that of 966, Test I., Result 1. H,P04, and is liable to be mistaken for it. * For Sulphuric acid, the reaction may he acid; for oxalic acid, the solution may have acetic acidulation; but for phosphoric acid it must not have an acid reaction. Acidity should be neutralized by ammonium hydroxide, which may be added in slight excess. Acids: Third Table. 307 To the neutral or slightly acid solution* add silver nitrate. Precipitate: Silver salts of a very large number of acids, including those in the residue next below, together with silver carbonate, sul- phide, tartrate, and other precipitates dissolved by nitric acid (409). Digest with dilute hot nitric acid. Residue: AgCl, Agßr, Agl, AgON, AglOs,f AgßrOs, AgCNS. Add Ammonium Hydroxide to a strong alkaline reaction (931) Residue: Agl, indicating Iodides. Test the original solution for iodides, by chlorine-water with carbon disulphide (or starch), as directed in 902 and 928. Folloiv the Table at 918, and consult the text at 920, 928, and 907. Solution: Ammonium silver compounds, representing HOI, HBr, HON, HBrOs, HIOs, HONS. Add dilute nitric acid to a slight acidulation. Precipitate; AgCl, AgBr, AgCN, AgIOs, AgBr03, AgONS. Test the original solution for chloride. By distillation for chro- mium dioxide dichlo- ride (845). By the solubilities of the silver precipitate in am- monium hydroxide and carbonate (843). Test the original solution for bromide, as direct- ed in 873 and in 920. Test the original solution for iodate, by reducing agents {911), and by pre- cipitations (918). Test the original solution for bromate, as direct- ed in 881. Test the original solution for cyanide, 681 or 682. Test the original solution for thiocyanate, by 700. Observe 701. Follow the Table of Comparative Reactions, at page 274. * If alkaline, the solution should be neutralized by nitric acid. Should effervescence result, boil until it is complete. Should a precipitate form, remove it by filtration, t By repeated digestion with HN03 the AgJOs may all be dissolved. 968. EXAMINATION OF SOME OF THE ACIDS PRECIPITATED BY SILVER NITRATE. 308 Solubilities of Metallic Salts. SOLUBILITIES OF THE SALTS OF EACH ACID. 969. Concerning salts insoluble in water, it is stated by what acids they are trans- posed, and from this it will be seen by what acids they may be dissolved, i.e., changed to compounds soluble in water. As to the solution of salts insoluble in water (and acids) by decomposition with alkalies, see 833. For more specific statements as to decomposing and dissolving agents, refer to the descriptions of the acids in question in the text. Acetates. All soluble in water; Silver and Mercurous are sparingly soluble. Arsenates. Closely resemble the (ortho)phosphates, both in solubilities and in transpo- sition with acids. Arsenites. Those of the Alkali bases are soluble. Those of Barium and Strontium, sparingly soluble ; the others are insoluble in water, but transposed by dilute acids. Borates. Only those of the Alkali bases are freely soluble in water ; many of the others being slightly soluble. They are transposed by all acids, except carbonic. Some of the metals form non-normal borates. Bromates. All soluble in water ; Silver, Lead, and Mercurous, sparingly soluble. Bromides. Silver, Lead, and Mercurous, insoluble ; Mercuric, sparingly soluble ; Bis- muth, instable ; all others soluble in water. The bromides insoluble in water are scarcely transposed with sulphuric acid, or with dilute nitric acid (693). Carbonates. Those of the Alkali bases only are soluble in water. The acid carbonates less abundantly than the normal. Most of the others are made slightly soluble by free carbonic acid. Carbonates are transposed by all acids, except hydrosulphuric and hydrocyanic. The pseudo-triads do not form carbonates ; some other heavy metals form basic carbonates in the wet way. Chlorates. All soluble in water. Potassium chlorate but moderately soluble. Chlorides. Silver and Mercurous, insoluble; Lead, slightly soluble; all others soluble in water ; antimonious, stannous and bismuth, soluble in acidulated water. Chromates. Those of the bases of the Alkalies, and Magnesium, Calcium, and Zinc, are soluble in water; Strontium and Mercuric, sparingly soluble; nearly all others insoluble. Iron, Manganese, and Copper form chromates not normal—some of which are soluble in water, but chiefly instable in solution. Citrates. Those of the Alkali bases are freely soluble in water ; of Iron, Copper, and Zinc, moderately soluble; the other {single) citrates, mostly insoluble; the double citrates mostly soluble. The insoluble citrates are transposed by dilute mineral acids. Cyanides. Mostly insoluble in water; except those of the Alkali and Alkaline earth metals, and double cyanides containing these. Barium cyanide is sparingly soluble. Cyanides are transposed by nearly all acids, even when dilute. Ferricyanides. Those of the Alkali and Alkaline earth bases are soluble in water; that of Barium, sparingly. A considerable number of the others are insoluble in water, and certain of the bases do not form ferricyanides. See 693. They differ as to transposition with acids, but those insoluble are transposed by alkalies. Ferrocyanides. Those of the Alkali bases and of Magnesium, Calcium (not the potassio calcium), and Strontium are soluble in water. See 689. Those insoluble differ as to transposition by acids, but are transposed by alkalies. Fluorides. Those of the Alkali bases are freely soluble in water; those of the Alkaline Solubilities of Metallic Salts. 309 earth metals insoluble ; of Copper, Bismuth, Cadmium, Ferricum, and Zinc, sparingly soluble; Silver, Tin, and Mercuric, soluble. The insoluble fluorides are transposed by strong sulphuric acid, and less easily by hydrochloric and nitric acids. Hypochlorites. All soluble in water. (Decomposed by all acids.) Hypophosphites. All soluble in water. (Decomposed by nearly all acids.) lodates. Only those of the Alkali bases are freely soluble; the others insoluble, or sparingly soluble. Calcium, sparingly soluble ; Barium, Silver, and Lead, insolu- ble. Transposed by moderately dilute mineral acids—those of Silver and Lead by nitric acid not dilute. lodides. Silver, Lead, Mercurous, Mercuric (and Palladous), insoluble in water. Bis- muth, and to some extent Copper iodides, are decomposed by water without solu- tion, The others are soluble. The insoluble lodides are transposed with diffi- culty, or not at all, by sulphuric acid or nitric acid. (894.) Nitrates. All soluble in water. Nitrites. All soluble in water ; Silver, sparingly. Nitrophenates. All soluble in water; Potassium, Ammonium, and Lithium, very sparingly ; most others, more freely. Oxalates of the Alkali bases are soluble; Chromium and Stannic oxalates, soluble ; Magnesium and Ferric oxalates, sparingly soluble; the others chiefly insoluble or slightly soluble. Transposed by sulphuric, hydrochloric, and nitric acids, not by acetic. Permanganates. All soluble in water; Silver, sparingly. A number of the bases de- compose the acid radical. Phosphates)ortho-). Of the di- and tri-metallic salts, only those of the ordinary Alkali bases are soluble %n water. (Lithium, insoluble.) Those two-thirds hydrie (“ acid phosphates”) are all soluble in water, to some extent. Acetic acid transposes most of the insoluble phosphates, except those of Iron, Aluminium, and Lead; and dilute hydrochloric, nitric, and sulphuric acids transpose all phosphates (partly or wholly, 761). Pyrophosphates are insoluble in water, except those of the common Alkali bases. They are scarcely at all transposed by acetic acid, but yield their bases to the stronger acid radicals. Metaphosphates of the common Alkali bases, only, are soluble in water. They are not transposed with acetic acid, and some of them not readily by other acids when dilute. Silicates. Those of the Fixed Alkali bases, only, are soluble in water. These, in solu- tion, are transposed by all acids. Of the silicates insoluble in water, many are transposed with hydrochloric or sulphuric acid, but the larger number of the natural silicates resist acids. All are decomposed by hydrofluoric acid, and by the fixed alkalies. Sulphates. Those of Barium, Lead, Strontium. Calcium, are insoluble in water, the last-named being slightly soluble. Argentic and Mercurous sulphates are spar- ingly soluble. Mercuric, Antiraonious, and Bismuth sulphates require acidulated water for solution. All others are soluble in water. Sulphates are not transposed with acids, at ordinary temperatures. Sulphites. Those of the Alkali bases are soluble; all others insoluble, or very sparingly soluble in water. Those of the Alkaline earth metals are somewhat soluble in 310 Solubilities of Metallic Salts. solution of sulphurous acid. All sulphites are transposed by acetic and the mine- ral acids. Sulphides. Of the bases of the Alkalies and Alkaline earths, soluble; the others insolu- ble in water. The earth metals do not form sulphides. Sulphides of the third- group metals are transposed with dilute acids; those of the second group metals (except Mercury), transposed or decomposed by hydrochloric, nitric, and sulphuric acids. Thiocyanates. Those of Alkali and Alkaline earth bases, and of Iron, Manganese, Zinc, Cobalt, and Copper, are soluble in water. Mercuric, sparingly soluble. The others are transposed by dilute acids. Tartrates. Those of the Alkali bases are soluble in water, the acid tartrates of Potas- sium, Ammonium, Rubidium, and Caesium but sparingly soluble. Manganous, Ferric, Cobalt, Stannous, and Antimonious tartrates are soluble: Calcium tartrate, slightly soluble. The other tartrates, not soluble in water, are mostly somewhat soluble in solution of tartaric acid, and mostly soluble in solutions of Alkalies (as double tartrates); also transposed by the mineral acids. Thiosulphates. All soluble in water; those of Barium, Lead, and Silver sparingly soluble in water, but made soluble as double salts. Decomposed by all acids, 786. Solubilities of Metallic Salts. 311 Aluminium. Ammonium. Antimonious. j Barium. Bismuth. Cadmium. Calcium. | Chromium. Cobalt. Copper. Gold (triad). Hydrogen. Iron (dyad). Iron (triad). Lead. Magnesium. Manganous. Mercuric. Mercurous. Nickel. Potassium. Silver. Sodium. Stannic. Stannous. Strontium. Zinc. Acetate S S S S S S S S S S S S S S S S S s S S s S s S S s Arsenate i S i i i i i i i i S i i i i i i i i S i S i i i i Arsenite s i s i i i i i i i i i i i s i s i s Borate i s i i s i i i i s i i i s i i s i s i i i Bromide S s S S Si S S S S S S s S S si S S s i S s I s S S S s Carbonate s i i i i i i i i i i i i i s i s i i Chlorate S s S S S S s S S s S S S S S S S S s S s S s Chloride s s Si S Si S S s S S S s S S si S S S i S s I s s Si s s Chromate s i i i i s i i s s i s s s % i s i s i s s Cyanide s s s i s i i i i s i i s i S i i s i s s i Ferricyanide s s i s I I s I S s s I i I s I s i i Ferrocyanide s s I s I I s I I i s i i I s I s i i s i Fluoride i s S i s s i s s s s s s s i i S S s s S s S S i 8 Hydroxide i s S i i s i i i i s i i i i i i s s i i i i Iodide s Si S Si S S S S Si i s S S i S S i i S s i s Si S S S Nitrate S s s Si s S s s S s S S S S S Si S S s S s Si Si S S Oxalate i s i i i i i s i i s i s i s i i i i s i s S i i i Oxide i i i i s i i i i s i i i i i i i i s i s i i i i Phosphate i s i i i i i i i i s i i i i i i i i s i s i i i i Sulphate S s Si I Si S si S S S s S S 1 S S Si s S s s s S S I S Sulphide s i S i i S i i i s i i i S i i i i s i s i i S i S, soluble in water, s, sparingly soluble in water, i, insoluble in water, but made soluble by acids. I, insoluble in water or acids, s I, sparingly soluble in water, not transposed by acids. S i, soluble in acidulated water. 970. SOLUBILITIES OF METALLIC SALTS. REAGENTS* [Aq. = H2O as crystallization water,] 971. Acid, Acetic, H(02H302). Sp. grav. 1.04. 30 p.c. acid. Hydrochloric, HOI. Sp grav. 1.12. 24 p.c. acid. Hydrosulphuric, H2S. Water saturated with the acid (777). Hydrofluosilicic, (HF)2SiF4. (741.) Nitric, HNOs. Sp. grav. 1.2 (32 p.c. acid). Nitro-hydrochloric, NOCI2 + Cl. About one part of concentrated Nitric to 3 parts of Hydrochloric acid. Nitrophenic, HO0H2(N02)30. (728.) Oxalic, H2C2C4. 2 aq. Crystals dissolved in 10 parts of water. Sulphuric, H2SO4. Concentrated, sp. grav. 1.843. Tartaric, H2(C4H406). Crystals dissolved in 3 parts of water. Chlorine-Water, Cl. Water saturated with chlorine (833). Alcohol, C2HOO. Sp. grav. .815. About 95 p.c. Ammonium Chloride, NH4CI. One part crystallized salt in 8 parts of water. Ammonium Carbonate, (NH4)2CO3, One part of crystallized salt in 4 parts water, with one part of solution of Ammonia. As a solvent for arsenious sulphide, the reagent is prepared without the addition of solution of ammonia, (NH4)4H2(C03)3. Ammonium Hydroxide, NH4OH. Sp. grav. .96. 10 p.c. NH3. Ammonium Molybdate, (NH4)2MoC4. Solution in Nitric acid. Ammonium Sulphide, (NH4)2S, colorless; (NH4)2S2 or NH4HS, yellow; solution of ammonia, treated with hydrosulphuric acid. Ammonium Oxalate, (NH4)2C ,04. One part of the crystallized salt (aq.) in 24 parts of water. Barium Chloride, BaCl2, One part of the ci'ystallized salt (2 aq.) to 10 parts of water. Barium Carbonate, BaCO3. (103.) Barium Hydroxide, Ba(OH)2. A saturated water solution (88). Barium Nitrate, 8a(N03)2. One part to 15 of water. Calcium Chloride, CaCl2. One part salt (6 aq.) dissolved in 8 parts of water. Calcium Hydroxide, Ca(OH)2. A saturated water solution (131), also the dry solid. * In the greater number of cases, reagents should be “ chemically pure.” Different uses require dif- ferent degrees of purity. An article of sodium hydroxide contaminated with chloride may be used in some operations; not in others. Those who have had training in analysis can do without specific directions, which cannot be made to cover all circumstances ; and the beginner must depend on others for the selection of reagents. 312 Rea gents. 313 Calcium Sulphate, CaS04. A saturated water solution (88). Carbon Disulphide, CS2. (873.) Cobaltous Nitrate, C0(N03)2. One part crystallized salt (5 aq.) dissolved in 8 parts of water. Copper Sulphate, CuSC4. One part of the crystallized salt (5 aq.) in 8 parts of water. Ether, (02H6)20. Sp. grav. not over .738—containing not over 5 p.c. alcohol. Ferrous Sulphate, FeSC4. One part crystallized (7 aq.) in 5 parts of water. Ferric Chloride, Fe2Cl0. One part of the solid salt (6 aq.) to 15 parts of water. Gold Chloride, AuCl3. Prepared by dissolving pure gold—which may be obtained by precipitation with Oxalic acid—in nitro-hydrochloric acid, evaporating to dryness on the water-bath, and dissolving in water. Lead Acetate, Pb(C2H3C2)2. One part of the crystallized salt (3 aq.) dissolved in 10 parts of water. Magnesium Sulphate, MgSC4., One part of the crystallized salt (7 aq.) to 10 parts of water. Mercuric Chloride, HgCl2. One part of the crystallized salt in 1G parts of water. Mercurous Nitrate, Hg2(NC3)2. One part of the crystallized salt (2 aq.) dissolved in 20 parts of water, acidulated with one part nitric acid, or prepared by dissolving mer- cury. Palladous Chloride, PdCl2. One part of the salt to 20 of water. Potassium Chromate, K2Cr04. One part dissolved in 10 parts of water. Potassium Dichromate, K2Cr2C7. One part dissolved in 10 parts of water. Potassium Chlorate, KCI03. The crystallized salt. Potassium Cyanide, KCN. One part dissolved in 4 parts of water. Potassium Ferrocyanide, K4Fe(CN)0. One part of the crystallized salt (3 aq.) dissolved in 13 parts of water. Potassium Ferricyanide, KOFe2(CN)]2. One part dissolved in 13 parts of water. Potassium lodide, KI. One part dissolved in 20 parts of water. Potassium Mercuric lodide. Kessler’s Solution. Dissolve 3.5 grams of KI in 10 c.c. of water ; dissolve I.G grams of HgCl2 in 30 c.c. of water; add the mercury solution gradually, and with constant stirring, to the potassium iodide solution, until the precipitate ceases to be redissolved ; then add 60 c.c. Potassium hydroxide solution and filter. Keep in small bottle, well stoppered. Potassium Nitrate, KNC3. The crystallized salt. Potassium Metantixnonate, KSbOs. (514.) Potassium Thiocyanate, KCNS. One part dissolved in 12 parts of water. Potassium Hydrogen Sulphate, KHSC4. Potassium Sulphate, K2S04. One part dissolved in 13 parts of water. Platinic Chloride, PtCl4. One part to 10 parts of water. Also prepared by dissolving the scrap-metal in nitro hydrochloric acid, and purifying by precipitation with am- monic chloride, dissolving again in the same acid, and evaporating to dryness. Sodium Acetate, Na(C2H3C2). One part crystallized salt (3 aq.) to 5 of water. Sodium Carbonate, Na2CC3. The dry salt. Also a solution of the crystals (10 aq.) in 5 parts of water. Sodium Diborate, Na20(82C3)2. The crystallized salt (10 aq.), or dried. 314 Rea gents. Sodium Hydroxide, NaOH. Solution in 9 parts of water. Sodium Thiosulphate, Na2S2O3 5H20, (Hyposulphite.) One part of the salt in 40 parts of water. Sodium Hypochlorite, NaOlO. Agitate one part of good bleaching-powder with ten parts of water; add solution of sodium carbonate as long as a precipitate is formed; allow the solid matter to subside, and siphon off. Sodium Phosphate, Na2HPQ4. (Disodium hydrogen phosphate.) One part of the crys- tallized salt (12 aq.) in 10 parts of water. Sodium Phosphomolybdate. (604.) Sodium Sulphide, Na2S. One part of the solution of soda saturated with Hydrosul- phurie acid, to one part unchanged soda solution. Sodium Sulphite, Na2503. One part of the salt to 5 parts of water. Silver Nitrate, AgNC3. One part crystallized salt in 20 parts of water. Stannous Chloride, SnCl2. One part of the crystallized salt (2 aq.) in 6 parts of water, acidulated with hydrochloric acid (529). Strontium Sulphate, SrS04. A saturated water solution (88). Zino, Zn. The granulated metal should be platinized according to 388. INDEX. Acetic Acid 176 Acids 9 Acid Salts 12 Acids, Preliminary Examination for... 302 “ Tables for the Determination of 302 Aluminium 40 Alkalies, Precipitation by 18 Alkali Metals 17 Alkaline Earth Metals 29 Ammonium 33 Analytical Tables 294 to 307 Anhydride defined 9 Antimonic Compounds 133 Antimonious “ 128 Antimony 127 Arsenic 115 “ Compounds 124 Arsenious “ 116 Arsenic and Antimony Spots 120 Atomic Weights, Table of 14 Basic Salts 12 Barium 30 Bead Reactions 291 Beryllium 77 Bismuth 90 Blow-pipe, Tests by 291 Bonds, Negative 165 “ and Structural Formulae 166 Boron and Boric Acid 173 Bromic Acid 363 Bromine 357 Cadmium 94 Caesium 26 Calcium 34 Carbon 175 “ Monoxide 179 Carbonic Anhydride. 183 Cerium '. 77 Chloric Acid 253 PAGE Chlorine 244 “ Group 245 “ Peroxide 253 Chlorous Aeid 253 Chromic Acid 45 Chromium... 43 Chlorochromic Acid 249 Citric Acid 177 Classification of Metals... 16 Cobalt 59 Copper 84 Cyanic Acid 195 Cyanogen 187 Decipium 81 Didymium 82 Dith ionic Acid 233 Dithionous Aeid 232 Barths 46 Equations, Rule for Balancing 169 Erbium 82 Ferricyanic Acid 193 Ferrocyanic Acid. 191 Fifth Group 17 First Group 83 Flame Reactions 291 Fluorine 212 Fourth Group 29 Gallium 81 Germanium 163 Glucinum—see Beryllium 77 Gold 153 Group I. and II 83 “ 111 38 “ TV 29 “ y 17 Grouping of Metals 16 Hydracids 9 Hydriodic Aeid 267 Hydrobroraie Acid 259 PAG® 315 316 Index. Hydrochloric Acid 247 Hydrogen 171 Hydrocyanic Acid 187 Hydrofluoric Acid 213 Hydrofluosilicie Acid 214 Hydrosulphuric Acid 226 Hydroxide defined 13 Hypobromous Acid 262 Hypoehlorous Acid 251 Hypophosphorous Acid 218 Hyposulphurous Acid—see Thiosul- phuric Acid 230 Indium 82 lodic Acid 271 lodine 264 Iridium 164 Iron 46 Lanthanum 82 Lead 96 Lithium 24 Magnesium 27 “ Separation of, from Alkalies 37 Manganates and Manganites 55 Manganese 54 Mercury . .. 105 Metals, Classification of 16 Molybdenum 160 Nickel 63 Niobium 82 Nitric Acid 203 “ Oxide 200 Nitrogen 198 “ Group 199 “ Peroxide 203 Nitro-hydrochloric Acid 250 Nitrous Acid.. 200 “ Oxide 200 Normal Salts 11 Norwegium 163 Notation of Metallic Compounds 9 Organic Substances, Removal of 282 Osmium 164 Oxyacids 9 Oxalic Acid 180 Oxidation, Balancing Equations by... 169 “ Bonds 165 “ Valence compared with Structural Valence 166 Oxygen 209 0z0ne.... . . 210 PAGE Palladium 158 Pentathionic Acid 234 Perchloric Acid 257 Periodic Acid 2G4 Permanganic Acid 55 Peroxide of Hydrogen 211 Phosphoric Acid 221 Phosphorous Acid 221 Phosphorus 216 “ Hydride 218 Platinum 156 Potassium 19 Preliminary Examinations 283 Pyrophosphoric Acid 222 Quanti valence 10 Reagents, List of 312 Rhodium 164 Rubidium 26 Rule for Balancing Equations 169 Ruthenium 164 Salts, Notation of 9 Samarium 81 Scandium 81 Selenium 164 Separation of First and Second Group. 141 ‘ ‘ Third-Group Metals.... 70, 72 “ Fourth-Gi'oup Metals 36 “ Bases without using H2S. 151 “ the Acids of Chlorine... 275 “ of the Thionic Acids 235 Silicic Acid 215 Silicon , 214 Silver 101 Sodium 21 Solubilities, List of 308 Strontium 33 Sulphocyanie Acid—see Thiocyanic... 196 Sulphur 225 Sulphuric Acid 240 Sulphurous Acid 236 Table of Atomic Weights 14 Tables, Analytical 70, 75, 294 to 307 Tantalum 82 Tartaric Acid 178 Tellurium 164 Terbium 81 Tetrathionic Acid 233 Thallium 80 Thiocyanic Acid 196 i Thiosulphuric Acid 230 PAGE Index. 317 PAGE PAGE Third-Group Bases 38 Thorium 83 Tin 135 Titanium 79 Trithionic Acid 333 Tungsten 164 Uranium 78 Valence 165 Vanadium 83 Yttrium 83 Ytterbium 81 Zinc 65 Zirconium 83 Catalogue of the Scientific Publications of D. 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No. 78. THE STEAM-ENGINE INDICATOR, AND ITS USE, By W. B. Le Van. No. 79. THE FIGURE OF THE EARTH. By Frank C. Roberts, C.E. No. 80. HEALTHY FOUNDATIONS FOR HOUSES. By Glenn Brown- SCIENCE SERIES. No. 81. WATER METERS : COMPARATIVE TESTS OF ACCURACY, DELIVERY, ETC. Distinctive features of the Worth- ington, Kennedy, Siemens, and Hesse meters. By Ross E. Browne. No. 82. THE PRESERVATION OF TIMBER BY THE USE OF ANTISEPTICS. By Samuel Bagster Boulton, C.E. No. 83. MECHANICAL INTEGRATORS. By Prof. Henry S. H. Shaw, C.E. No. 84. FLOW OF WATER IN OPEN CHANNELS, PIPES, CONDUITS, SEWERS, ETC. With Tables. By P. J. Flynn, C.E. No. 85. THE LUMINIFEROUS ffiTHER. By Prof, de Volson Wood. No. 86. HAND-BOOK OF MINERALOGY; DETERMINATION AND DESCRIPTION OF MINERALS FOUND IN THE UNITED STATES. By Prof. J. C. Foye. No. 87. TREATISE ON THE THEORY OF THE CON- STRUCTION OF HELICOIDAL OBLIQUE ARCHES. By John L. Culley, C.E. No. 88. BEAMS AND GIRDERS. Practical Formulas for their Re- sistance. By P, H. 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TRIPLE-EXPANSION ENGINES AND ENGINE TRIALS. By Prof. Osborne Reynolds. Edited, with notes, etc., by F. E. Idell, M. E. SCIENCE SERIES. No. TOO. HOW TO BECOME AN ENGINEER ; OR, THE THEORETICAL AND PRACTICAL TRAINING NECESSARY IN FITTING FOR THE DUTIES OF THE CIVIL ENGINEER. The Opinions of Eminent Authorities, and the Course of Study in the Technical Schools. By Geo. W. Plympton, Am. Soc. C.E. No. ioi. THE SEXTANT AND OTHER REFLECTING MATHEMATICAL INSTRUMENTS. With Practical Suggestions and Wrinkles on their Errors, Adjustments, and Use. With thirty- three illustrations. By F. R. Brainard, U.S.N. No. 102. THE GALVANIC CIRCUIT INVESTIGATED MATHEMATICALLY. By Dr. G. S. Ohm, Berlin, 1827. Translated by William Francis. With Preface and Notes by the Editor, Thomas D. Lockwood, M.I.E.E. No. 103. THE MICROSCOPICAL EXAMINATION OF POTA- BLE WATER. With By Geo. W, Rafter. No. 104. VAN NOSTRAND’S TABLE-BOOK FOR CIVIL AND MECHANICAL ENGINEERS. Compiled by Geo. W. Plympton, C.E, No. 105. DETERMINANTS, AN INTRODUCTION TO THE STUDY OF. With examples. By Prof. G. A. Miller. No. 106. TRANSMISSION BY AIR-POWER. Illustrated. By Prof. A. B. W. Kennedy and W. C. Unwin. No. 107. A GRAPHICAL METHOD FOR SWING-BRIDGES. A Rational and Easy Graphical Analysis of the Stresses in Ordinary Swing-Bridges. With an Introduction on the General Theory of Graphi- cal Statics. 4 Plates. By Benjamin F. Laßue, C.E. No. 108. A FRENCH METHOD FOR OBTAINING SLIDE- VALVE DIAGRAMS. 8 Folding Plates. By Lloyd Bankson, 8.5., Assist. Naval Constructor, U.S.N. No. 109. THE MEASUREMENT OF ELECTRIC CURRENTS. Electrical Measuring Instruments. By Jas. Swinburne. Meters for Electrical Energy. By C. H. Wordingham. Edited by T. Commerford Martin. Illustrated. No. no. TRANSITION CURVES. A Field Book for Engineers, containing Rules and Tables for laying out Transition Curves. By Walter G. Fox. No. in. GAS-LIGHTING AND GAS-FITTING, including Specifica- tions and Rules for Gas Piping, Notes on the Advantages of Gas for Cooking and Heating, and useful Hints to Gas Consumers. Second edition, rewritten and enlarged. By Wm. Paul Gerhard.