UNIVERSITY OF ILLINOIS 
 
 CHEMISTRY DEPARTMENT 
 
 ARTHUR WILLIAM PALMER 
 MEMORIAL LIBRARY 
 1904 
 
 ^ . 3 
 

 r... 
 
 u. 
 

Digitized by the Internet Archive 
 in 2017 with funding from 
 
 University of Illinois Urbana-Champaign Alternates 
 
 https://archive.org/details/victorvonrichterOOrich 
 
A TEXT-BOOK 
 
 OF 
 
 INORGANIC CHEMISTRY 
 
 RICHTER 
 
Richter’S Chemistries 
 
 AUTHORIZED TRANSLATIONS 
 
 BY EDGAR F. SMITH, M.A., Rm.D., Sc.D. 
 
 PROFESSOR OF CHEMISTRY IN THE UNIVERSITY OF PENNSYLVANIA; MEMIIER OF THE CHEMICAL 
 SOCIETY OF BERLIN, AMERICAN PHILOSOPHICAL SOCIETY, ETC., ETC. 
 
 Professor Richter’s methods of arrangement and teaching have proved their 
 superiority by the iarge sale of his books throughout Europe and America, trans- 
 lations having been made in Russia, Holland, and Italy. The success attending 
 tlieir publication in this country could only have been attained for good books that 
 have been found useful, practical, and tlioroughly up to the times. 
 
 The Chemistry of the Carbon Compounds, 
 or Organic Chemistry 
 
 Third American, translated from the Eighth German Edition by 
 Prof. R. Anschutz, University of Bonn. Thoroughly Revised, 
 Enlarged, and in many parts Rewritten. Illustrated. Two Volumes. 
 
 Vol. I. Aliphatic Series. 625 pages. . Cloth, net, 1^3.00 
 
 Vol. II. Carbocyclic and Heterocyclic Series. 671 pages. 
 
 Cloth, net, $3.00 
 
 INORGANIC CHEMISTRY 
 
 Fifth American from the Tenth German Edition by PROF. H. 
 Klinger, University of Konigsberg. Thoroughly Revised and in 
 many parts Rewritten. With many Illustrations and Colored Plate 
 of Spectra. Cloth, net, J^i.75 
 
 *** Special Catalogues of Books on Chemistry, Hygiene, Pharmacy, Medicine, 
 Dentistry, and Nursing sent free upon application. Correspondence solicited. 
 
 P. BLAKISTON’S SON & CO., PUBLISHERS 
 
 PHILADELPHIA 
 

 
 f 
 
 
 LIBRARY 
 Of THE 
 
 UNIVERSITY of ILLINOIS 
 
Lehman L BoltarrL,PhilcL, 
 
 Philadelphia^ PBlahistorvs Son S. Co. 
 
VICTOR VON RICHTER’S 
 
 TEXT-BOOK 
 
 OF 
 
 INORGANIC CHEMISTRY 
 
 EDITED BY 
 
 PROF. H. KLINGER 
 
 UNIVERSITY OF KOENIGSBERG 
 
 AUTHORIZED TRANSLATION 
 
 BY 
 
 EDGAR F. SMITH 
 
 PROFESSOR OF CHEMISTRY IN THE UNIVERSITY OF PENNSYLVANIA, PHILADELPHIA 
 
 (Assisted by WALTER T. TAGGART) 
 
 INSTRUCTOR IN CHEMISTRY 
 
 jfittb Hmerican trom tbe Uentb (Berman lEbition 
 
 CAREFULLY REVISED AND CORRECTED 
 
 WITH SIXTY-EIGHT ILLUSTRATIONS ON WOOD 
 
 AND 
 
 COLORED LITHOGRAPHIC PLATE OF SPECTRA 
 
 PHILADELPHIA 
 
 P. BLAKISTON’S SON & CO. 
 
 1012 WALNUT STREET 
 
 1900 
 
Copyright, 1900, by P. Blakiston’s Son & Co. 
 
 WM. F. FELL & CO., 
 ELECTHOTVPEH 8 AND PHINTER8, 
 1320-24 8ANS0M 8THEET, 
 PHILADELPHIA. 
 
PREFACE 
 
 
 TO THE 
 
 FIFTH AMERICAN EDITION. 
 
 The student of the present edition will discover that it differs very 
 materially from all preceding editions. This is largely due to the fact 
 that the editor has endeavored to give due consideration to the more 
 recent, well-established discoveries in chemical science ; hence additions 
 will be found relating to the general properties and the measurement of 
 gases, to the atmosphere and the interesting constituents lately observed 
 in it, to the theory of dilute solutions and electrolytic dissociation, to the 
 electrolysis of salts, to alloys, etc. Thus revised, it is hoped that the 
 book will continue to occupy the position it has so long held among 
 works devoted to the inorganic portion of chemical science. 
 
 The translator would take this opportunity to acknowledge his great 
 indebtedness and to return his sincere thanks to Mr. Walter T. Taggart, 
 upon whom devolved the task of arranging the crude manuscript for the 
 press and the revision of the proof-sheets. 
 
 The John Harrison 
 Laboratory of Che77iistry. 
 
 V 
 

PREFACE 
 
 TO THE 
 
 FIRST AMERICAN EDITION. 
 
 The success of Prof, von Richter’s work abroad would indicate its 
 possession of more than ordinary merit. This we believe true, inasmuch 
 as, in presenting his subject to the student, the author has made it a 
 point to bring out prominently the relations existing between fact and 
 theory. These, as well known, are, in most text-books upon inorganic 
 chemistry, considered apart, as if having little in common. The results 
 attained by the latter method are generally unsatisfactory. The first 
 course — that adopted by our author — to most minds would be the more 
 rational. To have experiments accurately described and carefully per- 
 formed, with a view of drawing conclusions from the same and proving 
 the intimate connection between their results and the theories based upon 
 them, is obviously preferable to their separate study, especially when they 
 are treated in widely removed sections or chapters of the same book. 
 Judging from the great demand for von Richter’s work, occasioning the 
 rapid appearance of three editions, the common verdict would seem to 
 be unanimously in favor of its inductive methods. 
 
 In the third edition, of which the present is a translation, the Periodic 
 System of the Elements, as announced by Mendelejeff and Lothar 
 Meyer, is somewhat different, in the manner of development and pre- 
 sentation, from that appearing in the previous editions. This was done 
 to give more prominence to and make more general the interesting rela- 
 tions disclosed by it. Persons examining this system carefully will be 
 surprised to discover what a valuable aid it really has been, and is yet, in 
 chemical studies. Through it we are continually arriving at new rela- 
 tions and facts, so that we cannot well hesitate any longer in adopting it 
 into works of this character. It is, indeed, made the basis of the 
 
 vii 
 
viii PREFACE TO THE FIRST AMERICAN EDITION. 
 
 present volume. In accordance with it, some change in the treatment 
 of the metals, ordinarily arbitrarily considered, has been made. 
 
 A new feature of the work, and one essentially enlarging it, is the 
 introduction of the thermo-chemical phenomena, briefly presented in the 
 individual groups of the elements and in separate chapters, together with 
 the chemical affinity relations and the law of ])eriodicity. “Hereby 
 more importance is attributed to the ])rinciple of the greatest heat 
 development than at present apjiears to belong to it, because it was 
 desired, from didactic considerations, by the explanation of the few 
 anomalies, to afford the student the incentive and o[)])ortunity of deduc- 
 tively obtaining the majority of facts from the thermal numbers, on the 
 basis of a simple principle, do facilitate matters, there is appended to 
 the volume a table containing the heat of formation of the most im- 
 portant. compounds of the metals.” 
 
 Trusting that the teachings of this work will receive a hearty welcome 
 in this country, and that they will meet a want felt and often expressed 
 by students and teachers, we submit the following translation of the 
 
 same. 
 
TABLE OF CONTENTS 
 
 INTRODUCTION. 
 
 Physics and Chemistry, 17. Physical and Chemical Phenomena, 1 8, 19. Chemical 
 Elements, 19. Principle of Indestructibility of Matter, 20. Principle of Conserva- 
 tion of Energy, 21. Forms and Equivalents of Energy, 22. Chemical Energy, 23. 
 Constitution of Matter ; Atom and Molecule, 24. Chemical Symbols and For- 
 mulas, 25. Atomic Weights, 26. Table of Atomic Weights, 26. Chemical Equa- 
 tions, 27. Conditions of Chemical Action, 28. T hermo-chemical Phenomena, 28. 
 Crystallography, 31. 
 
 HYDROGEN AND THE NON-METALS. 
 
 Classification of the Elements, 39. 
 
 Hydrogen, 40. Purifying and Drying of Gases, 42. Apparatus for the Generation 
 and Collection of Gases, 42. Physical Properties of Hydrogen, 43. Chemical 
 Properties of Hydrogen, 45. Condensation of Gases, 47. Critical Condition, 47. 
 Group of Halogens, 49. 
 
 Chlorine, 49, Bromine, 53. Iodine, 54. Fluorine, 56. General Characteristics 
 of the Halogens, 56. Compounds of the Halogens with Hydrogen, 57. Hydrogen 
 Chloride, 57. Acids ; Bases ; Salts, 60. Hydrogen Bromide, 61. Hydrogen Iodide, 
 63. Hydrogen P'luoride, 64. General Characteristics of the Hydrogen-halogen 
 Compounds, 65. Thermo-chemical Deportment of the Halogens, 66. Compounds 
 of the Halogens with one another, 68. 
 
 Weight Proportions in the Union of the Elements; Stoechiometric Laws; Atomic 
 Hypothesis; Choice of Atomic Weights, 69. General Properties of Gases; Atomic- 
 molecular Theory, 73. Avogadro’s Law, 75. Status nascens, 77. Principles of 
 the Atomic-molecular Theory, 79. Determination of the Atomic Weights, 80, 
 Oxygen Group, 80. 
 
 Oxygen, 80. Oxyhydrogen, 83. Oxidation and Reduction, 84. Ozone, 84. 
 Isomerism and Allotropy, 87. Compounds of Oxygen with Hydrogen, 88. Water, 
 88. Natural Waters ; Chemical Properties of Water, 91. Electrolysis of Water ; 
 Thermo-chemical Deportment, 92. Dissociation, 93. Kinetic Theory of Gases, 94. 
 Quantitative Composition of Water, 95. Molecular Formula of Water ; Atomic 
 Weight of Hydrogen and of Oxygen, 97. Hydrogen Peroxide, 99. Catalysis; 
 Thermo-chemical Dejiortment, 103. 
 
 Sulphur, 104. 
 
 Molecules of the Elements, 106. Hydrogen Sulphide, 107. Hydrogen Persulphide, 
 1 10. Compounds of Suljihur with the Halogens, no. Selenium, 112. Tellurium, 
 1 1 3. Summary of the Elements of the Oxygen Group, 113. Thermo-chemical 
 Deportment, 114. 
 
 IX 
 
X 
 
 TAI’.LF, OF CONTENTS. 
 
 Nitrogen Group, 114. 
 
 Nitrogen, 115. Alinospliere, n6. luKlioinetry, 120. Measuring Gases, 121. 
 Diffusion of (jases, 123. Gases Recently Discovered in the Atmosphere, 123. 
 Argon, 124. 1 leliuin, 125. Goinpouiufs of Nitrogen with Hydrogen, 125. Atn- 
 
 monia, 126, Amtnoniuin Salts, 129. Atoinie Weight of Nitrogen, 129. Ilydrox- 
 ylainine, 130, Dianiide, 131. Ilydrazoie Aeid, 132. Compounds of Nitrogen 
 with tlie Halogens, 133. IMiosphorus, 135. Compounds of J'hosphorus with 
 Hydrogen, 138. I'hosphoniuin Salts, 140. Comi)ounds of IMiosphorus with the 
 Halogens, 141. Arsenie, 143. Arsine, 144. Compounds of Arsenie with the 
 Halogens, 145. Antimony, 146, Stihine, 147. Comj)ounds of Antimony with 
 the Halogens, 147. Tabulation of the Elements of the Nitrogen Grouji, 148. 
 
 Carbon Group, 149. 
 
 Carbon, 149. Carbon Compounds of Hydrogen, 15 1. Methane, 152, Atomic 
 Weight of Carbon, 152. Ethane; Ethylene, 153. Acetylene, 154. Nature of 
 Elame, 155. Comixninds of Carbon with the Halogens, 159. Silicon, iCo. 
 Hydrogen Silicide, 161. Compounds of Silicon with the Halogens, 161, 162. 
 Hydrogen Silico-fluoride, 163. Silicon Carbide (Carborundum), 163. 
 
 Valence of the Elements; Chemical Structure of the Molecules, 165. 
 
 Oxygen Compounds of the Metalloids, 172. 
 
 Oxygen Compounds of the Halogens, 173. 
 
 Oxygen Compounds of Chlorine, 172. Ilypochlorous Oxide, 174. Hypochlorous 
 Acid, 174. Chlorine Trioxide ; Chlorous Acid ; Chlorine Tetroxide, 176. Chloric 
 Acid, 177. rerchloric Acid, 178. Oxygen Compounds of Eromine, 179. Oxygen 
 Compounds of Iodine, 180. Hydrate* of the Acids, 181. 
 
 Oxygen Compounds of the Elements of the Sulphur Group, 182. 
 
 Oxygen Compounds of Sulphur, 183. Sulphur Dioxide, 183. Sulphurous Acid, 
 185. Hydrosulphurous Acid, 186. Sulphur Sesquioxide, 187. Sulphur Trioxide, 
 
 187. Thermo-chemical De])ortment, 187. .Sulphur Heptoxide ; Persulphuric Acid, 
 
 188. Sulphuric Acid, 189 Disulphuric Acid, 193. Sulphuric Acid Chloranhydrides, 
 195. Chlorsulphonic Acid; Sulphuryl Chloride, 195. Amido-derivatives, 196. 
 Thiosulphuric Acid ; Polythionic Acids, 197. Oxygen Derivatives of Selenium and 
 Tellurium, 199. 
 
 Oxygen Compounds of the Elements of the Nitrogen Group, 200. 
 
 Oxygen Derivatives of Nitrogen, 201. Nitric Acid; Nitrogen Pentoxide ; Nitryl 
 Chloride; Nitrosyl Chloride; Nitramide, 204. Nitrogen Trioxide ; Nitrous Acid, 
 205. Nitrogen Tetroxide, 206. Nitrosyl-sulphuric Acid, 207. Nitric Oxide, 209. 
 Nitrous Oxide, 21 1. Hyponitrous Acid, 212. Compounds of Nitrogen and Sul- 
 phur, 213. Oxygen Compounds of Phosphorus, 213. Hypophosphorous Acid, 214. 
 Phosphorous Acid, 215. Phosphoric Acid ; Pyrophosphoric Acid ; Hypophosphoric 
 Acid, 216. Metaphosphoric Acid ; Phosphorus Pentoxide, 217. Chloranhydrides 
 of the Acids of Phosphorus, 218. Phosphorus Compounds with Sulphur, 219. 
 Oxygen Derivatives of Arsenic, 220. Arsenic Trioxide, 220. Arsenic Acid, 221. 
 Compounds of Arsenic with Sulphur, 222. Sulpho-salts, 223. Oxygen Derivatives 
 of Antimony, 223. Antimony Oxide ; Antimonic Acid, 224. Antimony Sulphides, 
 225. Vanadium ; Niobium ; Tantalum, 226. 
 
 Oxygen Compounds of the Elements of the Carbon Group, 226. 
 
 Oxygen Compounds of Carbon, 227. Carbon Dioxide, 227. Critical Pressure, 228. 
 Physiological Importance of Carbon Dioxide ; Percarbonic Acid, 231. Carbon Mon- 
 oxide, 231. Nickel Carbonyl, 233. Carbonyl Chloride, 234. Amido-derivatives of 
 Carbonic Acid, 234. Compounds of Carbon with Sulphur, 234. Cyanogen Com- 
 pounds, 235. 'rhermo-chemistry of the Carbon Compounds, 236. Oxygen Com- 
 pounds of Silicon, 236. Dialysis, 237. Crystalloids and Colloids, 238. Silicates, 
 238. M'itanium ; Zirconium; M'horium, 238. lk)ron, 24I. Boron Hydride, 241. 
 Honm (Chloride, 241. Boron Eluoride, 242. Boric Acid, 242. 
 
 The Periodic System of the Elements, 243. 
 
 I’criodicily of Chemical Valence, 2|8. Correction of Atomic Weights, 250. 
 
TABLE OF CONTENTS. 
 
 XI 
 
 THE METALS. 
 
 Physical Properties of the Metals, 251. Atomic Volumes, 252. Light and Heavy 
 Metals, 252. Melting Points of the Metals, 252. Electric Furnace, 253. Specific 
 Heat; Atomic Heat, 253. Thermal Atomic Weights, 254. Isomorphism, 255. 
 Chemical Properties of the Metals, 255. Alloys, 255. Amalgams, 257. Metallic 
 Carbides, 257. Halogen Compounds, 257. Oxides and Hydroxides, 258. Per- 
 oxides, 259. Salts, 259. Action of Metals on Salts and Acids, 261. Electrolysis 
 of Salts, 262. Faraday’s Law, 263. Solutions, 265. Theory of Dilute Solutions, 
 267. Theory of Electrolytic Dissociation, 268. Transposition of Salts, 270. 
 
 Group of the Alkali Metals, 272. 
 
 d'hermo-chemistry of the Alkali Metals, 273. Potassium, 273. Potassium Hydride, 
 274. Dissociation, 274. Potassium Oxides, 275. Potassium Hydroxide, 275. 
 Potassium Chloride, 276. Bromide, 276. Iodide, 276. Fluoride, 277. Cyanide, 
 277. Potassium Chlorate, 277. Potassium Hypochlorite, 278. Potassium Sulphate, 
 279. Potassium Sulphite, 279. Potassium Persulphate, 279. Potassium Nitrate, 
 279. Gunpowder, 280. Potassium Nitrite, 280. Potassium Phosphate, 280. Potas- 
 sium Carbonate, 280. Potassium Silicate, 282. Potassium .Sulphides, 283. Potas- 
 sium Amide, 283. Recognition of the Potassium Compounds, 283. Rubidium ; 
 Caesium, 284. Sodium, 284. Sodium Oxides, 285. Sodium Hydroxide, 286. 
 Sodium Chloride, 286. Sodium Bromide, 287. Sodium Iodide, 287. Sodium 
 Chlorate, 287. Sodium lodate, 287. Sodium Sulphate, 287. Supersaturated 
 Solutions, 287. Sodium Hyposulphite, 289. Sodium Carbonate, 289. Sodium 
 Nitrate ; Sodium Phosphates, 293. Borax, 294. Sodium Silicate, 295. Sodium 
 Nitride, 295. Recognition of Sodium Compounds, 295. Lithium, 295. Ammo- 
 nium Compounds, 296. Ammonium Chloride, 297. Ammonium Carbonate, 298. 
 Ammonium Phosphates, 298, Ammonium Nitride, 298. Ammonium Sulphide, 
 299. Ammonium Hydrosulphide, 299. Recognition of Ammonium Compounds, 
 299 - 
 
 METALS OF GROUP II, 299. 
 
 Group of the Alkaline Earths, 3C0. 
 
 Calcium, 300. Calcium Oxide, 301. Cement, 302. Calcium Chloride, 302. Cal- 
 cium Fluoride, 302. Chloride of Lime, 303. Calcium Sulphate, 304. Calcium 
 Phosphates, 304. Calcium Carbonate, 305. Glass, 306. Calcium Sulphides, 307. 
 Calcium Carbide, 307. Strontium, 307. Barium, 308. Barium Oxide, 308. Barium 
 Peroxide, 309. Barium Sulphate, 309. Barium Persulphate, 309. Barium Car- 
 bonate, 310. Recognition of the Compounds of the Alkaline Earths, 310. Diam- 
 monium Compounds, 310. Hydrazine Hydrate, 31 1. Diammonium Chloride, 312. 
 Diammonium Nitride, 312. Azides, 312. 
 
 Magnesium Group, 312. 
 
 Magnesium, 313. Magnesia, 314. Magnesium Chloride, 314. Magnesium Sul- 
 phate, 315. Magnesium Phosphates, 315. Magnesium Carbonate, 316. Mag- 
 nesium Nitride, 316. Recognition of Magnesium Compounds, 317. Beryllium, 
 317. Zinc, 318. Zinc Oxide, 318. Zinc Chloride, 319. Zinc Sulphate, 319. 
 Zinc Sulphide, 319. Cadmium, 320. Comparison of Zinc, Cadmium, and Mercury, 
 321. Mercury, 322. Amalgams, 323. Mercurous Compounds, 323. Mercuric 
 Compounds, 325. Heat of Formation of the Chlorides of Group II, 327. 
 
 Copper, Silver, Gold, 328. General Characteristics, 329. Forms of Combination, 329. 
 Copper, 330. Metallurgy of Copper, 331. Cuprous Compounds, 332. Cupric Com- 
 pounds, 334. Alloys of Copj)er, 335. Silver, 336. Metallurgy, 337. Silver Oxides, 
 338. Molecular Formulas, 339. Silver Chloride, 339. Photography, 340. Nitrate 
 of Silver, 341. Silver Nitride, 341, Silver Sulphide, 341. Silvering, 342. Gold, 
 342. Aurous Compounds, 343. Auric Compounds, 344. 
 
Xll 
 
 TABLK OF CONl'ENTS. 
 
 METALS OF GROUP III, 345. 
 
 Group of Earth Metals, 346. 
 
 Aluminium, 346, Aluminium Chloride, 348. Aluminium Oxide, 350. Aluminates, 
 350. Aluminium Sulphate, 352. Alum, 352. Aluminium Silicates, 353. Porcelain, 
 353. Ultramarine, 354. Rare Earth Metals, 354. 
 
 Gallium Group, 357. 
 
 Gallium, 358. Indium, 358. Thallium, 359. 'I'hallous Compounds, 360, Thallic 
 Compounds, 360. 
 
 METALS OF GROUP IV, 361. 
 
 Germanium, 362. Germanous Compounds, 362. Germanic Compounds, 363. 
 
 Tin, 363. Stannous Compounds, 364. Stannic Comi)ounds, 366. Stannates, 367. 
 Sulpho-stannates, 367. Lead, 367. Lead Oxide, 369. Plumbic Acid, 370. Lead 
 Chloride, 370. Lead Nitrate, 371. Lead Carbonate, 371. Lead Sulphide, 371. 
 Bismuth, 372. Pismuthic Acid, 373. Pismuth Nitrate, 373. 
 
 Chromium Group, 373. 
 
 Chromium, 374. Chromous Compounds, 375. Chromic Compounds, 375. Chro- 
 mium Alum, 377. Chromates, 377. Potassium Chromate, 379. Chromyl Chloride, 
 380. Molybdenum, 382. Tungsten, 384. Uranium, 384. 
 
 Manganese, 386. 
 
 Forms of Combination, 386. Manganous Compounds, 387. Manganic Compounds, 
 388. Manganese Peroxide, 389. The Acids of Manganese, 390. 
 
 METALS OF GROUP VIII, 392. 
 
 Iron Group, 392. 
 
 Iron, 393. Cast Iron, 394. Wrought Iron, 394. Metallurgy of Iron, 395. Fer- 
 rous Compounds, 397. Ferric Compounds, 399. Ferric Acid Compounds, 401. 
 Cyanogen Compounds, 401. Metallic Ions, 401. Nickel, 404. Cobalt, 406. 
 Platinum Metals, 408. 
 
 Ruthenium and Osmium, 410. Rhodium and Iridium, 41 1. Palladium, 412, 
 Platinum, 413. 
 
 Spectrum Analysis, 415. 
 
 INDEX, 421. 
 
A TEXT-BOOK 
 
 OF 
 
 INORGANIC CHEMISTRY. 
 
 INTRODUCTION. 
 
 The natural sciences are occupied with the investigation of the innu- 
 merable substances and changes by which we are surrounded. Physics 
 and chemistry differ in a sense from the other sciences, in that their domain 
 of research is not restricted to any very definite province of nature; 
 indeed, it is not confined to any one planet. This is due to the fact that 
 all changes, so far as they are perceptible to our senses, are referable to 
 chemical and physical causes. The best and mist satisfactory informa- 
 tion relative to the ]3roperties and composition of a substance, no matter 
 what its source, is afforded by jihysics and chemistry. For these reasons, 
 therefore, these two sciences — chemistry and physics — are regarded as the 
 general natural sciences in contradistinction to the other more special 
 sciences. 
 
 Physics deals with the doctrine of equilibrium and with that of 
 motions. The latter are visible, as those of mass — in fall, projection, 
 rotation, propagation in a plane, etc. ; or they are invisible, and are only 
 perceptible by their results — sound, heat, light, electricity. Chemistry, 
 on the other hand, reveals to us the composition of matter, and, in the 
 formation of new compounds, acquaints us with the rules and laws by 
 which its various forms act upon each other. The domain of chemistry 
 and physics consequently extends throughout all the natural kingdoms, 
 and each of the special natural sciences, even astronomy, avails itself 
 of the aid given by physics and chemistry to attain its own particular 
 goal. However, the.se two sciences are mutually dependent upon one 
 another for, so far as we know, there cannot be motion without matter, 
 nor matter without motion. 
 
 The influence exerted by chemi.stry and physic.s upon civilization corresponds to 
 their exalted position in the group of natural sciences. By means of these sister 
 sciences the y^roducts and forces of nature have been more completely utilized than 
 ever before, hence there are but few fields of human activity which in the course of the 
 pre.sent century have not been enriched by the accumulation of chemical and physical 
 observations. If the conquests of chemical and physical investigation of even the last 
 2 17 
 
cS 
 
 INORCIANIC CHEMISTRY. 
 
 five years were suddenly to disappear an almost unbearable retrogression would be fell in 
 commerce, manufacture, the various industries and in agriculture, and every individual 
 would find himself having recourse to innumerable advantages and facilities of wliic li at 
 present he is scarcely conscious. 
 
 The following figures give some idea of the importance of ( ierman chemical industries 
 In 1895, 114,581 operators were engaged in 5974 factories devoted to chemical manu 
 facture ; these represented in round numbers 100,000,000 M. or ^25,000,000 in wages, 
 etc. The exports of chemical i)roducts for the same year were valued at 290,000,000 M. 
 or $72,500,000. 
 
 Compare : Wickelhaus, Wirthschaftliche Hedeutung chemischcr Arbeit, lierlin, 1893 ; 
 Ferd, Fischer, Das Studium der technischen Chemie, Jlraunschweig, 1897. 
 
 A closer scrutiny of natural olijects discloses the fact that they in 
 time succumb to many more or less serious alterations or changes. 
 Although no abrupt boundaries are [tresented in nature, but gradual 
 transitions and intermediate steps throughout, two tolerably distinct 
 classes of phenomena may be observed. Some changes in the condition of 
 bodies are only superficial (external), and are not accompanied by material 
 alteration in sid'stance. Thus heat converts water into steam, which upon 
 subsequent cooling is again condensed to water, and at lower temperatures 
 becomes ice. In these three conditions, the solid, li(|uid, and gaseous, 
 the substance or the fuatter of water or ice is unchanged; only the sejja- 
 ration and the motion of the smallest particles — their states of aggregation 
 — are different. If we rub a glass rod with a piece of cloth, the glass 
 acquires the property of attracting light objects, e. g., particles of paper. 
 It becomes electrified. An iron rod allowed to remain suspended verti- 
 cally for some time slowly acquires the power of attracting small pieces of 
 iron. Through the earth’s magnetism it has become magnetic. In both 
 instances the glass and iron receive new properties; in all other respects, 
 in their external and internal form or condition, they have suffered no 
 perceptible alteration; the glass is glass, and the iron remains iron. All 
 such changes in the condition of bodies, unaccompanied by any real alter- 
 ation in substance, are known as physical phenomena. 
 
 Let us now turn our attention to the consideration of another class of 
 phenomena. It is well known that ordinary iron undergoes a change, which 
 we term rusting, i. e., it is transformed into a brown substance which is en- 
 tirely different from iron. On mixing finely divided copper filings with 
 flowers of suli)hur (pulverulent sulphur) there results an apparently uniform, 
 grayish-green powder. If this be examined, however, under a magni- 
 fying glass, we can very plainly distinguish the red metallic copper par- 
 ticles in it from the yellow of the sulphur; by treating with water, the 
 sjiecifically lighter sulphur particles can easily be separated from those of 
 tlie cop])er. Carbon bisulphide will dissolve out the sulphur particles. 
 Hence this powder represents nothing more than 2l mechanical mixture. 
 If, however, this mixture be heated, e. g., in a glass test-tube, it will com- 
 mence to glow, and on cooling, a black, fused mass remains, which differs 
 in all res])ects from coj)per and sulphur, and even under the strongest 
 microscojie does not reveal the slightest trace of the latter, and elutri- 
 ation witli water or treating with carbon bisuli)hide will not affect a sepa- 
 ration of the ingredients. By the mutual action of sulphur and copper in 
 |)resence of heat, a new body with entirely different ])roi)erties has been 
 
INTRODUCTION. 
 
 19 
 
 produced, and is named copper sulphide. Mixtures of sulphur with iron 
 or with other metals act in a similar manner ; the resulting bodies are 
 known as sulphides. 
 
 Such mutual action of different bodies occurs not only under the influ- 
 ence of heat, but frequently at ordinary temperatures. If, e.g., mercury 
 and sulphur are rubbed together in a mortar, there is produced a uniform, 
 black compound, called mercury sulphide. The action of gaseous chlorine 
 upon various metals is quite energetic. When finely divided antimony is 
 shaken into a flask filled with yellow chlorine gas, flame is produced; 
 each antimony particle burns in the chlorine with a bright white light. 
 The product of this action of solid metallic antimony and gaseous yellow 
 chlorine is a colorless, oily liquid, known as antimony chloride. Such 
 occurrences, therefore, in which a complete and entire alteration takes 
 ])lace in the bodies entering the reaction, are termed chemical phenomena. 
 
 In the previously described experiments we observed the phenomena 
 of chemical combination ; from two different bodies arose new homo- 
 geneous ones. The opposite may occur : the deco77iposition of com- 
 pound bodies into two or more dissi77iilar ones. If red mercuric 
 oxide be heated in a test-tube, it will disappear; a gas (oxygen) is 
 liberated, which will inflame a mere spark on wood ; in addition, we find 
 deposited upon the upper, cooler portions of the tube, globules of 
 mercury. From this we observe that on heating solid red mercuric oxide 
 two different bodies arise: gaseous oxygen and liquid mercury. We 
 conclude, then, that mercuric oxide holds in itself, or consists of, two 
 constituents — oxygen and mercury. This conclusion, arrived at by 
 decomposition, or analysis., may be readily verified by combination or 
 synthesis. It is only necessary to heat mercury for some time, at a some- 
 what lower temperature than in the preceding experiment, in an atmos- 
 ])here of oxygen, to have it absorb the latter and yield the compound we 
 first used — red mercuric oxide. The direct decomposition of a compound 
 body into its constituents by mere heat does not often happen. Generally, 
 the cooperation of a second substance is required, which will combine 
 with one of the constituents and set the other free. In this manner we 
 can, for example, effect the decomposition of the previously synthesized 
 mercury sulphide, viz., by heating it with iron filings; the iron unites 
 with the sulphur of the mercury sulphide, to form iron sulphide, while the 
 mercury is set free. 
 
 If, in a similar manner, natural objects be decomposed, bodies or sub- 
 stances are finally reached which have withstood all attempts to bring 
 about their division into further constituents, and which cannot be formed 
 by the union of others. Such substances are che7nical ele77ienis ; they 
 cannot be converted into one another, but constitute, as it were, the limit 
 of chemical change. Their number, at present, is about 70 ; some have 
 been only recently discovered. To them belong all the metals, of which 
 iron, copper, lead, silver, and gold are examples. Other elements do not 
 possess a metallic appearance, and are known as 7netalloids (from I 
 resemble). It would be more correct to term them non-7netals. To these 
 belong sulphur, carbon, phosphorus, oxygen, etc. The line between 
 metals and non-metals is not very marked. Thus, mercury, despite the 
 
20 
 
 INORGANIC CHEMIS’IRY. 
 
 fad that it is liquid at tlie ordinary tcini)craturc, must 1)C included among 
 the metals because of its chemical i)roj)ertics. 
 
 All the substances known to us arc made up of these elements. Water 
 is acomj)onnd of two gaseous elements — hydrogen and oxygen ; common 
 salt consists of the metal sodium and the gas chlorine, 'bhe elements 
 make up not only our own earth, but the heavenly bodies are comijosed 
 of them ; at least as far as has been proved by spectrum analysis. 
 
 THE PRINCIPLE OF THE INDESTRUCTIBILITY OF MATTER. 
 
 If the total weight of substances, wliich are to act chemically upon one 
 another, be determined, and the chemical action be then allowed to 
 occur, it will be discovered upon ascertaining the weight of the resulting 
 bodies, if due consideration be given for unavoidable errors of ex})eri- 
 ment, that no loss or increase in weight has occurred— no change in 
 mass, because mass and weight are strictly i)roj)ortional to each other for 
 one and the same j)lace. Jt is in these cases immaterial whether a com- 
 pound body be resolved into its elements, or whether elements unite to 
 produce com])ound bodies; the })roducts ])resent after the chemical 
 reaction will always weigh exactly what the bodies ])receding the reaction 
 weighed (comi:)are the ex])eriments of Landolt, Ber. 26 (1893), 1820). 
 Well-known, general phenomena apparently contradict this scientific 
 conclusion. We observe plants springing from a small germ and con- 
 stantly acquiring weight and volume. This spontaneous increase of sub- 
 stance, however, is only seeming. Closer inspection proves conclusixely 
 that the growth of plants occurs only in consequence of the alxsorjjtion 
 of substance from the earth and atmosphere. The opi)osite phenomenon 
 is seen in the burning of combustible substances, where an apparent 
 annihilation of matter takes place. But even in this, careful observation 
 will discover that the combustion phenomena consist purely in a trans 
 formation of visible solid or liquid bodies into non-visible gases. Carbon 
 and hydrogen, the usual constituents of combustible substances, e. g., 
 petroleum or wood, combine in their combustion with the oxygen of 
 the air and yield gaseous products — the so-called carbon dioxide and 
 water — which diffuse into the atmosphere. If these products be col- 
 lected, their weight will be found not less, but indeed greater, than that 
 of the consumed body, and this is explained by the fact that in addition 
 to the original weight they have had the oxygen of the air added. Such 
 a combustion must, therefore, be regarded as a conversion of visible 
 solid or liquid substances into invisible gaseous matter. 
 
 d'he production (creation) or annihilation of matter has never been 
 demonstrated as occurring in any change. A conq^ound body is com- 
 prised of certain elements, and contains a very definite quantity by weight 
 of each of them. If it decomixise it naturally breaks down into its con- 
 stituents, which iierhaps reunite in some other manner to form new com- 
 pounds, alwavs, however, pres'^rving their original nature, their original 
 weight and their masses. This fundnmcntal truth is ihe law of the inde- 
 stnictihilify of matter. 'The early (Irecian ])hilosophers arrived at this 
 conclusion by a keen observation of the changes of daily life (compare 
 
INTRODUCTION. 
 
 21 
 
 Lucretius — ‘‘Nature”), and since that time it has always been viewed 
 as an established fundamental principle of exact, scientific investigation. 
 Consult Debus, Ueber einige Fundamental-Satze der Chemie, Kassel, 1894. 
 
 THE PRINCIPLE OF THE CONSERVATION OF ENERGY. 
 
 All imaginable, appreciable, natural phenomena have their causes, which 
 escape simple observation, and are only realized by scientific research. 
 We observe that iron rusts on exposure to the air. By chemical investi- 
 gation we learn that rust is a compound of iron and oxygen. This latter, 
 as we also learn from chemistry, is a constituent of the air. If it be re- 
 moved from the latter, the iron will cease to rust. Hence the tendency 
 of iron and oxygen to combine with one another is designated as the cause 
 of the rusting. It is generally said that a force exists between them, 
 which effects their union and this force is called chemical affinity. In 
 changes of other kinds our ex[)lanation assumes the presence of some 
 force. The falling of bodies is attributed to the force of gravitation — 
 a universal attracting force, which even influences the course of the stars. 
 The decomposition of a body by external efforts acts in opposition to its 
 cohesive force. Two different bodies directly in contact with one another 
 are held together by the adhesive force existing between them. By assump- 
 tions similar to these it is possible to refer an almost infinite number of 
 changes to but comparatively few causes. But, in doing this, it must be 
 remembered that the nature of the force continues enigmatical, even if 
 we know the laws, by which it acts, as well as we know those of gravity. 
 
 At present the movements of the ])arts of matter are considered to be 
 the cause of other phenomena which formerly were ascribed to the actions 
 of special forces. It has long been known that a sounding body com- 
 pletes certain vibrations, which are imparted to the surrounding air and 
 arrive thereby to the tympana of our ears. The same is true of the phe- 
 nomena of light and heat. It is supposed that the heat phenomena are 
 due to an energetic oscillating motion of the smallest particles of a body. 
 Upon grasping a warm substance, motion present in it is partially trans- 
 ferred to us and we experience warmth. If, on the contrary, heat motion 
 be passed from the hand to the object touched, the latter appears to us to be 
 cold. At sufficiently elevated temperatures, and indeed even under other 
 conditions (phosi)horescence, fluorescence) the heat motions of bodies 
 produce in the surrounding aether (a hypothetical medium, enigmatical 
 in its nature and capable of penetrating everything) motion extending 
 in all directions in a wave-like manner and known as radiant heat. If 
 these waves follow each other rapidly enough they become light waves 
 and are perceived as such by the retina of the eye. The source of electric 
 phenomena and of the x or Rontgen rays is also supposed to be due to 
 motion, of an unexplained nature, present in the aether. H. Hertz in 
 1888 demonstrated by the radiation of a periodic electric force reversing 
 its direction that these rays were subject to the same laws as those of light 
 and radiant heat. They are not only propagated in direct lines with all 
 the velocity of light but are similarlv reflected and refracted. Conse- 
 quently, in the periodic change of direction we find the same motion 
 
22 
 
 INORGANIC CHEMISTRY. 
 
 alterations of the aether to be at the basis of liglU and radiant heat, as 
 they are the cause of electric phenomena. 
 
 Numerous physical investigations have demonstrated that the various 
 forms of motion can be transferred mjt only from one body to another 
 but that they can be converted into one another. The bullet which in its 
 flight is sustained by resistance becomes hot ; the visible motion of the 
 entire mass ceases and is then converted into the invisible motion of the 
 minutest particles perceptible to us as heat. Conversely, the motion of 
 the smallest particles is changed to that of large masses, if by means 
 of the steam engine we ])roduce driving force by heat, and the latter by 
 combustion, a chemical change. 
 
 It will be discovered upon attempting to measure these changes that 
 the different forces or modes of motion bear a fixed ratio of trans- 
 formation to one another. According to the proposition, causa cequat 
 cffectum, they are subject to equivalent convertibility. At present the 
 indestructible portion of all these changes is termed energy, which mani- 
 fests itself as (i) jnechanical energy, (2) thermal energy, (3) electric and 
 jna;netic energy, (4) chetnical and internal e 7 iergy and (5) radiant energy. 
 
 The power of a body to do work is called energy. It is distinguished as potential and 
 kinetic. Examples are as follows : The mass ?n, acted upon by the constant force p, 
 traverses the distance J and acquires the speed v. This would be represented by the 
 
 equation ps = As the force p has acted through the distance s, it has performed 
 
 the work ps. If a hammer, weighing p, be raised to the height s, the work done, 
 opposed to gravity, equals /s-. This work is present as potential energy in the raised 
 hamirrer, so far as the presence of a force and its removal through space can possibly make 
 a mechanical action present. If the hammer be allowed to fall it acquires kinetic energy 
 
 -^^>2 (active force), which is equal to the work expended in raising the hammer, or the 
 
 potential energy of the raised hammer. In the fall the potential is transformed into 
 kinetic energy [Kiveu — to move ; kvepyeu — to act). 
 
 A given mass-motion can be converted into a definite heat-quantity. 
 By the application of the latter, work can be performed again equivalent 
 to the mass-motion (first principle of the mechanical theory of heat). 
 The quantity of heat sufficient to raise i kilogram of water at the ordinary 
 temperature 1° C. is taken as the unit in the determination of heat and 
 is called calorie (large). To produce this unit by mass-motion would 
 require mechanical work (at the average geographical latitude) equal to 
 426 kilogram-meters. This quantity signifies that if under the condi- 
 tions mentioned as to plate i kilogram falls through 426 meters, or 426 
 kilograms through i meter, and the resulting total kinetic energy be 
 transformed into heat, the latter will raise i kilogram of water from 15° 
 to 16° in temperature. In the conversion of heat into mass-motion, 
 however, one calorie will disappear for every 426 kilogram-meters of 
 work performed. This magnitude (or constant) is known as the tne- 
 chanical equivalent of heat. 
 
 Chemical energy can be measured by tlie heat or electricity developed in chemical 
 changes. While all forms of energy can be converted without difficulty into heat, this 
 particular form is only altered with limitations into the other forms. It is only when 
 lieat passes from a warmer to a colder substance that a tlelinite portion of it can by proper 
 
INTRODUCTION. 
 
 23 
 
 appliances be changed into mechanical work. Could this so occur that the greatest possible 
 portion of the heat could be utilized in the performance of mechanical work, then the 
 process might be reversed. If, however, the greatest possible quantity of heat is not con- 
 sumed in doing such work but is lost by conduction and radiation from a higher to lower 
 temperature, then the change is not reversible. For a portion, at least, of the heat there 
 is no possibility of change to mechanical work ; it is permanently lost {degradation of 
 energy). The reversibility is made impossible by such changes (conduction, radiation, fric- 
 tion, etc.), which take place under given conditions, of their own accord. As the op- 
 posite conversions do not occur spontaneously (the passage of heat from a body of 
 lower to one of higher temperature), in nature the changes of the first order must exceed 
 those of the second order. The sum of all the conversion-values (calculated as proceeding 
 positively in the sense of the first group) is termed entropy {rpoirri — conversion). The 
 entropy of all nature is therefore in the act of constant increase (Clausius, Second Principle 
 of the Mechanical Theory of Heat). 
 
 The law of the conservation of energy, according to which the energy 
 in nature is of a convertible character but unalterable in quantity, con- 
 stitutes one of the most important foundations of the science of nature. 
 
 It was first clearly explained and definitely enunciated by Julius Robert Mayer (1842), 
 a physician of Heilbronn (Annalen, 42, 233 ; see also his “ Mechanik der Warme gesam- 
 melte Schriften ”), Not knowing of Mayer’s work, or of the treatise of Colding (1843), 
 a Dane, in which the principle of energy was also developed, Hermann v. Helmholtz 
 (1847) announced the law as empirical, followed and developed it mathematically 
 through all the domains of natural phenomena (Ueber die Erhaltung der Kraft, Berlin, 
 1847). Mayer was also the first to discover the mechanical equivalent of heat, which 
 shortly afterwards James Prescott Joule definitely determined by accurate experiments. 
 
 Heat almost invariably appears in chemical union ; even light and 
 electricity can be produced by chemical processes, or work can be per- 
 formed in opposition to external pressure by increase in volume. All 
 these forms of energy owe their origin to the potential energy of the 
 chemical forces, which in the process of chemical change do work. 
 Hence, we may speak of chemical energy or of chemical tension. In the 
 chemical decomposition of a compound body into its components, on the 
 other hand, heat is usually absorbed and disappears as such and becomes 
 chemical energy. Thus, in the union of approximately i kilogram of 
 hydrogen with 8 kilograms of oxygen to produce 9 kilograms of water 
 a quantity of heat, equivalent to 34,200 calories, is set free, and this 
 corresponds to work equivalent to 34,200 X 426 = 14,569,200 kilogram- 
 meters. In the decomposition, on the other hand, of 9 kilograms of water 
 into hydrogen and oxygen, the same force or quantity of heat is neces- 
 sary. Therefore, the same quantity of force or motion must be contained 
 in the form of chemical energy in the liberated hydrogen and oxygen. 
 
 Although all bodies, and the elements especially, possess chemical 
 energy, they do not manifest it in the same way. Some of them react 
 readily with one another and others with difficulty or not at all. The 
 cause of this variation in behavior is entirely unknown to us. It is cus- 
 tomary to express the fact by saying the bodies have a strong, a feeble, or 
 no affinity for one another. Formerly, bodies which combined chem- 
 ically, were supposed to be related to each other, and it was assumed that 
 their affinity — their tendency toward one another — was satisfied by their 
 union. This choice of terms was unfortunate. The nature of the chemical 
 attraction, which produces and holds chemical compounds together is just 
 
24 
 
 INORGANIC CHEMISTRY. 
 
 as enigmatical as that of gravity. Even the laws in accordance with 
 which affinity acts are scarcely known to us. It is only in recent times 
 that advances have been made in this direction, since the positions and 
 the motions of atoms and of molecules have been regarded as affording 
 hints as to the cause of a reaction. 
 
 CONSTITUTION OF MATTER. ^ 
 
 >- 
 
 ATOM AND MOLECULE. 
 
 If we seek to give expression to the constitution of the chemical ele- 
 ments and the bodies composed of them we return, if guided by experience, 
 to the ancient atomic hypothesis which alone is justified by the present 
 condition of chemical and ])hysical investigation. It apjiears that the 
 Indian and Grecian natural philosophers established this same hypothesis 
 in a purely inductive manner [Kanada (founder of the Vaiseshika system), 
 Lucippus, Democritus (500 b. c.) and Ei)icurus (400 b. c.)]. 
 
 The foundation of this hypothesis is brilliantly set forth in the first book of Lucre- 
 tius (died 55 B. c. ) “ Ueber die Natur der Dinge.” Numerous observations are re- 
 corded. The following occurs there in substance : All bodies can be divided into in- 
 finitely small parts, no longer recognizable by sight or taste. Invisible aqueous vapor 
 separates from the air as water upon cold objects and again disappears on the ap])roach of 
 heat. The ring that is constantly worn on the finger becomes thinner in the course of 
 years. Dropping water wears away the stone. The smooth pavement is made rough by 
 walking. All these things occur without our perceiving what at any one time departs 
 from the ring, etc. Hence we conclude the bodies are composed of invisible, extremely 
 small parts, which to us are without mass. These particles, the atoms (arofioQ — indivisible), 
 are indestructible and cannot be created. There is nothing beyond them and the vacant 
 space between. The difference in things is due to the difference in number, size, form, 
 and arrangement of the atoms. There is no qualitative difference in the atoms, they act 
 upon one another only by contact and pressure. Change is only a union and separation 
 of atoms ; nothing occurs by chance but everything by reason and with necessity. Thus 
 far Lucretius. 
 
 In the first half of the 17th century the atomic idea, which until then had been 
 driven into the background by the Aristotelian philosophy, was resuscitated by Daniel 
 Sennert, a German physician, and a P'rench ecclesiastic named Pierre Gassendi. They 
 adopted the Greek atomic doctrine ; and, from the point of view of the atomists, constitute 
 the connecting link between the past and the present. Since, however, our modern 
 atomic notions have developed step by step from the ideas of Sennert and Gassendi, their 
 beginnings go back to and have their origin in Lucippus and Democritus. Atoms were 
 introduced into chemistry by Robert Boyle, a contemporary and follower of Gassendi. 
 Boyle was the first chemist to devote his experiments to the noble purpose of investi- 
 gating nature. (Compare F. A. Lange, Geschichte des Materialismus, 3 Aufl. 1876 ; 
 Grie.sbach physik.-chem. Propadeutik (1895), and the fasciculus of Debus referred to 
 on page 2I.) 
 
 'The scientific foundations of our present atomic doctrine will be 
 described later. Its fundamental ideas alone will be given now. We 
 assume that an element consists of atoms perfectly similar to one another, 
 but differing from those of other elements. We must grant that there are 
 as many kinds of atoms as there are different elements. A compound body 
 like iron sulphide, according to this view, is produced by the combination 
 of sulj)hur atoms with iron atoms in a definite ratio. Those ])articles of 
 a compound, representing the limit of divisibility so far as similarity 
 
. INTRODUCTION. 
 
 25 
 
 goes, are called molecules (j?iolecula, diminutive of moles — the mass), by 
 further division they are resolved into dissimilar parts. Hence, iron 
 sulphide is made up of molecules, which in turn consist of atoms of iron 
 and of sulphur. We shall learn later that the elements — with few excep- 
 tions— are composed, at ordinary temperature, not of a collection of free 
 atoms, but of an aggregation of atom-groups — of molecules. A molecular 
 structure is the rule. In the case of the elements the molecules are com- 
 posed of like kinds of atoms, in compounds they consist of dissimilar 
 atoms. 
 
 The greatest advancement of the atomic theory is due to John Dalton. 
 By assuming that the atoms combined with one another in definite pro- 
 portions, he laid the basis for the determination of the relative atomic 
 weights, and thereby became the founder of the chemical atomic theory 
 based upon definite weight proportions (1804). If two elements form 
 but a single compound by union with one another, it may be assumed 
 with Dalton, as long as no other reason to the contrary exists, that their 
 molecules consist of an atom of each of the two elements. Should two 
 compounds of the elements A and B be known, then the molecule of the 
 one compound would consist of an atom each of A and of B, consequently 
 of two atoms, while the other compound might be composed of three 
 atoms (2 A -f B or 2 B -f A), etc. With these premises clearly enunci- 
 ated it is possible to determine the relative atomic weights of the elements. 
 
 One hundred parts of the previously mentioned iron sulphide consist 
 in round numbers of 63.6 parts of iron and 36.4 parts of sulphur. If, 
 however, in accordance with the assumptions of Dalton, there is in this 
 compound one atom of sulphur for one atom of iron, then the atomic 
 weights of iron and sulphur must be to one another as 63.6 : 36.4. The 
 ratios between the atomic weights of the elements may be determined in 
 this manner. If for any element a number be taken for its atomic 
 weight, it can readily be calculated in what ratio the atomic weights of 
 all the other elements stand to this arbitrarily chosen standard, and we 
 thus obtain the relative ato 77 iic weights. That element which combines 
 with the majority of the other elements to form compounds capable of 
 the most accurate analysis, is chosen as the standard of comparison. 
 Finally, the number or value assigned this standard element as its atomic 
 weight is a matter of consensus of opinion. 
 
 Later we shall become acquainted with physico-chemical methods of 
 testing the atomic numbers, derived in a chemical way, and especially for 
 establishing whether it is not a fraction or a multiple of the true, relative 
 atomic weight. 
 
 CHEMICAL SYMBOLS AND FORMULAS. 
 
 The chemical elements are simjfiy and conveniently represented by the 
 initials of their Latin or Greek names in accordance with the suggestion 
 of the great Swedish chemist, John Jacob Berzelius (1779-1848) to 
 whom we are also indebted for the first accurate atomic weight determina- 
 tions. Thus hydrogen is designated by the letter H, from the word 
 hydrogenium ; nitrogen by N, from nitrogenium. Elements having the 
 same initials are distinguished by adding a second letter; thus, Na indi- 
 3 
 
26 
 
 INORGANIC CHEMISTRY. 
 
 cates natrium (sodium), Ni — nickel, Hg — mercury (from hydrargyrum), 
 Pd — palladium, Pt — platinum, etc. 
 
 The following table contains the names of 71 known chemical ele- 
 ments, together with their symbols and their atomic weights, refer7'cd to 
 oxygen equal to 16. 
 
 It may be added : 
 
 That as a rule the atomic weights are given with only as many decimals as are accu- 
 rate to the last figure. In the case of bismuth, nickel and tin, indicated by *, this rule 
 is not adhered to. Also in the ca.se of hydrogen the more accurate value 1. 008 is given 
 as 1. 01 for ordinary use. 
 
 The elements whose names have ? attached are in doubt either as to their simplicity, 
 or as to entire units in their atomic weights. 
 
 Elements. 
 
 j 
 
 Symbol. 
 
 Atomic 
 
 Wekjiit. 
 
 Elements. 
 
 Symbol. 
 
 Atomic 
 Weigh r. 
 
 Aluminium, 
 
 A 1 
 
 27. 1 
 
 Neodymium (?), . . . 
 
 Nd 
 
 144 
 
 Antimony (Stibium), 
 
 Sb 
 
 120 
 
 Nickel, 
 
 Ni 
 
 58.7* 
 
 Argon (?), 
 
 A 
 
 40 
 
 Niobium, 
 
 Nb 
 
 94 
 
 Ai’senic, 
 
 As 
 
 75 
 
 Nitrogen, 
 
 N 
 
 14.04 
 
 Barium, 
 
 Ba 
 
 137-4 
 
 Osmium, 
 
 Os 
 
 191 
 
 Beryllium, 
 
 Be 
 
 9-1 
 
 Oxygen, 
 
 0 
 
 16.00 
 
 Bismuth, 
 
 Bi 
 
 208.5* 
 
 Palladium, 
 
 Phosphorus, 
 
 Pd 
 
 106 
 
 Boron, 
 
 B 
 
 II 
 
 P 
 
 31.0 
 
 Bromine, 
 
 Br 
 
 79.96 
 
 Platinum, 
 
 Pt 
 
 194.8 
 
 Cadmium, 
 
 Cd 
 
 1 12 
 
 Potassium (Kalium), 
 
 K 
 
 39-15 
 
 Caesium, 
 
 Cs 
 
 133 
 
 Praseodymium (?), . . 
 
 Pr 
 
 140 
 
 Calcium, 
 
 Ca 
 
 40 
 
 Rhodium, 
 
 Rubidium, 
 
 Rh 
 
 103.0 
 
 Carbon, 
 
 C 
 
 12.00 
 
 Rb 
 
 854 
 
 Cerium, 
 
 Ce 
 
 .140 
 
 Ruthenium, 
 
 Ru 
 
 101.7 
 
 Chlorine,' 
 
 Cl 
 
 35-45 
 
 Samarium (?), . . . . 
 
 Sm 
 
 150 
 
 Chromium, 
 
 Cr 
 
 52.1 
 
 Scandium, 
 
 
 44-1 
 
 Cobalt, 
 
 Co 
 
 59 
 
 Selenium, 
 
 
 79-1 
 
 Copper, 
 
 Cu 
 
 63.6 
 
 Silicon, 
 
 Si 
 
 28.4 
 
 Erbium (?), 
 
 Er 
 
 166 
 
 Silver (Argentum), . . 
 
 Ag 
 
 107.93 
 
 Fluorine, 
 
 FI 
 
 19 
 
 Sodium (Natrium), . . 
 
 Na 
 
 23-05 
 
 Gallium, 
 
 Ga 
 
 70 
 
 Strontium, 
 
 Sr 
 
 87.6 
 
 Germanium, 
 
 Ge 
 
 72 
 
 Sulphur, 
 
 S 
 
 32.06 
 
 Gold (Aurum), .... 
 
 All 
 
 197.2 
 
 Tantalum, 
 
 Ta 
 
 183 
 
 1 leliuin (?), 
 
 He 
 
 4 
 
 Tellurium, ...... 
 
 Te 
 
 127 
 
 Hydrogen, 
 
 H 
 
 1. 01 
 
 Thallium, 
 
 T 1 
 
 204. 1 
 
 Indium, 
 
 In 
 
 1 14 
 
 Thorium, 
 
 Th 
 
 232 
 
 Iodine, 
 
 I 
 
 126.85 
 
 'bin (Stannum), .... 
 
 Sn 1 
 
 118.5* 
 
 Iridium, 
 
 Ir 
 
 193 
 
 Titanium, 
 
 Tungsten (Wolfram), • 
 
 I'i 
 
 48. 1 
 
 Iron (Ferrum), .... 
 
 Fe 
 
 56.0 
 
 W 
 
 184 
 
 Lanthanum, 
 
 La 
 
 138 
 
 Uranium, 
 
 Ur 
 
 239-5 
 
 Lead (Plumbum), . . . 
 
 PI) 
 
 206.9 
 
 Vanadium, 
 
 Yd 
 
 51.2 
 
 I.ithium, 
 
 Li 
 
 7-03 
 
 Ytterbium, 
 
 Yb 
 
 173 
 
 Magnesium, 
 
 Mg 
 
 24.36 
 
 Yttrium, 
 
 Y 
 
 89 
 
 Manganese, 
 
 Mn 
 
 55-0 
 
 Zinc, 
 
 Zn 
 
 65.4 
 
 Mercury, 
 
 Molybdenum, .... 
 
 Hg 
 
 Mo 
 
 200. 3 
 96.0 
 
 Zirconium, 
 
 Zr 
 
 90.6 
 
 In nddilioi) to the elements mentioned in the table the following are believed to have 
 been observed in certain rare minerals: terbium {^niosandrium^ philippiitnt^ 7 m, 
 
INTRODUCTION. 
 
 27 
 
 gadolinium^ decipium^ holmhcm, thuliiun and dysprosium. They occur associated with 
 cerium, lanthanum, scandium, ytterbium and yttrium. They are so much alike chemic- 
 ally that it is difficult to separate them and obtain them pure ; they are probably mixtures 
 of unknown elements. 
 
 Consult the paragraphs under Air for an account of what seem to be elementary gases 
 (metargon, neon, krypton, xenon) which have been isolated from the atmosphere. For 
 the reasons leading to the adoption of oxygen as the element of comparison and estab- 
 lishing its atomic weight as 16, see p. 73. 
 
 Compounds produced by the union of the elements are represented by 
 placing their corresponding symbols together and designating these 
 chemical formulas. Common salt, a compound of sodium and chlorine, 
 is represented by the formula NaCl ; mercuric oxide, a compound of 
 mercury and oxygen, by HgO; iron sulphide by FeS ; hypochlorous 
 acid, a compound of hydrogen, chlorine and oxygen, by HCIO. 
 
 By these premises chemical formulas acquire a very precise and evident 
 importance. The formula NaCl represents the union of i atom of sodium 
 with I atom of chlorine, and indicates that in it 23.05 parts, by weight, 
 of sodium are combined with 35.45 parts by weight of chlorine to yield 
 58.50 parts of sodium chloride (common salt). If several atoms of an 
 element are present in a compound, this is denoted by numbers which are 
 attached to the symbol of the atom : 
 
 HCl H2O NHg CH^ 
 
 Hydrochloric acid. Water. Ammonia. Methane. 
 
 The formula of water (HgO) means that its molecule consists of 2 
 atoms of hydrogen (2.02 parts by w^eight) and i atom of oxygen (16 
 parts by weight). The formula of sulphuric acid (H2SOJ indicates it to 
 be a compound consisting of i atom of sulphur (32.06 parts), 4 atoms of 
 oxygen (4 X 16= 64 parts), and 2 atoms of hydrogen (2 X i-oi = 2.02 
 parts), from which the composition of the acid may be at once calculated 
 into per cent., or into any desired quantity by weight. 
 
 Atomic Composition. In Per Cent. 
 
 Hydrogen, Hg = 2.02 2.06 
 
 Sulphur, . . S = 32.06 32.69 
 
 Oxygen, .. O4 = 64.00 65.25 
 
 H2SO4 = 98.08 100.00 
 
 A chemical change is represented by arranging these symbols in the 
 form of an equation. The left side of the equation indicates the sub- 
 stances present before the reaction occurs, while the right side shows the 
 products. Thus the chemical equation : 
 
 HgS -f Fe = FeS -f Hg 
 
 means that mercury sulphide (232.3-6 \parts) and iron (56 parts) have 
 combined to form iron sulphide (88.06 parts) and mercury (200.3 parts). 
 The equation 
 
 zH + O = HgO 
 
 indicates that i molecule of water has been formed by the union of 2 
 atoms of hydrogen with i atom of oxygen. 
 
28 
 
 INORGANIC CHEMISTRY. 
 
 It is true these equations give no idea of the conditions surrounding 
 these transpositions, nor of the alterations in energy which accompany 
 tliLMii. 'I'hey represent tlie i)urely material side of the change. 'I'hey indi 
 cate tiie quantities by weight of the substances entering the reaction aiul 
 also of the products: the weight of the bodies entering the reaction is 
 equal to that of the resulting })roducts. Iwery chemical equation is there- 
 fore an ex])ression of the law of the indestructibility of matter (p. 20). 
 
 CONDITIONS OF CHEMICAL CHANGE. 
 
 THERMO-CHEMICAL PHENOMENA. 
 
 The first requisite for bodies to act chemically iq)on one another is 
 that they be brought into most intimate contact because chemical action 
 does not occur at great distances. In the case of solids this intimate 
 contact, so essential for comi)lete chemical transformation, cannot ordi- 
 narily be attained by mere mechanical mixing : the necessary condition 
 is best reached by liquefying the bodies, or at least one of them, by 
 fusion or solution in some solvent. Hence the old saying corpora non 
 agiifit nisi fluida. 
 
 In many instances, however, the chemical transposition does not take 
 place even with the most intimate contact. An external physical im- 
 pulse occasioned by light, by electricity, by change in pressure (Spring, 
 van’t Hoff), but more especially by the temperature, is required for its 
 occurrence. Thus, for example, hydrogen and oxygen at the ordinary 
 temperature are wholly indifferent to one another despite the fact that as 
 gases they may be mixed as completely as it is possible. It is only when 
 they are heated that they combine — slowly at 200°, but with violent ex- 
 plosion about 700° — to water. The same occurs upon passing the electric 
 spark through the mixture. A mixture of hydrogen and chlorine will 
 remain unchanged in the dark, while in diffused sunlight the gases will 
 slowly unite to hydrochloric acid, but in direct sunlight, upon the appli- 
 cation of heat, or by passing the electric spark they will unite at once with 
 great violence. A mixture of iron and mercury sulphide requires the aid 
 of heat to bring about the transposition to iron sulphide and mercury. 
 At the ordinary temperature they appear to exercise no visible chemical 
 action upon one another. By the application of external energy — heat, 
 light, electricity, etc. — the atomic structure of the molecules of hydro- 
 gen and oxygen, of chlorine and hydrogen, of mercury sulphide and 
 iron, etc. — is first loosened or disintegrated and then the chemical action 
 between the several com})onents takes place. However, the experiments 
 of Raoul Pictet [1892; compare Ber. 26 ( 1893), iv, i and A. Welter, Die 
 tiefen 'rem])eraturen, Crefeld 1895] show that chemic al reactions do not 
 occur at temjicratures below — 125°. Substances which at the ordinary 
 temperature react with the greatest readiness — 0. suljihuric acid and 
 barium cliloride, hydrochloric acid and silver nitrate, sodium and alcohol 
 — appear at — 80° to be as indifferent toward one another as mercury sul- 
 phide and iron, and hydrogen and oxygen at the ordinary temperature. 
 Even such delicate tests as those of blue litmus and sulphuric acid or hydro- 
 
INTRODUCTION. 
 
 29 
 
 chloric acid do not take place below — 1 10°. It is obvious, therefore, that 
 the power of substances to act chemically upon one another is entirely de- 
 pendent upon the external conditions, particularly the temperature, pre- 
 vailing at the moment of their contact. This has been further demonstrated 
 by the experiments which it has been possible to conduct at very high tem- 
 peratures. It is well known that hydrogen and oxygen unite to form 
 water above 200°, but at 2000° and above, water breaks down again into 
 hydrogen and oxygen (compare Dissociation of Water). The higher 
 the temperature the more complete will be the decomposition, and 
 eventually a point will be reached at which the hydrogen and oxygen will 
 exert as little chemical action upon one another as they would below 200°. 
 
 . A chemical compound is, therefore, only wholly stable within a certain 
 range of temperature. The latter may change, however, with the pres- 
 sure and in the case of some substances we may not be able to obtain it 
 with our present facilities. The compound will begin to separate into 
 its constituents — its elements — just as soon as this limit is exceeded. The 
 rapidity of the decomposition, the magnitude to which it may extend, 
 is also dependent upon pressure and temperature. This is therefore due 
 to the fact that a portion of the product of decomposition reunites to 
 form the original body and in this way a state of chemical equilibrium is 
 produced. The extent of the decomposition is always definite for given 
 external relations. Similarly, the action of unlike bodies upon one an- 
 other is frequently complete only within definite ranges of temperature, 
 and the rapidity with which this action proceeds is in like manner influ- 
 enced not only by temperature but also by pressure and by quantity- 
 relations. Very often opposite reactions, union and decomposition, oc- 
 cur simultaneously, and occasion a state of equilibrium. 
 
 Every chemical change is invariably accompanied by a change of 
 energy — by the disengagement or absorption of heat (electricity, etc.). 
 The customary chemical equations, such as are employed upon page 27, 
 represent merely the material side of a chemical reaction, the nature and 
 quantities by weight of the reacting and resulting substances. But when, 
 for example, 2.02 grams of hydrogen and 16.00 grams of oxygen unite to 
 yield water, there is an accompanying dynamical change, a definite and 
 considerable quantity of heat is disengaged — in this instance equaling 
 68.4 calories (p. 23). An equation showing the union of hydrogen with 
 oxygen to form water — an equation which would include both weight xtdiC- 
 tions and those of energy — would read as follows: 
 
 2H + O = HP -f 68.4 Cal. 
 
 Similarly, in the union of i.oi grams of hydrogen with 35.45 grams of 
 chlorine, forming 36.46 grams of hydrogen chloride, we have a liberation 
 of 22 calories : 
 
 H-f-Cl =HC1 T 22 Cal. ; 
 
 while the formation of hydrogen iodide (127.86 grams) from hydrogen 
 (i.oi grams) and iodine (126.85 grams) is accompanied by an absorption 
 of energy — an absorption of 6 calories : 
 
 H -y I + 6 Cal. = III. 
 
30 
 
 INORGANIC CHEMISTRY. 
 
 But this e(|uati()n is in no manner to he understood as meaning that l)ydro^en and 
 iodine reciuire the addition of l)ut 6 ealories in order that hydrogen iodide may result. 
 The application of that amount of heat would only produee a warmed mixture of hydro- 
 gen and iodine vapor. 'The (juantity of energy e(juivalent to 6 calories of heat can only 
 be taken up by the mixture under certain delinite physical conditions in such a way that 
 hydrogen iodide is the product. 
 
 When we desire to decompose water into hydrogen or oxygen we mtist 
 restore all the energy which has esca^ted. This energy is contained in 
 the free elements as chemical (potential) energy. The decomposition of 
 hydrogen iodide into its elements, on the contrary, occurs with a disen- 
 gagement of heat. This heat is equivalent to the energy which the iodine 
 and hydrogen took up in their i)assage into hydrogen iodide. 
 
 Reactions are distinguished as exothermic and endothermic^ i. e., they 
 have positive or negative the7'7?ial values depending upon whether heat 
 (energy) is liberated or absorbed. 
 
 The heat modulus, attending a chemical reaction, offers important con- 
 clusions as to the concurrent alterations in energy, as well as to the nature 
 of the substances entering the reaction. 
 
 These relations are most evident in those reactions in which only two 
 ele77ients participate. Compounds, formed from their elements with the 
 liberation of heat, contain less energy than the elements themselves, are 
 more stable than their mixture, and can be resolved into the original 
 elements by the consumption of energy. The conditions requisite for 
 their formation are equally present in the mixture of the parent sub- 
 stances, because every system of bodies, as taught by mechanics, strives to 
 attain that state of equilibrium, in which the content of energy, convert- 
 ible into work, of tension or, as Helmholtz terms it, of free energy^ is as 
 low as possible. It is for this reason, therefore, that reactions of this 
 description occur almost immediately on bringing the respective bodies 
 together. An example of this kind has already been given : antimony 
 and chlorine unite instantaneously to antimony chloride (p. 19). As a 
 rule, however, an external impulse from heat, electricity, or light is 
 needed to start the reaction. Several examples illustrating this have been 
 presented. This behavior must be regarded as a sort of liberation of the 
 chemical tension. It is necessary that a portion of the molecules — the 
 physical individuals — first be resolved into their atoms — the chemical ‘in- 
 dividuals — the combination of the atoms in the molecules must be made 
 less intimate at least. When once the reaction has been in this manner 
 begun at one point of the mixture, it generally continues and proceeds of 
 itself, according to the amount of heat developed, with greater or less 
 intensity, which may even reach to explosion, as in the production of 
 water or hydrochloric acid from their elements. 
 
 Com])ounds which absorb heat when produced from their elements, e. g., 
 hydrogen iodide and nitrogen chloride, contain more free energy, more 
 tension, than the parent sul)stances, and are consequently less stable than 
 their mixture, d’hey cannot be formed from their elements without the 
 simultaneous addil ion of energy. In this instance to simply initiate the 
 reaction is not sufficient, for energy must be continually added, otherwise 
 the chemical action will cease. 
 
INTRODUCTION. 
 
 31 
 
 Viewed thermally, the decomposition of compounds into their elements 
 proceeds oppositely to their formation. If the latter was accompanied 
 by heat evolution then the decomposition would occur with heat (energy) 
 absorption, would require the constant addition of energy, advance very 
 gradually, never in an explosive manner, and would be circumscribed by 
 the opposing combination-tendency of the products of the decomposition 
 (compare Dissociation of Water). Just as every transition of a system 
 from a condition of stability to one of less stability requires an expen- 
 diture of work, so does the decomposition of such a body as indicated in 
 the preceding lines. Compounds, however, which like hydrogen iodide 
 and nitrogen chloride are produced with heat absorption, generally break 
 down with ease and completely into their elements. In this way the 
 system passes into a more stable condition. Frequently, an external im- 
 pulse is all that is necessary to start such a decomposition. It then pro- 
 ceeds of its own accord and may increase even to explosion. Many of 
 the bodies belonging in this group are explosive ; indeed, some of them, 
 e. g., chlorine monoxide and nitrogen iodide explode when touched or 
 when warmed, others again require a more energetic concussion. Thus, 
 nitric oxide, acetylene, and cyanogen explode if a slight amount of 
 mercury fulminate be ignited in them. 
 
 The heat modulus is not a measure of the affinity of the elements which 
 combine with one another. Even the formation of water from hydrogen 
 and oxygen, an apparently very simple process, is really the product of a 
 number of chemical and physical changes, proceeding by grades, which in 
 turn are accompanied partly by positive and partly by negative thermal 
 values— breaking down of the molecules into atoms, union of the differ- 
 ent atoms to molecules; diminution of the number of molecules, lique- 
 faction of the aqueous vapor. The quantity of heat observed merely 
 represents the algebraic sum of all these heat moduli (thermal values). 
 
 If more than two elements, if several compound substances, take part 
 in a chemical change the meaning of the thermal value accompanying it 
 is more difficult to comprehend. The majority of reactions of this kind 
 proceed in harmony with the principle of greatest heat development 3 .cc.or&- 
 ing to which from a given system of bodies, without the introduction of 
 external energy, that new compound will result, in whose formation there 
 is the greatest heat development (Berthelot). This principle is not uni- 
 versally acknowledged, neither do all the facts support it, nor is it justified 
 from the standpoint of the mechanical theory of heat. The entropy- 
 principle is a more acceptable substitute (p. 23). When the thermal 
 relations of the various groups of elements and compounds are discussed 
 this point will be more minutely considered. ^ ^ 
 
 the PRINCIPLES OF CRYSTALLOGRAPHY. yv, ^ 'b t 
 
 Solid, homogeneous bodies are composed of the smallest particles, mole- 
 cules, which we imagine as being irregularly placed or arranged in a 
 regular net-like manner. In the first instance they are amorphous, show- 
 ing like physical properties in every direction, while in the second case 
 they are crystallhed, manifesting a similar physical deyiortment in all 
 parallel directions, but which in general varies in different directions. 
 
32 
 
 inor(;anic chemistry. 
 
 In crystallized bodies the molecules are so arranged in layers, that all 
 the particles are similar. These layers form i)lanes, crystal faces, whose 
 form depends on the nature of the molecules; in crystals differing from 
 one another chemically, the form of the faces is also different. 
 
 A combination of all the crystal faces (planes) circumscribing a crys- 
 tallized body constitutes its crystal form, which is always definite for 
 every crystalline chemical comiiound. I'heform and the extension of the 
 individual faces may vary, but the angles produced by the faces remain 
 unchanged. (Law of the constancy of angles formed by crystal faces.) 
 
 A zone is produced by three or more iilanes cutting one another in paral- 
 lel edges ; the zone axis is the direction with whi( h the edges run parallel. 
 
 We can imagine most crystals so divided by planes that each crystal 
 face upon the one side corres[)onds to a similar face on the ojjposite side, 
 producing a like angle with the intersecting plane. Such a i)lane is desig- 
 nated di symmetry-plane ; its i)erpendicular is known as the sy^nmetry axis. 
 
 The symmetry ratios vary with the individual crystals. One may 
 exhibit greater symmetry — more symmetry-planes — than another. All 
 those crystalline forms in which there is a like number of symmetry- 
 planes constitute a system of crystallization. 
 
 There are two kinds of symmetry-planes. Thus crystals exhibit in 
 part one or several symmetry-planes, with which several symmetry axes 
 are parallel and which may be exchanged one for the other without 
 altering the crystalline form. Symmetry axes of this description are said 
 to be equivaleiit, and the symmetry-planes with which they are parallel 
 are called the principal symmetry-pla?ies and their perpendiculars are the 
 principal axes. 
 
 Six systems of crystallization are distinguished on the basis of relations 
 in symmetry. 
 
 I. The regular or isometric system with three principal symmetry- 
 planes and six secondary symmetry-planes. 
 
 II. The hexagonal system with one principal symmetry-plane and six 
 symmetry-planes. 
 
 III. The quadratic or tetrag07ial systetn with one principal symmetry- 
 plane and four secondary symmetry-planes. 
 
 IV. The rhombic system with three secondary symmetry-planes. 
 
 V. The mo7ioclinic system with one secondary symmetry-plane. 
 
 VI. The triclinic system, in which there is no symmetry-plane. 
 
 I. The Isometric System. — The forms of this system are referred 
 to an axis system consisting of three axes of equal length (^principal 
 sym?fietry axes) and at right angles to one another. Two of the axes lie 
 in a ])rincii)al symmetry-})lane. The various fundamental forms are de- 
 rived by imagining these axes a, a, a, cut by planes at equal or unequal 
 distances from tlieir point of intersection. The axial section of a plane 
 is termed its |)arameter. If the smallest of these be designated by a, and 
 the other two by na and ma, then the values n and m (the coefficients of 
 the parameters) become in accordance with experience rational numbers. 
 'I'here are seven fundamental forms: 
 
 I. Octahedron {O) (h'ig. i). 'Fhe faces intersect the axes at equal 
 distances a : a : a from the center. 
 
INTRODUCTION. 
 
 33 
 
 2. Cube (ood?oo) (Fig. 2). The faces intersect one axis and are 
 
 parallel to the other two : a : : oo^r. 
 
 3. Dodecahedron ( 00 (Fig. 3). The faces intersect two axes at the same 
 distance from the center, while they are parallel to the third axis : a \ a: ooa. 
 
 4. Trisoctahedron (mO) (Fig. 4). The faces intersect two of the axes 
 at unity and the third at a greater distance : a : a : ma. 
 
 5. Trapezohedron or Jcositetrahedron {mOin) (Fig. 5). The faces 
 intersect one axis at unit distance and the other two at equal but greater 
 distances: a : nia : ma. 
 
34 
 
 INORGANIC CHEMISTRY. 
 
 6. Tetrahexahedro 7 i ( o/^Oti) (Kig. 6). Tlic faces intersect two axes at 
 different distances and the thiixl axis at infinity : a : 7 ux : cc/'/. 
 
 7. Ilexoctaliedron (xnOn) (h'ig. 7). The faces cut all three axes at 
 different distances : a : ;ia : xna. 
 
 These siixiple forms usually occur in combination; this is also true of 
 the remaining systems. Thus Fig. 8 rej)resents the combination of 
 dodecahedron ( 006?), cube ( oo(9co), and octahedron (6^). 
 
 II. The Hexagonal System. — 'I'here are seven symmetry-j)lanes in 
 this system, six of which are at right angles to the seventh — the iirinciiial 
 symmetry jilane. The six planes referred to cut one another at an angle 
 of 30°. The intersecting lines of three alternating ])lanes (cutting one 
 another at 60°), together with the princii)al symmetry-j)lane, are regarded 
 as the secondary axes (these are equal a : a : a). The fourth axis is the 
 axis (c) at right angles to the principal symmetry-plane. The ratio : c 
 is irrational and is very definite for every substance crystallizing in the 
 hexagonal system. 
 
 P 
 
 Fig. 9. 
 
 The forms of this system are either double consisting of twelve 
 
 or six faces converging above and below at the terminals of the principal 
 axis, or they are twelve- or six-sided prisms, the faces of which are parallel 
 to the principal axis. To these must be added the pair of faces perpen- 
 dicular to the principal axis — the basal planes. Fig. 9 represents the 
 hexagonal pyra 7 nid (^P) and Fig. 10 the hexagonal pyra 77 iid (^P) with the 
 pris 77 i ( ooP). The angle produced by the intersection of two prism faces 
 ecpials 60°. 
 
 III. Tetragonal or Quadratic System. — This system like the 
 hexagonal is characterized by a single principal symmetry-plane to which 
 four alternating similar symmetry-planes cutting one another at 45° are 
 fierpendicular. d'here are three axes at right angles to one another. 
 'I'he ))rinci])al axis (c) is normal to the principal symmetry-plane, while 
 the other two (a, a) correspond to the intersecting directions made by two 
 ecpial symmetry-])]ancs with the principal symmetry-plane. 
 
 'The forms of this system, as in the hexagonal system, consist of pyra- 
 mids, jirisms, and basal i)lane (})inacoid). The pyramids are made up of 
 
INTRODUCTION. 
 
 35 
 
 four or eight faces above and below. The i)risms also have four or eight 
 laces; e. Fig. ii: tetragonal pyratnid (^P)\ Fig. 12: tetragonal 
 prism ( 00^) with the pyramid {P"). The angle formed by two prism 
 faces is 90°. 
 
 IV. The Rhombic System. — In this system there are three dis- 
 similar symmetry-planes at right angles to one another. Their directions of 
 intersection are assumed to be the axes a, I?, c. The rhombic pyramid (P) 
 (Fig. 13) is the fundamental form. It cuts the three axes (of unequal 
 lengths) at their unit distances (^zis the brachy-axis, b the macro-axis, and 
 c the vertical axis'). There are other pyramids in which the faces intersect 
 
 Fig. 12. 
 
 Fig. 13. 
 
 Fig. 14. 
 
 the axes at other than unit distances from the center. Prismatic forms re- 
 sult when the faces are parallel to one axis. If they are parallel to c the 
 rhombic prism ( ccP) (Fig. 14), and macro- or brachydomes if the faces 
 run parallel to b or a, e. g., Fig. 15, the 7 nacrodome. 
 
 The pair of faces perpendicular to the vertical axis constitutes the basal 
 planes, while the other pairs of faces at right angles to the other two axes 
 are termed the macro- and brachy-pinacoids. 
 
 V. The Monoclinic System. — There is but one symmetry-plane in 
 this system. It is directed toward the observer when he examines the 
 
36 
 
 INORGANIC CHEMISTRY. 
 
 crystals. The axis system consists not (mly of tlie perpendicular (/^-axis) 
 to the syminetry-])lane, but also (d' the intersecting lines und c), \vhi( li 
 form two crystal faces with and at right angles to the symmetry-plane and 
 are inclined to one another at any arbitrary angle (fi). The vertical axis 
 is c, while a, directed toward the observer, is designated the clino-axis. 
 The forms are like those of the rhombic system, l^iach ])yramid, Innvever, 
 resolves itselt into two independent / these are distinguished 
 as positive (-(- P) and negative ( — P) (Fig. i 6 ). d'here are positive and 
 negative orthodomes, each arising from a jiair of faces, which correspond 
 to the rhombic macrodomes. The prismatic form, i)arallel to the clino- 
 axis, is termed the cli nodome. 
 
 VI. Triclinic System. — Any three crystal planes are selected for 
 the axis planes. Tneir intersecting lines constitute the axes zz (brachy- 
 diagonal), b (macrodiagonal), c (vertical axis). All the axes intersect 
 one another at oblicpie angles. The angle between b and c is designated 
 a, that between a and c (3, and that between a and b y. There is no 
 symmetry-plane present; each parallel pair of faces constitutes a crystal 
 form, so that a pyramid intersecting the three axes at the distances a: b: c 
 consists of four different crystal forms which are marked as in Fig. 17 
 with P', P^, 'P, ^P. All the other forms, as in the monoclinic system, are 
 derived from these, with the difference that in this system the distinction 
 between right and left must be observed. 
 
 In nature, crystals rarely occur so regularly developed as represented 
 in the preceding forms. Usually they are more or less elongated in one 
 or more directions and this causes the faces of one and the same form to 
 become, as regards their spacial extension, unequal and the whole crystal 
 appears to be distorted. Such a distortion, however, never influences 
 the position of the faces with reference to .the axes. The angles produced 
 by the faces remain unchanged so long as the temperature is the same. 
 (Law of the constancy of the angles formed by the faces.) The measure- 
 ment of the angle formed by the faces by means of the goniometer is the 
 only means we possess of unraveling complicated combinations, and even 
 this aid is not in all cases satisfactory in determining definitely and surely 
 that a form belongs to one or another system. For this purpose a careful 
 investigation of the physical properties is frequently essential, because it 
 occurs that crystals apparently develop a higher symmetry geometrically 
 than belongs to them when their physical properties are considered 
 (pseudo-symmetric crystals). 
 
 Another irregularity in their development consists in the fact that the 
 faces themselves are not even and smooth. They frequently occur bent, 
 e. g., in the diamond and dolomite; drusy, due to the protuberance of 
 numerous little solid angles, uj)on the faces, of other forms differently 
 orientated, or striated by the frequent alternating appearance of two 
 faces— g., the striation of the cube faces of pyrite due to the oscillatory 
 a])))earancc of a ])cntagonal dodecahedron — rough, corroded, etc., etc. 
 
 'I’wo or more crystals are often grown together; this growth being 
 either regular or irregular. When regular the crystals are ]xarallel to one 
 aiu>ther and are designated pa7’anel growths, or they are not ])arallel but 
 yet limited by a very definite regularity. Such growths are called twins. 
 
INTRODUCTION. 
 
 37 
 
 Their regularity consists in the fact that the two crystals forming the 
 twin are symmetrically developed with one another in relation to a crys- 
 tallographically possible plane — the iw inning-plane. Fig. i8 represents 
 an example of such a twin crystal. Two octahedra, or better, two half 
 octahedra, are here symmetrical in reference to an octahedral face. 
 Spinel and magnetite are examples of this class. The plane along which 
 the two individuals have developed is the growth-plane. It is not neces- 
 sary that this should at the same time be the twinning-plane, for two 
 crystals can be symmetrical to one another in referenc e to one plane and 
 yet be developed along another plane. This, for example, is often the case 
 with what are termed the Carlsbad orthoclase twins. In them the clino 
 pinacoid is the growth-plane and the orthopinacoid the twinning-plane. 
 
 'T'/ 
 
 Fig. I 8. 
 
 Fig. 19. 
 
 oP 
 
 Fig. 21. 
 
 Fig. 20. 
 
 Crystals may also grow on both sides beyond the growth-plane, 
 tration-twins or crosses are the result. 
 
 In addition to the simple forms appearing with all their faces, hence 
 termed holohedral for7n^ others occur in which only half the possible 
 faces are present — the heniihedral form. These result if we imagine a 
 holohedral form to be divided by symmetry-planes into congruent parts, 
 and then have one of every two parts, which are symmetrical in reference 
 to these planes, fall away. In this manner every holohedral figure will 
 yield two heniihedral forms which differ from one another in their 
 position, and are either congruent or symmetrical. Thus, the octahedron 
 of the isomeric system yields the teti'aliedron, Fig. 19, and the hexagonal 
 pyramid, the rhombohedron (Fig. 20). 
 
 A comparatively rare phenomenon has been observed in the crystals of 
 
38 
 
 INORGANIC CHEMISTRY. 
 
 some substances — namely, a different development at the opposite ends of 
 an axis. This is known as he77iimorphi5m. Struvite (magnesium ammo- 
 nium phosphate Mg(NH 4 )r 04 + 6 HjO), Fig. 21 , is an exampleof this. 
 On the upper end of the vertical axis occur the domes P 00 and P ^ , 
 while below there is only the dome P^ and the basal plane oP. Tour- 
 maline, calamine, cane sugar, etc., exhibit the same phenomenon. 
 
 Substances crystallizing variously in the same, or in two or three 
 different systems are said to be di77iorphous, tnTtiorphous, etc. Titanium 
 dioxide (TiOa), for example, is found in nature in crystals of the quadratic 
 system as the minerals anastase and rutile, and in those of the rhombic 
 system as brookite (see also Sulphur). 
 
 Various substances which crystallize in similar or in forms very much 
 alike in the same system are called iso77io7phous {sqq also Isomorphism). 
 

 SPECIAL PART. 
 
 ■ - " CLASSIFICATION OF THE ELEMENTS. 
 
 Ordinarily we are accustomed to divide the elements into two groups : 
 metals and non-metals (seep. 19). The former possess metallic appear- 
 ance, are good conductors of heat and electricity; the latter, the metal- 
 loids or non-metals, do not have these properties, or at least in less degree. 
 In chemical respects the metalloids have the tendency to combine with 
 hydrogen, forming volatile, generally gaseous, compounds; their oxygen 
 derivatives form acids with water. The metals, on the contrary, do not 
 unite with hydrogen, or at least do not form volatile compounds with it, 
 and their oxygen derivatives yield chiefly the so-called bases with water. 
 Further, the compounds of metals with the non-metals are so decomposed 
 by the electric current that the metal separates at the electro-negative, 
 and the non-metal at the electro-positive pole. From this we observe 
 the metals are more electro-positive — more basic; the metalloids more 
 electro-negative — of an acid-forming nature. A sharp line of difference 
 between metals and metalloids does not exist. There are elements, like 
 antimony, which in their external appearance resemble metals, while in 
 a chemical respect they deport themselves throughout as metalloids, and 
 vice versa. Thus hydrogen, a gaseous element, is like the metals in its 
 entire chemical character, while metallic antimony arranges itself with 
 the metalloids. 
 
 It is therefore better to divide the elements into separate natural groups, 
 based upon their chemical analogies. The best and only correct classifi- 
 cation of all the elements depends on the law of periodicity, according 
 to which the properties of the elements and of their compounds present 
 themselves as a periodic function of the atomic weights. Later we shall 
 treat of the periodic system more at length ; it forms the basis of this 
 text-book, and in accordance with this doctrine we consider the elements 
 in single natural groups of similar chemical deportment. The first of 
 these groups, comprising almost all the so-called non metals, are the 
 
 following : 
 
 Fluorine 
 
 Oxygen 
 
 Nitrogen 
 
 Carbon 
 
 Boron 
 
 Chlorine 
 
 Sulphur 
 
 Phosphorus 
 
 Silicon 
 
 
 Bromine 
 
 Iodine 
 
 Selenium 
 
 Tellurium 
 
 Arsenic 
 
 Antimony 
 
 
 
 . Hydrogen does not belong to any of these groups; uniting the metal- 
 lic and non-metallic characters in itself, it represents, as it were, the type 
 of all elements, and therefore it will receive first attention. Boron 
 
 39 
 
40 
 
 INORGANIC CHEMISTRY. 
 
 occupies an isolated position. It has been classed with the non-metals, 
 but differs somewhat from them in chemical deportment. It forms the 
 transition to the metallic elements, beryllium and aluminium. 
 
 HYDROGEN. 
 
 Atom: H = i.oi. Molecule: Hg = 2.02. 
 
 Hydrogen (hydrogeniuin), a gaseous body, occurs rarely in a free con- 
 dition upon the earth’s surface, — in the gases from volcanoes and in 
 
 those issuing from the earth, 
 — as an enclosure in minerals, 
 a product of decay, and, from 
 recent statements, it is found 
 in very small amount in the 
 atmosphere. It is, however, 
 present in considerable quan- 
 tity in the photosphere of the 
 sun and fixed stars. In com- 
 bination, it is found chiefly 
 as water, and in substances of 
 vegetable and animal origin. 
 Paracelsus first observed this 
 element in the sixteenth cen- 
 tury, and called it inflamma- 
 ble air; and in 1766 Cav- 
 endish, recognizing it as a 
 peculiar gas, named it inflam- 
 mable air. In 1 783 Lavoisier 
 proved that hydrogen was a 
 constituent ot water — a chemical compound of the elements hydrogen 
 and oxygen — by conducting steam over ignited metallic iron. 
 
 Preparation . — It may be readily obtained from water. The decompo- 
 sition of the same by the removal of oxygen can be effected by some metals, 
 like sodium and potassium, at the ordinary temperature. Both metals act 
 very energetically upon it, liberating gaseous hydrogen. To perform the 
 experiment, take a piece of sodium, roll it up in a piece of wire gauze, 
 and shove it, with nippers, under the mouth of a glass cylinder filled with 
 and inverted over water (Fig. 22). Bubbles of hydrogen are at once 
 disengaged, displace the water and collect in the upper part of the cylin- 
 der. The reaction occurring between the sodium and water is expressed 
 by the following chemical equation : 
 
 11,0 -f Na NaOH + H. 
 
 Water. Sodium. Hydrogen, 
 
 I he compound NaOH, known as sodium hydroxide, remains dissolved in 
 the excess of water. 
 
 Other metals decom))ose water in a similar manner, at an elevated 
 temperature. To effect this with iron allow steam to pass through a tube 
 
CLASSIFICATION OF THE ELEMENTS. 41 
 
 filled with iron filings, exposed to a red heat in a combustion furnace, 
 d'he iron withdraws oxygen from the water, combining with it, while the 
 hydrogen set tree is collected. Magnesium powder reacts similarly, but 
 at a much lower temperature, upon steam (Ber. 26 (1893), I, 59). 
 
 For laboratory purposes, hydrogen is prepared by the action of zinc 
 upon hydrochloric or sulphuric acid. The reaction with the latter acid 
 is as follows : 
 
 Zn + H^SO, ZnSO, + H.,. 
 
 Sulphuric acid. Zinc sulphate. 
 
 Place granulated zinc (obtained by dropping molten zinc into water) 
 in a double-necked flask (Fig. 23), and introduce sulphuric acid (diluted 
 with about 3 vols. of H2O) through the funnel tube, b. The liberation 
 of gas begins immediately, and the hydrogen, escaping through the exit 
 tube, f, is collected as previously described. 
 
 Fig. 23. 
 
 The hydrogen thus formed has a faint odor due to a slight admixture 
 of foreign substances (the hydrides of sulphur, arsenic, phosphorus 
 and carbon — if these elements are contained in the metal used in the gas 
 evolution). It is therefore conducted through a solution of potassium 
 permanganate to purify it. 
 
 Many other metals, e. g., iron, behave like zinc with dilute acids. 
 Some metals — zinc, iron, and aluminium — dissolve, when finely divided, 
 in sodium or potassium hydroxide with the liberation of hydrogen : 
 
 Zn -f 2NaOH = Zn(ONa )2 + 
 
 Pure hydrogen may be obtained by heating potassium formate with 
 potassium hydroxide: CHO^K -f- KOH =: K2CO3 -(- H2 ; for technical 
 purposes by heating zinc or iron with calcium hydroxide (slaked lime) in 
 a combustion tube : Zn -(- Ca02H2 = ZnO -j- CaO -f- Hj, or with coal 
 (anthracite) : 
 
 2Ca02ll2 + c = CaCOg + CaO -f- 2H2. 
 
 4 
 
42 
 
 INORGANIC CHEMISTRY. 
 
 A very important method for tlie i)rei)aration of hydrogen consists in 
 decomposing certain arpieons solutions by means (jf tlie electric current. 
 This decom])osition usually proceeds as if the water liroke down into its 
 constituents : 2H./ ) — 2I I2 -|- 0 . 2 . The hydrogen ajipears at the negative 
 pole. The electrolysis of such solutions will be discussed exhaustively later. 
 
 Purifying atuf D)yjng of Gases. — To free gases of the substances mechanically carried 
 along during their disengagement, it is best to conduct them through variously constructed 
 wash-bottles, filled with water or liipiids, which will absorb the impurities. Ordinarily the 
 so-called WoullT bottles are employed (compare Figs. 31 and 35). 'J’he open tube, 
 
 placed in the middle tubulure, is called the 
 safety tube. It serves to equalize the inner 
 pressure with that of the external atmos- 
 phere. Wash-bottles of various construction 
 will be represented in the several sketches. 
 
 Gases liberated from an aqueous liquid 
 are always moist, as they contain aqueous 
 vapor. To remove this conduct them 
 through vessels or tubes filled with hygro- 
 scopic substances (see Fig. 30). Calcium 
 chloride, burnt lime, sulphuric acid, etc., 
 are used for this purjjose. 
 
 Apparatus for the Gene?-ation and Col- 
 lection of Gases. — In the ap])aratus pictured 
 in Fig. 23, the liberation of hydrogen con- 
 tinues uninterruptedly as long as zinc and 
 sulphuric acid are present. To control the 
 generation of the gas we have recourse to 
 different forms of apparatus. One of the 
 most practicable of these is that of Kipp. 
 It consists of two glass spheres, d and 
 Fig. 24, in the upper opening of w'hich 
 there is a third sphere, c, fitting air-tight 
 and provided with an elongated tube. 
 It serves as a funnel. Granulated zinc is 
 placed in the middle sphere through the 
 tubulure <?, and dilute sulphuric acid is 
 poured into the spherical funnel, which 
 first fills d, then ascends to where it comes in contact with the zinc ; at once the evo- 
 lution of hydrogen commences and the gas escapes through e. Upon closing the stop- 
 cock of the tube fixed in e, the hydrogen which is set free presses the sulphuric acid out of 
 b, and consecpiently the liberation of the gas ceases. On again opening the stop-cock, the 
 acid rises in b to the zinc, and the evolution of gas commences anew. The vessel a con- 
 tains water or concentrated sulphuric acid to wash the escaping hydrogen. 
 
 The .somewhat complicated Kipp apparatus may be advantageously replaced by the 
 following siinjde contrivance recommended by Debray (Fig. 25), Two bottles provided 
 with openings near their bottom, in which are glass tubes, are connected by a rubber tube. 
 4 'he bottle A is filled with granulated zinc, and B with dilute sulphuric acid. The cock 
 R clo.ses A. When this is opened the sulphuric acid flows from B to A, to the zinc, 
 and the evolution of gas commences. D represents a form of wash-bottle suggested 
 by Kemi)f. The gas evolution can be quickly regulated by raising or lowering the 
 Inktle /)’. 
 
 'fhe Mohr ap])aratus ( Fig. 26) is very convenient for the evolution of gases. Into the 
 second cylinder containing the acid another narrower cylinder closed by a stop-cock is 
 introduced, 'fhe zinc (iron, marble, etc.) is ]daced above the contraction of the inner 
 cylinder. On opening the stop-cock the acid enters aiuk evolution of gas commences. 
 On closing the stopcock the hydrogen pres.ses the acid back from the inner cylinder; tlae 
 evolution of gas cease.s. 
 
 S 
 
 Fig. 24. 
 
CLASSIFICATION OF 'I HE ELEMENTS. 
 
 43 
 
 Gasometers of various construction serve to collect and preserve gases. In Fig. 27 we 
 have the ordinary gasometer of Pepys. It is constructed of sheet copper or zinc, and 
 consists of two cylindrical vessels, the lower one closed, the upper open, communicating 
 with each other by the two tubes a and d. The tube c is only a support. To this end, 
 pour water into the upper cylinder, and open a and tr ; the water then flows through a, 
 nearly reaching the bottom of the lower cylinder, while the air escapes through e. When 
 the lower cylinder is filled with water close a and e (the last traces of air can be removed 
 by opening b'). To fill the gasometer with gas, remove the cover of the side tubulure 
 d, and introduce the tube from which the gas is escaping. The latter rushes up into the 
 cylinder, while the water flows out the tubulure. When the water is displaced by the 
 gas, close d, after filling the upper cylinder, and then, if desired, open a, and the 
 gas can be set free, either by e or b. 
 
 In addition to the gasometer described, various other forms are employed ; gas-bags 
 and tubes are very well adapted for preserving gases. 
 
 Physical Properties . — Hydrogen is a colorless, odorless, and tasteless 
 gas. It conducts heat and electricity better than all other gases. This 
 
 Fig. 25. 
 
 Fig. 26. 
 
 may be readily proved by the following experiment : A current of elec- 
 tricity is sent through a thin platinum spiral, and while the latter remains 
 in the air or some other gas, it will glow, but in an atmosphere of hydro- 
 gen the spiral will not become luminous, or at once cease glowing. 
 The metallic nature of hydrogen is shown by this behavior. 
 
 Of all gases hydrogen is the most difficult to liquefy, because its criti- 
 cal temperature is the lowest (about — 220°); it must, therefore, be ex- 
 posed to the most intense cold. 
 
 The procedure first adopted by Olszewski and more recently by Dewar 
 for the production of liquid hydrogen (also ajiplied to other gases) con- 
 sists in suddenly lowering the pressure of the sufficiently cooled and 
 powerfully compressed gas, when it will rapidly expand and a portion of 
 it will deprive another portion of so much heat that the second portion 
 will liquefy. Dewar permitted hydrogen cooled to — 205° and under a 
 pressure of 180 atmospheres to enter a vacuum-vessel also cooled to — 200°. 
 
44 
 
 INORGANIC CHKMISTKY. 
 
 'File liquid hydrogen i)rei)ared in this way is colorless, very nioliile and 
 has the specific gravity of 0.07. It boils at about —240° under the ordi- 
 nary atmospheric pressure. 
 
 Like all gases, coercible with difficulty, hydrogen is but slightly soluble 
 in water, 1000 volumes dissolving 19 volumes of the gas at 14° C. 'Fhe 
 coefficient of absori)tion of hydrogen by water is therefore 0.0193. It is 
 the lightest of all gases. That is, a definite volume of it weighs less than 
 an equal volume of any other gas at the same temperature, under 
 
 like iiressure, and subject to a like 
 action of gravity. It is generally 
 understood that the comparison 
 of the weight of equal gas volumes 
 is made at 0°, 760 mm. barometric 
 pressure at sea-level and in the 45° 
 of geographical latitude. These 
 are normal conditions, and under 
 them I liter of hydrogen weighs 
 0.08988 gram, i liter of oxygen: 
 1. 4291 grams, and i liter of air: 
 1.2930 grams. The thousandth 
 part of each number, the weight of 
 one cubic centimeter, is the 
 specific gravity of the gas referred 
 to water at 4°.* 
 
 Fig. 27. Fig. 28. 
 
 From this it follows that air is 14.4 (= times heavier than 
 
 hydrogen, and oxygen 15.9 ^ ) times heavier than the latter. 
 
 'Fhese numbers are the densities of these gases referred to hydrogen as unit. 
 If, however, the Aveight of the air volume is selected as unit we obtain 
 0.06951 (— ^1^.2 9 3^ ) density of hydrogen, and 1.T052 (— 
 
 as that of oxygen. It has become the rule quite recently in chemical 
 discussions to refer densities to that element which is taken as the standard 
 in atomic weights : oxygen, and for reasons which will be given subse- 
 
 * f'or llic; (let(‘nnitiatif)ii of these numbers compare Morley, Z. j^hys. C^i. 17 (1895') 
 87; 20 (]H()G) 68, 242, 417, and 4 'Iiom.sen, Z. f. anorg. Cli. 12 (1896) I ; also Kohl- 
 rau.sch, Lcilfaden der prakt. Physik, 8 Aull. (1896), 459. Morley gives as the most prob- 
 able values for the weight of a liter of liydrogeii and that of oxygen : 0.089873 and 
 1.4290; 'riuun.sen : 0.089947 and 1.42906. 
 
CLASSIFICATION OF THE EI.EMENTS. 45 
 
 quently, the weight of its volume is placed at 32. Then for hydrogen 
 there would result the number 2.0126 and for air 28.954. 
 
 The volume expressed in cubic centimeters which i gram of gas would 
 occupy under normal conditions is its specific volume. If the weight of one 
 cubic centimeter of gas is represented by a, its specific volume would be 
 Vgrrr-T- Heuce the specific volume of hydrogen is = 1 1, 1 26 ; 
 
 that is, one gram of hydrogen under normal conditions occupies 11,126 
 cubic centimeters ( 1 1. 1 2 liters). The specific volume of air is q ^ 293 
 = 773-4 c.c,, and that of oxygen =699-7 c.c. 
 
 That hydrogen is lighter than air is shown by a balloon of collodion 
 or gum filled with the former rising in the latter ; this can also be seen 
 in soap bubbles filled with hydrogen. In consequence of its levity, 
 hydrogen may be collected in inverted vessels (opening turned down) by 
 displacing the air, and can also be poured from one cylinder into another, 
 as represented in Fig. 28. The hydrogen flows from the inclined 
 cylinder into the one held vertically and filled with air, which it expels. 
 Such a separation of gases, based on their varying specific gravity, is only 
 temporary, as they soon mingle with each other by diffusion. By virtue 
 of its levity and mobility, which the kinetic gas theory attributes to the 
 great velocity of the gas ])articles, hydrogen penetrates porous bodies with 
 ease, and diffuses througli both animal and vegetable membranes, as well 
 as through gutta-percha (consult Air, upon diffusion of gases). Metals, 
 e. g., iron, platinum, palladium, permit a free passage to hydrogen, when 
 they are raised to a red heat. They are impenetrable to other gases (see 
 Dissociation of Water). This behavior is in part probably dependent upon 
 the chemicalattraction of these metals for hydrogen. 
 
 Chemical Properties . — Hydrogen is character- 
 ized by its ability to burn in the air, when it com- 
 bines with the oxygen of the latter and forms 
 water; hence its name hydrogenium, due to Lavo- 
 isier (from water, and y&wdoi, I produce). 
 
 Its flame is faint blue, and almost non-luminous, 
 but possesses a very high temperature. When a 
 mixture of hydrogen and air is ignited a violent 
 explosion ensues ; therefore, before bringing a 
 light in the vicinity of hydrogen disengaged in a 
 vessel filled with air, allow the latter to escape 
 completely, otherwi.se the vessel will be shattered 
 to pieces by the explosion. 
 
 As hydrogen itself is inflammable, it cannot 
 sustain the combustion of other bodies which will 
 burn in the air. If a burning candle be introduced into an inverted 
 cylinder containing the gas (Fig. 29) the latter will ignite at the mouth 
 of the vessel, but the candle will be extinguished. 
 
 Water is the product of the combustion of hydrogen in the air. It is 
 a chemical compound containing hydrogen and oxygen. To render the 
 formation of it visible, by the combustion of hydrogen, the flame of the 
 latter is made to burn under a cold glass jar (Fig. 30). The sides of the 
 
46 
 
 INORGANIC CHKMISTRY. 
 
 latter are soon covered vvitli moisture, whicli collects in drojis. d'o avoid 
 any deception the hydrogen is first conducted through sulphuric acid or a 
 tube filled with calcium chloride, to absorb all moisture. 
 
 At the ordinary temperature hydrogen and oxygen do not react with- 
 out the intervention of a third substance. 
 
 The union of hydrogen with oxygen to form water occurs at about 
 200° (compare p. 28), if brought in contact with aflame, or by the pas- 
 sage of an electric sjiark through the mixture. The combination can be 
 effected at ordinary temperatures with the aid of platinum sponge or ])al- 
 ladium (platinum sponge, palladium asbestos) ; the jilatinum sponge con- 
 sists of finely divided metal, obtained by the ignition of ammonio- 
 idatinum chloride (see Platinum). If a stream of hydrogen be directed 
 upon a piece of freshly ignited platinum sponge, the gas will at once 
 ignite. This is due to the power of the platinum to condense hydrogen 
 
 Fig. 30, 
 
 and oxygen upon its surface, and thereby increase their ability to unite 
 (^Dxbereiner lamp). 
 
 The absorption of hydrogen by the metal palladium is very character- 
 istic. As already known, water is so decomposed by the electric current 
 that hydrogen separates at the electro-negative pole and oxygen at the 
 electro-positive. Now, if a piece of palladium, in sheet or wire form, be 
 attached to the electro-negative pole, the disengagement of hydrogen 
 does not occur, because it is absorbed by the palladium, in a quantity 
 about eight hundred and fifty times the volume of the latter. Palladium 
 also absorbs hydrogen when it is heated to 200°. The palladium expands, 
 becomes lighter in weight, but retains its metallic appearance. Its 
 tenacity and ])ower of conducting heat and electricity are but little im- 
 paired. 'The compound of palladium and hydrogen, therefore, conducts 
 itself like an alloy of two metals. From the specific gravity of the com- 
 f)ound (according to Oraham), the specific gravity of the condensed 
 hydrogen is found to be 0.62 (water i), and is, therefore, somewhat 
 
CONDENSATION OF GASES. 
 
 47 
 
 heavier than the metal lithium (compare Palladium). The metals 
 potassium and sodium absorb hydrogen when heated from 200° to 400°, 
 forming alloys (Na4H2 and K4H2) in which the density of hydrogen is 
 again equal to 0.62. The experiments of Winkler and of Moissan indi- 
 cate that Other metals also unite with hydrogen: barium, calcium, stron- 
 tium, yttrium, lanthanum, cerium, and thorium. These facts prove the 
 metallic character of hydrogen, which is also indicated by its ability to 
 conduct heat and electricity. Later, we will observe that this element 
 displays the character of a metal in its entire chemical deportment, and 
 that it must be regarded as a gaseous metal at ordinary temperatures. 
 
 CONDENSATION OF GASES. 
 
 CRITICAL CONDITION. 0 
 
 Hydrogen and several other gases (oxygen, nitrogen, carbon monoxide, 
 methane, nitric oxide), were considered until 1877 non-condensable — 
 permanent gases, inasmuch as all attempts to liquefy the same were fail- 
 ures, notwithstanding Natterer (1852) had employed a pressure of 3,600 
 atmospheres for this purpose. These negative results find their explana- 
 tion in a general property of gases, first recognized by Andrews (1869), 
 and called by him the critical conditio 7 i of matter. There is a temperature 
 common to all gases, above which they cannot be condensed — this is the 
 critical temperature. It was first observed with carbon dioxide (see this). 
 It can be liquefied at 0° under a pressure of 35.4 atmospheres, at 13° 
 under 48.9 atmospheres, and at 30° under 73 atmospheres. Again, Cag- 
 niard de la Tour (1821) showed that all liquids when heated above a cer- 
 tain temperature (the same critical temperature), would be transformed 
 into gases (absolute boiling point of Mendelejeff), although they were 
 subjected to intense pressure (in sealed tubes). The pressure exerted by 
 the gas at the critical temperature (at which it would immediately con- 
 dense upon lowering the temperature) is called the critical pressure ; the 
 volume occupied by the substance at this time is the critical volwiie. 
 
 An additional interesting fact is the behavior of a liquid when it is 
 gradually heated under a pressure which exceeds its critical pressure. 
 When this is done the temperature can be increased to any extent with- 
 out producing a noticeable separation into liquid and gas. At the critical 
 temperature the liquid changes to gas without becoming heterogeneous 
 and without sudden increase in volume. At the critical temperature and 
 pressure the volume of the gas (or saturated vapor) is equal to the volume 
 of an equal quantity, by weight, of the liquid. A difference between the 
 gaseous and liquid condition no longer exists. 
 
 The critical condition also furnishes us with a criterion for the distinc- 
 tion between a gas, in the strictest sense of the word, and a vapor. A 
 vapor by increase of pressure with constant temperature may become a 
 liquid (or solid), while this is not possible with a gas. A gaseous body 
 is, according to this idea, a permanent gas above the critical temperature, 
 while below the same it is a condensable vapor. 
 
48 
 
 INORGANIC CIIKMIS'J’RY. 
 
 'To licjiiefy gases, we need nut only pressure, but also a definite tem- 
 perature, and this must be lower than the critical. I>y this means 
 Cailletei and Pictet ( 1 877) succeeded in condensing nearly all the ptr- 
 manent gases. Pictet })ursued the method of Paraday, who had con- 
 densed various gases in sealed tubes (see Condensation of Cddorine, p. 51). 
 'J'he gases were generated in a powerful iron retort (oxygen from potas- 
 sium chlorate ; hydrogen from sodium formate) by the ai)plication of heat. 
 They were then com])ressed under their own jjressure in a coi)per tube 
 attached to the retort. Solid carbon dioxide surrounded the tube, and 
 by its evaporation under the air i)ump its temperature was reduced to 
 — 140°. On opening the stoi)-cock of the coi)per tube, the liquefied 
 gas escaped in a stream which rajiidly evaporated. Cailletet employed a 
 capillary glass tube, provided with a reservoir and a pressure ])ump. 
 d’he strongly conqiressed gas was cooled by opening a stoj^-cock and [ler- 
 mitting it to expand suddenly. In its exj^ansion and in overcoming the 
 external pressure it i)erforms work and there follows an absorj)tion of an 
 ai)})reciable quantity of heat, which is taken from the gas. This causes a 
 partial liquefaction of the gas in the form of a dense cloud, or in small 
 drops. 
 
 With these facts as a basis Wroblewsky, Olszewsky, and more especially 
 Dewar, Linde, and Ramsay have succeeded in condensing and retaining 
 all gases in liquid form ; indeed most of them have been solidified. 
 Usually Cailletet’s method has been pursued, care be taken to thoroughly 
 cool the gas before the sudden decrease in pressure occurs (see Hydrogen 
 and Air). Pictet’s suggestion to strongly chill a greatly condensed gas 
 has proved satisfactory for the condensation of several permanent gases 
 to liquid form. Carbon dioxide, liquid ethylene, liquid oxygen, and 
 nitrogen, which vaporize at slight pressure, have also been employed to 
 produce the reduction in temperature. With a pressure of 10 mm. the 
 temperature of liquid ethylene is reduced to — 150°, liquid oxygen at 
 9 mm. to — 211°, and liquid nitrogen to — 225°. Lower temperatures 
 than these, about — 250°, are only attainable by the evaporation of liquid 
 hydrogen. 
 
 Liquefied gases can be preserved quite well in Dewar bulbs : open, 
 double-walled glass vessels, the intermediate space of which is a vacuum 
 and the inner wall coated on its exterior with a thin layer of mercury. 
 The conduction and radiation of heat are thus reduced to a minimum. 
 
 The critical temperatures (T) and critical pressures in atmospheres (P) 
 of the gases condensed with difficulty are as follows: 
 
 Carl)on Dioxide, CO^, 
 I'khylene, C.^lf,, . . . . 
 Nitric Oxide, NO, . . . 
 Marsh Oas, ClI^, . . . 
 Oxygen, O.^, . . . 
 
 C’arhon Monoxide, CO, . 
 
 Nitrogen, N.„ 
 
 Hydrogen, 11^, . . . . 
 
 Air, 
 
 T 
 
 P 
 
 + 31° 
 
 77 atmos. 
 
 + 10° 
 
 51 “ 
 
 - 93 ° 
 
 71 “ 
 
 — 82° 
 
 55 “ 
 
 — 118° 
 
 50 “ 
 
 —140° 
 
 35 “ 
 
 — 146° 
 
 35 “ 
 
 — 220° 
 
 15 “ 
 
 — 140° 
 
 39 “ 
 
 Low temiieratiires are ascertained by means of a hydrogen or helium 
 
HALOGEN GROUP— CHLORINE. 
 
 49 
 
 thermometer, a thermo-electric element composed of copper and German 
 silver, or a platinum resistance thermometer. 
 
 The critical temperature, pressure and volume can be determined not only experiment- 
 ally, but also may be deduced from the variations of the gases from the laws of Boyle and 
 Gay-Lussac, by a theory developed by van der Waals (1873) (compare Air : measurement 
 of gases). Students of chemistry desiring more detailed information on this subject 
 are directed to the works of Ostwald : Grundriss der allgemeinen Chemie, 2 Aufl. 
 1890, and Lothar Meyer : Chemical Theories, 1893, and also Ostwald’ s Lehrbuch, 
 vol. I, 289 (1891). 
 
 ,, ■ ' ' 
 
 ■ "’halogen group. 
 
 Chlorine, bromine, iodine, and fluorine constitute this group. These 
 elements show a similar chemical deportment. They are termed halogens 
 or salt producers, because by their direct union with the metals salt-like 
 derivatives result. Chlorine, bromine, and iodine will be discussed first 
 because of their great importance. 
 
 1 . CHLORINE. 
 
 Atom ; Cl = 35.45. Molecule : CL = 70.90. 
 
 It does not occur free in nature. As it acts very energetically upon 
 most bodies containing hydrogen, especially those of an organic nature, 
 to form new derivatives with them or with their constituents, chlorine can 
 only be obtained free in the transitional state. Its most important deriv- 
 ative is sodium chloride, or rock salt, which is composed of chlorine and 
 sodium. The Swedish chemist, Scheele, discovered chlorine in 1774. 
 It was regarded as a compound body, until its elementary character was 
 established by Gay-Lussac and Thenard in France (1809), and by Davy 
 in England (1810). 
 
 Preparation . — To obtain free chlorine, heat a mixture of black oxide 
 of manganese (MnOa) and hydrochloric acid in a flask (Fig. 31), pro- 
 vided with a so-called Welter safety-tube to equalize the gas pressure. 
 The escaping gaseous chlorine is washed and freed from acid that is 
 carried along mechanically by passing it through water in a three-necked 
 Woulff bottle, and then collecting it over water. The reaction which 
 occurs above is indicated in the following equation : 
 
 MnO^ 4 - 4 HCI = MnCh + 21120 -f Clj. 
 
 The manganous chloride formed dissolves in the water. 
 
 At low temperatures the mangane.se peroxide dissolves in the hydrochloric acid without 
 the evolution of chlorine, the solution being brown in color : 
 
 Mn02 + 4IICI = MnCq + 2II2O. 
 
 Chlorine is expelled on the aj)plication of heat : 
 
 MnCb = MnCl2 -f Cl^. 
 
50 
 
 INOKflANFC CHEMISIRV. 
 
 The evolution of tlie chlorine proceeds more regularly if a mixture of manganese oxide, 
 sodium chloride, and suliduiric acid is employed. 
 
 The hydrochloric acid is formed from the last two (together with sodium acid sul- 
 phate) : 
 
 NaCl f II, SO, zzzz IICl + NallSO,, 
 
 and then acts upon the manganese peroxide. It is advisable to use 5 parts of manga- 
 nese peroxide (90 j^er cent. ), ll parts of sodium chloride, atid 14 parts of concentrated 
 sulphuric acid diluted with 7.5 parts of water [Klason, l>er. 23 (1890), 330]. 
 
 'File second method is more advantageous for laboratory jnirposes ; the first (action of 
 hydrochloric acid upon manganese dioxide), however, is jireferred in jiractice, as it is 
 cheaper. 
 
 d'he resulting manganous chloride (MnCl,) is converted by the proctss of Weldon 
 into manganese peroxide (see this). 
 
 Other technical methods for the preparation of chlorine are based upon the liberation 
 of chlorine from hydrochloric acid or metallic chlorides (particularly calcium chloride and 
 magnesium chloride) by the oxygen of the air at elevated temperatures. Ily-products of 
 
 Fig. 31. 
 
 the soda industry, mother liquors of sea salt and tailings from the .Stassfurth mines, all of 
 which contain chlorine, are utilized in this way. 
 
 1. Deacon- JIurte)' Process: Hydrochloric acid mixed with air is conducted over 
 jiorous substances saturated with copper chloride or sulphate and heated to 370-400°. 
 
 2. Weldo)i- Pechiney Process : Magnesium chloride or oxychloride is heated to 1000° 
 in an air current : MgCl, | O = MgO -|- Cl, ; the magnesia is reconverted into chlo- 
 ride by solution in hydrochloric acid. 
 
 3. Solvay Process: Calcium chloride or magnesium chloride, mixed with silica or 
 clay, is heated in an air current: CaCl, + SiO, I O =CaOSiO., 4- Cl.,. 
 
 4. Pcychlcr-de Wilde Mel hod : A mixture of magnesium chloride, manganous chlo- 
 ride and manganc.se sulphate is exposed to air heated to about 520° until tlie evolution 
 of chlorine ceases, when it is restored by hydrochloric acid to its original condition and 
 chlorine again liberated. 
 
 t;. Mond Method : Ammonium chloride vapor is allowed to act upon magne.sium 
 oxide or olliei' metallic oxides .and the ])roducts are ammonia and metallic chlorides, 
 'file latter aie heated to a high tempc'iatuie in air when chlorine and oxiiles result. 
 
CHLORINE. 
 
 51 
 
 The electrolytic decomposition of salt solutions is especially important ; 
 indeed it is likely to supplant the other technical i)rocesses in the near 
 future. The current decomi)oses sodium chloride (and other metallic 
 chlorides) into its components, so that chlorine separates at the positive 
 and sodium at the negative pole. The sodium at once acts upon water 
 forming sodium hydroxide and hydrogen. This universally important 
 process will be again referred to in connection with soda, potassium chlo- 
 rate and bleaching lime.* 
 
 An excellent laboratory method for the preparation of chlorine con- 
 sists in allowing dilute hydrochloric acid to act upon bleaching lime (see 
 this). The latter is previously mixed with burnt gypsum ()^ part), and a 
 little water added, when the mass can be formed into cubes or stout sticks 
 which are introduced into a Kipp generator (p. 42). 
 
 Very pure chlorine may be obtained by digesting potassium bichro- 
 mate with concentrated hydrochloric acid : 
 
 KaCr^O^ + 14HCI = Cr2Cl6 + 2KCI -f 7H2O + 3CI2. 
 
 As chlorine dissolves readily in cold water it is advisable to collect it over warm water. 
 It cannot be collected over mercury, as it readily combines with the latter. When perfectly 
 dry chlorine is sought, conduct the liberated gas through Woulff bottles containing con- 
 centrated sulphuric acid, to absorb the moisture, then collect in an empty upright flask. 
 As chlorine is so much heavier than air it will displace the latter. 
 
 Physical Properties. — Chlorine is a yellowish-green gas- (hence its name 
 irom yXwpoq), with a penetrating, suffocating odor. Its specific gravity 
 compared with oxygen (02 = 32) is 70.9 ; 
 with air (— i) it is 2.45. A liter of chlo- 
 rine at 0° and 760 mm. pressure at the sea- 
 level in the latitude of 45° weighs 3.167 
 grams. At 15° C., and a pressure of 57 at- 
 mospheres (at — 40° C., under the ordinary 
 pressure) it condenses to a yellow liquid, 
 boiling at — 33 6°. By pressure and cold, 
 liquid chlorine has been made applicable 
 in technical operations (see Knietsch, Ann. 
 
 Chem. 259 (1890), 100). The specific 
 gravity of liquid chlorine at its boiling 
 point is 1.5575. At very low tempera- 
 tures it solidifies to a yellow crystalline 
 mass, which remelts at — 102°. The 
 critical temperature of chlorine is 146°, and its critical pressure 93.5 
 atmospheres. It may be condensed in the following manner as demon- 
 strated by Faraday: Take a bent glass tube (Fig. 32), introduce into 
 the leg closed at one end crystals of chlorine hydrate (-CI2 + SHjO, see 
 p. 52), then seal the o])en end. The limb containing the comi)ound is 
 
 Fig. 32. 
 
 ■^‘Compare N. Caro, Dar.stellung von Cblor unci Salzsaure, Tlerlin, 1893; also L. 
 Mond, flesdiichtliclier Ue])erblick der verscbiedenen Arten der Cldordarstelliing, Cliem- 
 iker-Zeitung, 1896, 928, and R. Masenclever, Die Entwickelung der Sodafabrikation, 
 Ber. 29 (1896), III, 2861. 
 
52 
 
 INORGANIC CHEMISTRY. 
 
 placed in a water-bath; the other is cooled in snow. Uj)on heating the 
 water a little above 30° the chlorine hydrate is decomposed into water 
 and chlorine gas, which condenses to a liquid in the cooled limb. On 
 reversing the position of the limbs and cooling the one previously 
 warmed, the chlorine distils back and is reabsorbed by the water. Ohar- 
 coal saturated with chlorine may be substituted for the chlorine hydrate. 
 One volume of charcoal takes up 200 volumes of chlorine, which are 
 disengaged again on heating. 
 
 One volume of water, at 20° C., absorbs 2 volumes of chlorine; at 
 10° C., 2.5 volumes. The aqueous solution is known as chlorine water 
 {aqua chlorata), and possesses almost all the ])roperties of the free gas; it 
 is therefore frequently employed for laboratory uses as a substitute for 
 chlorine. 
 
 Recent investigations show that chlorine water, by virtue of the action of the chlorine 
 upon water, contains hydrochloric and hypochlorous acids, which in turn react upon one 
 another with the production of chlorine and water, so that in accordance with the strength 
 of the chlorine water and external prevailing conditions equilibrium occurs among these 
 four substances (p. 29). This may be expressed as follows : 
 
 Cb + H-P HCl + HCIO. 
 
 The yellow, scale-like crystals of chlorine hydrate (Cb -j- 8H2O) sepa- 
 rate when water saturated with the gas is cooled below 0°. This com- 
 pound is regarded as one of chlorine with water. At ordinary tempera- 
 tures it decomposes into water and chlorine. 
 
 Chemical Properties . — When thin sheet copper (false gold leaf), or, 
 better, pulverized antimony or arsenic, is thrown into a vessel filled with 
 dry chlorine, it will burn with a bright light; a piece of phosphorus will 
 also inflame in an atmosphere of the gas. Perfectly dry chlorine does 
 not manifest this energetic combination-tendency. 
 
 Chlorine unites just as energetically with hydrogen. A mixture of 
 equal volumes of the gases obtained by the electrolysis of hydrochloric acid 
 combines in direct sunlight with violent explosion. In diffused sun- 
 light the action is only gradual ; in the dark it does not occur if perfectly 
 dry gases are employed (Baker). The course of this reaction and the 
 laws of the chemical action of light have been ably discussed in a series 
 of celebrated papers by Bunsen and Roscoe [Pogg. Ann. d. Physik, Bd. 
 100, lOi, 108 (1857-1859)]. Chlorine also manifests great affinity for 
 hydrogen derivatives, most of them being so decomposed by the chlorine 
 that hydrogen is removed from them, and hydrochloric acid is formed. 
 Thus water is decomposed, especially in sunlight, by chlorine into hydro- 
 chloric acid and oxygen : 
 
 11,0 + CI2 2HCI + O. 
 
 If a glass cylinder be filled with and inverted over chlorine water and 
 exj)osed to direct sunlight, a gas will be evolved, and will collect in the 
 ui)i)er |)ortion of the vessel ; this isoxygen. In diffused light the decom- 
 position will not be so rapid ; it is hastened by heat. Chlorine oxygen 
 acids, IK'IO, IICK)., fsec above), are also jiroduced. 
 
 Chlorine alters the hydrocarbons, in that it abstracts hydrogen. The 
 
BROMINE. 
 
 53 
 
 reaction is sometimes so violent that carbon is separated in a free condi- 
 tion. A piece of tissue paper saturated with freshly distilled turpentine 
 oil, and introduced into a dry chlorine atmosphere, is immediately car- 
 bonized. An ignited wax taper immersed in chlorine burns with a smoky 
 flame, with separation of carbon. 
 
 The organic dyestuffs (containing carbon and hydrogen) are decolor- 
 ized by moist chlorine gas. The same occurs with the dark-blue solutions 
 of indigo and litmus ; colored flowers are rapidly bleached by it. On this 
 princii)le depends the application of chlorine in bleaching, and the de- 
 struction of decaying matter and miasmata in chlorine disinfection are 
 dependent upon this reaction (see Bleaching Lime). 
 
 The bleaching action of chlorine is mostly influenced by the presence of water. It 
 probably depends on the oxidizing action of the oxygen liberated by the chlorine or upon 
 the chlorine oxy-acids found when chlorine and water react (see p. 52). This property, 
 free oxygen does not possess ; it does, however, very probably belong to that which is in 
 the act of fonning, — of becoming free. We will learn, later, that many other elements, 
 at the moment of their birth {in statu nascendi), act more energetically than when in 
 the free condition. 
 
 2. BROMINE. 
 
 Atom : Br = 79.96. Molecule : Br2 = 159.92. 
 
 Bromine, the perfect analogue of chlorine, was discovered by Balard, 
 in 1826. It is not found free in nature for the same reasons which 
 were given under chlorine. It occurs in sea-water as sodium bromide, 
 accompanied by sodium chloride, but in much smaller quantity than the 
 latter (especially in the water of the Dead Sea), and in many salt springs, 
 as at Kreuznach and in Hall. When sea-water or other salt water is 
 evaporated, sodium chloride first separates; in the mother liquor, among 
 other soluble salts, are found sodium and magnesium bromides. Bromine 
 is found in greatest abundance in the upper layers of the rock-salt deposits 
 of Stassfurth, near Magdeburg, where it exists in the form of bromides 
 together with other salts. At present, large quantities of bromine are 
 obtained in America. The method of its preparation is similar to that 
 employed under chlorine. A mixture of manganese dioxide and sodium 
 bromide is warmed with sulphuric acid : 
 
 MnOg + 2NaBr -f 2H2SO^ = MnS04 -|- Na^SO^ -|- Br2 -f 2H2O. 
 
 The bromine condenses in the well-cooled receivers. When free chlo- 
 rine is conducted into an aqueous solution of sodium bromide, bromine 
 separates. This is the method adopted at Stassfurth and in America in 
 preparing bromine from the magnesium bromide saline liquors : 
 
 MgBr2 -f CI 2 = MgCb -f Br2. 
 
 Bromine is a heavy, reddish-brown liquid, with an exceedingly pene- 
 trating, chlorine-like odor (hence the name Bromine, from [ipcbij.oc;, stench). 
 When bromine is strongly cooled it solidifies to a dark-brown mass, of 
 delicate needles with a slightly black metallic luster, which remelt at 
 — 7.3°. Liquid bromine at 0° has the sjiecific gravity 3. 18 (water = i) ; 
 it is very volatile, forming dark-brown vapors at the ordinary tempera- 
 
54 
 
 INORGANIC CHEMISTRY. 
 
 lure, and boils at about 6o°, clianging at the same time into a dark 
 brown-red va^xir. Its density eijiials 159.9 (oxygen 32), (jr 5.52 (air 
 = 1). 
 
 One part of bromine dissolves in about 35 ])arts of water of medium 
 temjierature; it is therefore more soluble in water than chlorine. Cooled 
 below 4° C., the hydrate ( I Ir.^ -f roll./)) crystallizes out: this is analogous 
 to the chlorine hydrate. It is decomi)osed at 6.2°. J>romine dissolves 
 with ease in alcohol, and esi)ecially in ether, chloroform and bisulphide. 
 
 Chemically, bromine is extremely like chlorine, combining directly 
 with most metals to form bromides; but it possesses a weaker affinity 
 than chlorine, and is liberated by the latter from its compounds. With 
 hydrogen it only combines on warming, not in sunlight. Ujion hydro- 
 carbons it acts like chlorine, withdrawing hydrogen from them, llromine 
 water gives starch-paste an orange color. 
 
 Solid bromine is a mixture of licpiid bromine with silicious earth. It 
 occurs in trade pressed into cubes or sticks and is used for disinfecting 
 purposes. Bromine and its salts are api)lied in medicine, in photography 
 and in the manufacture of aniline colors. 
 
 3. IODINE. . \ 
 
 Atom : I = 126.85. Molecule : I2 = 253.70. 
 
 Iodine, as well as bromine, occurs in combination with sodium, in sea- 
 water and some mineral springs, especially at Hall, in Austria, and the 
 Adelheit spring in Bavaria. In these springs the iodine can easily be de- 
 tected ; in sea-water it is, however, only present in such minute quantity 
 that its separation, practically, is disadvantageous. Sea algae absorb iodine 
 compounds from the water, and these are then thrown by the tide on various 
 coasts, where they are burned, yielding an ash (known as kelp in Scotland, 
 as varec in Normandy) which is the jirincipal source for the manufacture 
 of iodine. It was in this ash that the element was accidentally discovered 
 l)y Courtois in 1811; in 1813, it was investigated by Davy and Gay- 
 Lussac, and its elementary character established. To obtain the iodine, 
 the ash is treated with water, the solution concentrated, and separated 
 from the chlorides of sodium and potassium which have crystallized out, 
 the final liquors, in which the readily soluble iodides have accumulated, 
 being then distilled with manganese dioxide and sulphuric acid. Iodine, 
 therefore, is set free from its compounds in the same manner as chlorine 
 and bromine. It is more convenient, however, to pass chlorine (or 
 better, nitrous acid) through a solution of the iodides, when all the iodine 
 will separate : 
 
 KI 4- Cl KCl 4- 1 , 
 
 or the iodide solution is digested with ferric chloride, when the latter is 
 reduced to ferrous chloride : 
 
 Nal 1- FeCly =1 -- 1 + NaCl |- FeCl.,. 
 
 The grayish-black jiowder thus liberated is collected on a filter, dried, 
 and then sublimed. 
 
IODINE. 
 
 55 
 
 In recent years the greatest quantities of iodine have ])een obtained 
 from the mother liquors of crude Chile saltpeter (NaNO;,). The iodine is 
 present in this salt as sodium iodate (NalOa), from which it is set free by 
 nitrous or sulphurous acid : 
 
 aNalOg + SH.SOg = 2I -f 2NaHS04 + SH.SO^ + Hp. 
 
 An interesting occurrence of iodine is that observed by Baumann in 
 1895 thyroid (thyroidea, from Oopeoq, shield, and e'ldof;, form) gland, 
 
 the active constituent of which, containing 9 per cent, of iodine, is mixed 
 with milk-sugar and sold in the markets under the name thyroidin. 
 
 Iodine is a gray-black solid, subliming in large rhombic crystals, pos- 
 sessing strong metallic luster. It has a peculiar odor, reminding one 
 somewhat of that of chlorine; it stains the skin brown, and is corrosive, 
 although not as strongly so as bromine. Its specific gravity is 4.95 at 
 15°. It fuses at 114° to a dark-brown liquid, and boils near 184.3°, 
 passing at the same time into a dark-violet vapor (hence the name Iodine, 
 suggested by Gay-Lussac, from loetdrjq, violet-like). 
 
 The vapor density of iodine equals 8.7 up to 600° C. (air= i) or 254 (O2— 32), cor- 
 responding to the molecular weight I2 = 253.7. Above 600° the vapor density gradually 
 diminishes, and at about 1500° it is only half the original. This is explained by the 
 gradual decomposition (see Dissociation of Water) of the normal diatomic molecule L 
 into the free atoms I T I. In like manner the bromine molecules Br2 suffer a separation 
 into the free atoms. The dissociation of bromine vapor (diluted with ii volumes of nitro- 
 gen) commences at about 1000° and is complete at 1600°. The vapor density of chlorine 
 is still normal at 1200°, and it is only at 1400° that it sustains a slight diminution. Oxy- 
 gen and nitrogen on the contrary show no alteration in their vapor density even at 1690° 
 (C. Danger and V. Meyer). 
 
 Iodine is very slightly soluble in water (i : 3600), more readily in 
 alcohol (^Tinctura iodi'), very easily in aqueous potassium iodide, in ether, 
 in chloroform and in carbon bisulphide (10 c.c. of the latter will dissolve 
 1.85 grams of iodine at 18°), the last two assuming a deep red-violet color 
 in consequence. It affords a particularly beautiful crystallization, consist- 
 ing of forms of the rhombic system, when it separates from a solution of 
 glacial acetic acid. 
 
 In chemical deportment iodine closely resembles bromine and chlorine ; 
 it possesses, however, weaker affinities and for this reason is liberated from 
 its compounds by those elements. With the metals it usually combines 
 only when warmed ; with hydrogen it does not combine directly, and it 
 does not remove it from its carbon compounds. 
 
 The deep-blue color it imparts to starch is characteristic of iodine. 
 On adding starch-])aste to the solution of an iodide, and following this 
 with a few drops of chlorine water, the paste will immediately be colored 
 a dark blue by the separated iodine. This reaction serves to detect the 
 smallest quantity of it. 
 
 Iodine is largely employed in medicine, photography, and in the 
 preparation of aniline colors. 
 
56 
 
 INORGANIC CHKMIS'IKY. 
 
 4. FLUORINE. 
 
 Atom ; r'l = 19. Molecule : l''l2 = 38. 
 
 Fluorine has such great affinity for nearly all substances, that des])ite 
 numerous efforts it has only recently been possible to obtain it in a free 
 condition. It was Moissan (1886) who electrolyzed anhydrous, strongly 
 chilled hydrofluoric acid to which some sodium lluoride had been added, 
 and obtained fluorine as a slightly greenish-yellow colored gas. The de- 
 composing vessel was a j)latinum tube with sto])pers of fluorspar. 'J'he 
 ])ositive pole, at which the fluorine ai)peared, consisted of an alloy of 
 platinum and iridium. Dewar and Moissan have very recently succeeded 
 in liquefying fluorine at — 187° by the use of boiling oxygen as the refrig- 
 erating substance. Under these conditions it loses its chemical activity 
 almost completely, does not attack glass, iodine, sulphur, or metals. 
 It is only the affinity for hydrogen which remains; at least benzene is 
 entirely decomposed with production of flame by fluorine. 
 
 Fluorspar (CaFla) is the most important fluorine derivative. It is most 
 frequently employed in the preparation of other fluorine compounds. 
 
 Fluorine combines with hydrogen, iodine, sulphur, silicon, boron, pow- 
 dered arsenic and antimony, finely divided iron and manganese even in 
 the dark ; organic substances, like oil of turpentine, alcohol and cork 
 burn in the gas. The metals, with the exception of gold and platinum, 
 are energetically attacked by fluorine. The latter decomposes water with 
 the production of hydrofluoric acid and ozonized oxygen, and liberates 
 chlorine, bromine and iodine from their metallic derivatives. The 
 density of fluorine referred to hydrogen equals 1.32 (1.26 found). 
 
 These four elements, fluorine, chlorine, bromine and iodine, exhibit 
 gradual differences in their properties; and, what is remarkable, this gra- 
 dation stands in direct relation to the specific gravity of the elements in 
 the state of gas or vapor. 
 
 FI Cl Br I 
 
 Specific gravity, 19 35-45 79 - 9^ 126.85 
 
 A simultaneous condensation of matter occurs with the increase of spe- 
 cific gravity. This expresses itself in the diminished volatility. Fluorine 
 is a gas down to — 187° ; chlorine can readily be condensed to a liquid ; 
 bromide is a liquid at ordinary temperatures, and iodine is a solid. 
 Other ])hysical properties, as seen in the following table, are also in 
 accord with the preceding: 
 
 
 Fluorine. 
 
 Chlorine. 
 
 Bromine. 
 
 Iodine. 
 
 Fusing point, 
 
 
 — 102° 
 
 — 72° 
 
 4 114° 
 
 Foiling ]X)inl, 
 
 -187° (?) 
 
 - 33° 
 
 + 60° 
 
 4184° 
 
 Specific gravity in licpiid .... 
 
 1. 14 
 
 1.47 
 
 3.18 
 
 
 or .solid condition, .... 
 
 
 
 
 4-95 
 
 Color, 
 
 Green - yel- 
 low 
 
 Yellow- 
 
 green 
 
 Frown 
 
 Flack- 
 
 violet 
 
HYDROGEN CHLORIDE. 
 
 57 
 
 Just such a gradation, as we have seen, is observed in the chenaical 
 affinities of these four elements for the metals and hydrogen ; fluorine is 
 the most energetic, iodine the least. Iodine is therefore separated from 
 its soluble metallic and hydrogen compounds by the other three, bromine 
 by chlorine and fluorine, and chlorine by fluorine (p. 56). We shall dis- 
 cover, later, that the halogens are displaced in exactly the reverse order 
 of their oxygen compounds ; that iodine has the greatest and chlorine the 
 feeblest affinity for oxygen. Oxygen derivatives of fluorine are not 
 known. 
 
 COMPOUNDS OF THE HALOGENS WITH 
 HYDROGEN. 
 
 The halogens form gaseous acids, readily soluble in water, with hydro- 
 gen : the halogen hydrides or haloid acids. But one such compound is 
 known for each halogen. 
 
 1. HYDROGEN CHLORIDE. 
 
 HCl = 3646. 
 
 The direct union of chlorine with hydrogen takes place through the 
 agency of heat, and by the action of direct sunlight or other chemically 
 active rays (magnesium light ) ; in diffused light the action is only 
 gradual, and does not occur at all in the dark. When both gases are 
 perfectly dry they do not react in direct sunlight. On introducing a 
 flame of hydrogen ignited in the air into a cylinder filled with chlorine 
 it will continue to burn in the latter with the production of hydrogen 
 chloride. The opposite, the combustion of chlorine in an atmosphere of 
 hydrogen, may be shown easily by the following experiment : An 
 inverted cylinder is filled with hydrogen by displacement, the gas is 
 ignited at the mouth, and a tube immediately introduced which will con- 
 duct dry chlorine into the cylinder. The burning hydrogen will inflame 
 the chlorine, which will continue to burn in the former. From these 
 experiments, we perceive that combustion is a phenomenon which ac- 
 companies a chemical change ; in this instance the union of hydrogen 
 with chlorine; if hydrogen is combustible in chlorine (or air), so, 
 inversely, is chlorine (or air) combustible in hydrogen for the same 
 reason. By the term combustion, in chemistry, is understood every 
 chemical union of a body with a gas, which is accompanied by the phe- 
 nomenon of light. 
 
 A mixture of ecjual volumes of chlorine and hydrogen is called chlor- 
 detonating gas ; it explodes with very great violence under the conditions 
 given above for the union of the gases. The product is gaseous hydrogen 
 chloride. 
 
 The formation of the latter compound succeeds best by allowing sul- 
 
58 
 
 INORGANIC CHEMISTRY. 
 
 l)hiiric acid to act upon sodium chloride when solid sodium bisulphate 
 and hydrogen chloride gas will result : 
 
 NaCl + 1 1 , SO, -r NallSO, + IICl. 
 
 Pour over 5 parts sodium cliloride, 9 })arts of concentrated sulphuric acid, diluted with 
 water (2 parts), and gently warm the tlask on a sand-bath (Fig. 33). 'I'he escaping 
 hydrogen chloride is conducted through a Woulff bottle containing .sulphuric acid or 
 through the cylinder ll ((died with pumice-stone saturated with sul[)huric acid), intended 
 to free it from all moisture, and atlerward collected over mercury in the cylinder C. 
 
 At a red heat, the acid sodium sulphate reacts anew upon the sodium chloride with the 
 formation of neutral sodium sulphate and hydrochloric acid: 
 
 NallSO, -f NaCl = IICl -f Na,SO,. 
 
 These are the reactions by which hydrochloric acid, in the preparation of sodium sulphate 
 by the Le Blanc soda process, is technically made. As this method, however, is being 
 
 Fig. 33. 
 
 replaced by the electrolytic method, in which not hydrochloric acid, but chlorine is evolved, 
 it seems very important that the latter should be converted by some convenient procedure 
 into hydrochloric acid, always of the greatest importance technically. Lorenz contends 
 that this may be accomplished by conducting the chlorine together with steam over ignited 
 coke, when the chief products will be hydrochloric acid and carbon dioxide (CO.^) : 
 
 2CI, -f 2ll,0 -(- C == 4HCI -f CO.,. 
 
 [Comj)are Per. 30 (1897), l, 347]. 
 
 'I’lu! suggestion of Davy to allow .sulphuric acid to act upon pieces of ammonium 
 chloride will give a regular current of hydrogen chloride: 
 
 Nil, Cl + II.,SO,, = (NII,)IISO, f IICl. 
 
 A Norblad generator as modilied by Kreussler (Fig. 34) answers well for this purpose. 
 
HYDROGEN CHLORIDE. 
 
 59 
 
 Physical Properties . — Hydrogen chloride is a colorless gas, with a 
 suffocating odor. In moist air it forms dense clouds as it combines with 
 the aqueous vai)or to form hydrochloric acid. Its critical temperature is 
 about -1-52.3°, and the critical pressure 86 atmospheres, /. e., for its con- 
 densation at the tenqjerature just given it requires a pressure of 86 atmos- 
 pheres. There is no i^ressure which will condense it 
 above this temperature. Liquid hydrogen chloride is 
 colorless, has a specific gravity at 15° of o 83, and freezes 
 at — 115-7° to a white crystalline mass, which begins 
 to melt at — 112.5°. boils at — 80.3° under the 
 ordinary pressure. 
 
 The specific gravity (density) of the gas is 36.46 (O2 = 
 
 32), or 1.26 (air= i). One liter of it weighs 1.6285 grams 
 at 0°, 760 mm. pressure, in latitude 45° and at sea-level. 
 
 Hydrogen chloride possesses an acid taste, and colors 
 in the presence of water blue litmus-paper red ; it is, 
 therefore, an acid, and has received the name hydro- 
 chloric acid gas. It dissolves very readily in water, and 
 on that account cannot be collected over it. One volume 
 of water at 0° C. and 760 mm. pressure dissolves 505 
 volumes, and at ordinary temperatures about 450 volumes 
 of the gas. At higher pressure water dissolves more 
 hydrochloric acid, and at low pressures less (compare 
 Carbon Dioxide, and also Solutions). This great solu- 
 bility is very nicely illustrated by filling a long glass 
 cylinder with the gas and then just dipping its open end 
 into water ; the latter rushes up into the vessel rapidly (as 
 into a vacuum), as it quickly absorbs the gas. The aque- 
 ous solution of hydrogen chloride is commonly known 
 as muriatic or hydrochloric acid (^Acidiim hydrochIo 7 'icu 7 )i). For its prep- 
 aration the gas is passed through a series of Woulff bottles (Fig. 35) con- 
 taining water. The small bottle B, in which there is but little water, 
 serves to wash the gas — free it of any mechanically admixed sulphuric 
 acid. The same apparatus may be employed in the manufacture of chlor- 
 ine water, and is generally used in the saturation of liquids with gases. 
 
 A solution saturated at 15° C., contains about 42.9 per cent, hydrogen 
 chloride, has a specific gravity of 1.212, and fumes in the air. On the 
 a])plication of heat, the gas again escapes, and the temperature of the 
 liquid rises to 110° C., when a liquid distils over, containing 20.24 ])ei- 
 cent, of hydrogen chloride, having a specific gravity of 1. 104 and corre 
 sponds a])proximately to the formula HCl -f- 8H.2O. The composition of 
 the distillate varies somewhat with the pressure. A dilute acid, upon 
 distillation, loses water, until finally that boiling at 110° C. ])asses over. 
 On conducting hydrogen chloride into concentrated hydrochloric acid 
 cooled to — 22°, crystals of the formula HCl -f 2H.2O separate; these 
 fuse at — 18° and then decompose. 
 
 Hydrochloric acid finds an extensive industrial application, and is 
 obtained in large quantities, as a by-product, in the soda manufacture 
 (Le Blanc process). 
 
 Fig. 34. 
 
6o 
 
 INORGANIC CHEMISTRY. 
 
 CJiemical Properiies of Jlydroi^cn Chloride. — Acids — Bases — Salts . — 
 Hydrogen chloride is a very stable compound ; it is only l)eyond 1500'’ 
 that it sustains a partial decomposition (see Dissociation of Water). Its 
 com])Osition is readily established fpiantitatively in the following way : 
 Pass hydrochloric acid gas over a piece of sodium or iiotassium heated in 
 a glass tube, and hydrogen will escaiie from the latter: 2Na -f- 21101 — 
 2NaCl -f- Il'i- If, on the other hand, manganese ])eroxide be heated in 
 it, chlorine will be disengaged : MnO.^ -[- 41101 = IVInOlj -f 2H2()-}-Ol2. 
 If the electric current be allowed to act upon an a(pieous solution of 
 hydrochloric acid the latter will be decomposed so that chlorine se])arates 
 at the electro-positive and hydrogen at the electro-negative j)ole (j). 77). 
 Hydrogen chloride, as well as its solution, possesses all the properties of 
 
 acids, and can well figure as a prototype of these; it tastes intensely sour, 
 reddens blue litmus-pa})er, and saturates the bases (oxides and hydrox- 
 ides, i. e., which bodies impart a blue color to red litmus-paper) form- 
 ing chlorides. There are some bases, the alkalies and alkaline earths, 
 which are soluble in water. These solutions react basic, alkaline, i. e., 
 red litmus is colored blue by them. If we add hydrochloric acid to a 
 solution of a base, e. sodium hydroxide, until the reaction is neutral, 
 we will obtain (besides water) a neutral compound — sodium chloride, 
 which remains in a crystalline form when the solution is evaporated: 
 
 NaOir I IICI NaCI 1 II 2 O. 
 
 Sodium Sodium 
 
 hydroxide. chloride. 
 
HYDROGEN BROMIDE. 
 
 6l 
 
 Hydrogen bromide, iodide and fluoride deport themselves similarly to 
 hydrogen chloride. These halogen compounds of hydrogen are termed 
 haloid acids, to distinguish them from those which, in addition to hydrogen, 
 contain oxygen, hence called oxygen acids. The latter conduct themselves 
 like the former, and saturate bases, forming salts and water : 
 
 KOH + IINO3 ^ + H^O. 
 
 Potassium Nitric Potassium Water, 
 
 hydroxide, acid. nitrate. 
 
 In the same manner the acids act upon the basic oxides, to form salt; 
 and water : 
 
 ZnO + 2HCI == ZnCL 4 - H^O. 
 Zinc Zinc 
 
 oxide. chloride. 
 
 ZnO + 2HNO3 = ZnfNOg)^ + H^O. 
 Zinc Zinc 
 
 oxide. nitrate. 
 
 Usually when acids act upon metals, the hydrogen of the former is 
 directly displaced ; salts and free hydrogen are produced. Thus, by the 
 action of hydrochloric acid upon sodium, its chloride and hydrogen 
 result : 
 
 HCl + Na = NaCl + H ; 
 
 and when zinc and hydrochloric acid react, zinc chloride and water 
 (see p. 41) : 
 
 2HCI + Zn = ZnCb + H2. 
 
 From the examples cited it is manifest that acids are hydrogen com- 
 pounds which yield salts, by the replacement of their hydrogen by metals 
 (by the action of metallic oxides, hydroxides, and by the free metals). 
 The metallic oxides and hydroxides like* sodium hydroxide, capable of 
 forming water and salts by the saturation of acids, are called bases. 
 Finally, by the term salts, we understand such compounds as are analogous 
 to sodium chloride, and are formed by the mutual action of bases and 
 acids with the exit of water. Salts are distinguished as haloid salts a.nd 
 oxygen salts. The first have no oxygen, and arise in the direct union of 
 the halogens with the metals: 
 
 Na+ Cl =NaCl. 
 
 Zn -j- CI2 = ZnCI^. 
 
 2. HYDROGEN BROMIDE. 
 
 HBr = 80.97. 
 
 Hydrogen bromide is perfectly similar to the corresponding chlorine 
 compound. As there is but slight affinity between bromine and hydrogen 
 their direct union will only occur at a red heat or in the presence of 
 ])latinum sponge (see p. 46). Like hydrogen chloride, hydrogen bro- 
 mide can be obtained by the action of some acids, e. g , phosphoric acid, 
 upon bromides; concentrated sulphuric acid would not answer as the re- 
 
62 
 
 INORHANIC CHEMIS'I’RY. 
 
 suiting hydrogen bromide is apin i)artly decomposed by it. Ordinarily 
 it is prepared by the action of phosi)horus tribromide upon water ; 
 
 Vlh, + 3II/) ll3P()3 -p 3iip,r. 
 
 PIiosi)li()nis I’hospliorous 
 
 tribroniide. acid. 
 
 Place water (i part) in a flask (Fig. 36), gradually admit through the 
 funnel, supplied with a stop-cock, the licjuid j)h()S[)horus tribromide (3 
 parts), and warm gently, 'bhe escaping gas is collected over mercury or 
 conducted into water. To free it perfectly from accom])anying ])hos- 
 ])horus bromide vapors it is passed through water (the U-shai)ed tube in 
 Fig. 36 contains pieces of pumice-stone or glass beads, which are moist- 
 ened with water). 
 
 Instead of employing prepared phosidiorus bromide, we may let bromine vapors act 
 
 u[)on (red) phosphorus, d'hismay be done 
 by ])ouring water (2 parts) over the phos- 
 phorus ])laced in a flask; bromine ( 10 parts) 
 is added gradually while cooling and heat is 
 then applied. To free the hydrogen bro- 
 mide gas from the bromitie carried along 
 mechanically, conduct it through a tube 
 containing glass wool and moist red phos- 
 phorus. 
 
 Claseous hydrobromic acid can also be 
 prepared by allowing bromine to act upon 
 crude anthracene. The resulting hydro- 
 bromic acid gas is freed from the accom- 
 panying bromine by pa.ssing it through a 
 tube filled with anthracene. 
 
 To obtain an aqueous solution of the 
 gas, pour 15 parts of water over i part of 
 red phosphorus, and then add bromine (10 
 parts) drop by drop. Finally the .solution 
 is heated, filtered, and distilled. Bromides 
 (sodium bromide, potassium bromide) yield 
 hydrogen bromide by distillation with dilute 
 Fig. 36. sulphuric acid in the presence of phosphorus. 
 
 Hydrogen bromide is a colorless gas, fuming strongly in the air. 
 Under great pressure it is condensed to a liquid, solidifying at — 120°, 
 melting at — 87°, and boiling at — 73°. Its density is 80.97 (O^ = 32) or 
 2.79 (air = i). 
 
 In water the gas is very readily soluble, its solution saturated at 0° 
 having a siiecific gravity of 1.78, and containing 82 ]ier cent, of hydro- 
 gen bromide. Its composition closely approximates the formula HBr -|- 
 1 1/) ; at 15° it contains 49.8 per cent, of acid and has the specific gravity 
 1.5 15. At 125° a solution distils over, containing 48.2 per cent, of hy- 
 drogen bromide; its composition corresponds very nearly to the formula 
 1 1 l>r 5 1 ; its specific gravity is 1.49 at 14° C. 
 
 On conducting hydrogen bromide into a solution of the same cooled 
 to — 20°, crystals of the formula Hllr -|- 2ll./) separate and melt at 
 — i Chemically, hydrogen bromide is the jierfect analogue of hydro- 
 gen chloride ; it is, however, less stable, and suffers a partial decomposi- 
 tion at 800° C. 
 
HYDROGEN IODIDE. 
 
 63 
 
 3. HYDROGEN IODIDE. 
 
 HI = 127.86. 
 
 The attraction of iodine for hydrogen i.s very slight. The two elements 
 combine at higher temperatures, between 400 and 500°, very incom- 
 pletely with one another, because at this point hydrogen iodide is 
 partly resolved into its elements (see below). Their union is more 
 complete if both elements, in the form of vapor, are conducted over 
 heated platinum sponge. It cannot be obtained by acting upon iodides 
 with sulphuric acid, because the resulting hydrogen iodide decomposes 
 more easily than the bromide. It is formed, however, similarly to the 
 latter, by acting on phosphorus iodide with water : 
 
 PI3 + 3H,0 = H3PO3 + 3 HI. 
 
 A more convenient procedure consists in adding 15 parts of iodine (to which lo parts 
 of water have been added) gradually and while cooling to a mixture of one part of red 
 phosphorus and four parts of water, and then gently heat the same ; or allow an emul- 
 sion of red phosphorus (5 parts) with water (10 parts) to flow gradually and at first very 
 slowly upon iodine (100 parts) moistened with water (lo parts) (Lothar Meyer, Ber. 20 
 (1887), 3381). Hydrogen iodide prepared in this way is invariably contaminated with 
 phosphorus compounds. The pure gas can only be made by the method mentioned 
 above : by conducting iodine vapor and hydrogen over heated platinum sponge. Water 
 is then saturated with the escaping gas and on heating this fuming hydriodic acid a steady 
 current of hydrogen iodide is easily obtained. It is dried by passing it over phosphorus 
 pentoxide. (See Bodenstein, Zeit. f. phys. Ch. 13 (1894), 59.) 
 
 Another method of obtaining aqueous hydrogen iodide consists in 
 passing hydrogen sulphide into water to which finely pulverized iodine 
 is added, as long as decolorization occurs : 
 
 H^S + 12 = 2HI + S. 
 
 Filter off the separated sulphur and distil the liquid. Dry iodine and 
 dry hydrogen sulphide do not react upon one another. 
 
 Hydrogen iodide is a colorless gas; it fumes strongly in the air; its 
 density is 128 (02 = 32) or 4.4 (air= i). Under a pressure of 4 atmos- 
 pheres (at 0°) it is condensed to a liquid which boils at — 34°-* It 
 solidifies at lower temperatures and remelts at — 51°. It is easily soluble 
 in water, i volume of the latter dissolving 450 volumes of the gas at 10°. 
 The solution saturated at 0° C., has a specific gravity of i 99, and fumes 
 strongly in the air. If the solution be heated hydrogen iodide is ex- 
 pelled, the temperature rises and at 126° a solution of 1.70 specific gravity, 
 containing 57 ])er cent, of hydrogen iodide, distils over. Its composi- 
 tion corresponds closely to the formula HI 5H2O. 
 
 Hydrogen iodide is a rather unstable compound. Its decomposition 
 takes place at all temperatures at which it exists as a gas ; the speed of its 
 disintegration increases rajiidly with the temperature. While only the 
 two-thousandth part of the hydrogen iodide separates into hydrogen and 
 iodine in ninety days at a temperature of 100°, almost one-fourth of it 
 
 * Consult Estreicber, Zeit. f. phys. Cli., 20 ('1896), 605, upon the behavior of hydrogen 
 chloride, bromide and iodide at low temperatures. 
 
64 
 
 INORGANIC CHKMISTRV. 
 
 will be decomposed in fifteen minutes at 518°, and tlie maximum decom- 
 position for tins temperature will then have been attained (see Ibssoci- 
 ation of Water). At high temperatures oxygen decomposes into water 
 and iodine : 
 
 2lII -t- O^IbO + Ij. 
 
 On bringing a flame near the mouth of a vessel containing a mixture 
 of hydrogen iodide and oxygen, violet iodine vapors will be liberated. 
 The same will be noticed when fuming nitric acid is drojijied into a vessel 
 containing the gas; in this reaction the oxygen of the acid oxidizes the 
 hydrogen and liberates iodine. The nitric acid breaks down into com- 
 pounds containing less oxygen. All oxidizing bodies behave in the same 
 way; the hydrogen iodide abstracts their oxygen and reduces The 
 
 same fact is noticed with concentrated suliihuric acid when theattemjit is 
 made to apply it in liberating hydrogen iodide from an iodide. The 
 oxygen of the air gradually decomposes aqueous hydrogen iodide at the 
 ordinary temperature, and es])ecially in sunlight. The solution, at first 
 colorless, becomes brown, owing to seiiaration of iodine, which in the 
 beginning dissolves; subsequently, however, it separates in beautiful 
 crystals. 
 
 At ordinary temperatures mercury and silver decompose hydrogen 
 iodide, with separation of hydrogen: 
 
 2lIl4-2Ag = 2Agl4-H2. 
 
 Chlorine and bromine liberate iodine from hydrogen iodide (see 
 P- 54 )- 
 
 This compound is employed as a powerful reducing agent in laboratory 
 work. 
 
 4. HYDROGEN FLUORIDE. 
 
 HFl = 20.01. 
 
 It is obtained, like hydrogen chloride, by decomposing fluorides with 
 sulphuric acid. Finely pulverized fluors-par (CaFb) is mixed with con- 
 centrated sulphuric acid and heated gently : 
 
 CaFb -f H.,S04 = CaSO^ + 2HFI. 
 
 Calcium Calcium 
 
 fluoride. sulphate. 
 
 The oiieration is executed in a lead or platinum retort, as the hydrogen 
 fluoride attacks glass and most of the metals. The esca])ing gas is con- 
 ducted into water. To get perfectly anhydrous hydrogen fluoride, heat 
 hydrogen ])otassium fluoride, which then decomposes according to the 
 following equatifin : 
 
 IIKFI, = KFl + 1 1 FI. 
 
 Anhydrous hydrogen fluoride is a colorless, very mobile liquid, fuming 
 strongly in the air, and attracting moisture with avidity; it boils at 
 -{-19.4° (’., and has a specific gravity of 0.98 at 12°. To recondense 
 llie gas it must be cooled to — 20°. Hydrogen fluoride solidifies at 
 — 102.5° and rcmells at — 92.5°. 
 
HYDROGEN FLUORIDE. 
 
 65 
 
 The gas density of hydrogen flouride equals 20.01 (O,^ — 32) at 100°, corresponding to 
 the molecular formula HFl. At 30°, however, it is twice as large, equaling 40. It fol- 
 lows, therefore, that the molecules of the gas at the latter temperature correspond to the 
 formula H2Fl2» consist of two chemical molecules of HFl (compare Arsenic Tri- 
 oxide), 
 
 The concentrated aqueous solution fumes in the air ; when heated, 
 hydrogen fluoride escapes; the boiling temperature increases regularly 
 and becomes constant at 120°, when a solution distils over, the specific 
 gravity of which is 1.15, and its percentage of hydrogen fluoride is 35.3. 
 'The vapors as well as the solution are poisonous, extremely corrosive, 
 and produce painful wounds upon the skin. 
 
 Hydrofluoric acid dissolves all the metals, excepting lead, gold and 
 platinum, to form fluorides. It decomposes all oxides, even the anhy- 
 drides of boric and silicic acids, which it dissolves to form boron and 
 silicon fluorides. Glass, a silicate, is also acted upon; hence the use of 
 the acid for etching this substance (compare Silicon Fluoride). To do 
 this, coat the glass with a thin layer of wax or paraffin, draw any figure 
 upon it with a pin, and then expose it to the action of the gaseous or 
 liquid hydrogen fluoride. The exposed portions appear etched ; gaseous 
 hydrogen fluoride furnishes a dim, and liquid hydrogen fluoride a smooth, 
 transparent etching. 
 
 Vessels of lead, platinum, or caoutchouc are employed for the preser- 
 vation of hydrofluoric acid, as they are not affected by it. 
 
 These halogen derivatives of hydrogen show great resemblance to one 
 another. At ordinary temperatures they form strongly smelling and 
 fuming gases, which can be easily condensed to liquids. Their fuming 
 in moist air is due to the fact that they are dissolved by the water 
 vapor and the resulting solutions appear as a cloud consisting of very 
 minute drops. Being readily soluble in water, they are only partly ex- 
 pelled from their saturated solutions by boiling; solutions of definite 
 composition distil over, but these cannot be regarded as definite chemical 
 combinations of the halogen hydrides with water, because their com- 
 position depends upon the pressure at which they are boiled (Roscoe). 
 
 As acids they neutralize the bases and form haloid salts, which also re- 
 sult by the direct union of the halogens with metals. 
 
 The densities of the halogen hydrides exhibit a gradation similar to 
 that of the densities of the halogens (p. 56) : 
 
 HFl HCl HBr HI 
 
 Densities, 20.01 36.46 80.97 127.86 
 
 The difference in chemical deportment corresponds to this gradation. 
 Hydrogen fluoride is the most stable, and acts most energetically ; fluo- 
 rine unites in the dark with hydrogen ; chlorine combines with it in sun- 
 light, while bromine and iodine require higher temperatures for their com- 
 bination with it. On the other hand, hydrogen iodide is decomposed at a 
 gentle heat (180°), into its constituents; the more stable hydrogen bro- 
 mide at 800°, while hydrogen chloride remains unaltered up to 1500° C. 
 
66 
 
 INORGANIC CHEMISTRY. 
 
 Corresponding to this we liave the very energetic action of fluorine, and 
 the tolerably ready action of chlorine iii)on water, oxygen sej)arating at 
 the same time : 
 
 n/) -p Cl, ^ 2IICI -f O. 
 
 Iodine does not act upon water. The opjjosite reaction occurs : oxygen 
 decomi)oses hydrogen iodide into water and iodine: 
 
 2lII -f O = II/) + I,. 
 
 Bromine occupies an intermediate position between chlorine and 
 iodine; in dilute acpieous solution it decomposes water into hydrogen 
 bromide and oxygen, while a concentrated solution of hydrogen bromide, 
 on the contrary, is partly decomposed by oxygen into water and free 
 bromine. 
 
 From all theabove it is evident that the affinity of fluorine for hydrogen 
 is the greatest; then follow chlorine and bromine, and finally, as the 
 least energetic element, we have iodine (see j). 57). Fluorine holds 
 an exceptional position with the other halogens in that its hydride is a 
 liquid at the ordinary temperature, and many of its metallic derivatives 
 show a solubility directly oiiposite to that of the metallic chlorides, bro- 
 mides, and iodides. This will be discussed later. 
 
 THERMO-CHEMICAL DEPORTMENT OF THE HALOGENS. 
 
 The quantities of heat, disengaged or absorbed in chemical reactions, 
 afford the most satisfactory explanations of the deportment of the 
 halogens with hydrogen, and indeed of all the chemical elements and 
 compounds toward one another. These heat changes are also called 
 positive and negative thermal values (Jieat fnodulus') (see p. 30). 
 
 The quantities of heat are estimated in heat units or calories. The quantity of heat 
 required to raise one gram of water 1° C. from 15° (measured with an air-thermometer), is 
 taken as the heat unit (small calorie), or a thousand times this quantity can be taken ; 
 then it would be the quantity of heat needed to raise i kilogram of water from 15° to 16° 
 (large calorie. Cal.). Large calories will be used in the following pages. (See further, 
 Nernst, Theoretische Chemie, 2 Aufl. (1898), 10.) 
 
 To obtain data that may be easily compared, the quantities of heat are 
 referred to qtiantities in grams corresponding to the atomic or molecular 
 weights of the elements entering into combination. Thus, in the union 
 of 19 grams of fltiorine (FI 19) with i.oi grams of hydrogen (H = i.oi) 
 to form 20.01 grams of hydrogen fluoride (HFl = 20.01), 37.6 Cal. are set 
 free, and in the formation of 36.46 grams of hydrogen chloride (HCl = 
 36.46) from its elements 22.0 Cal. When 79.96 grams of bromine 
 ( Br = 79.96) combine with i.oi grams of hydrogen to form hydrogen 
 bromide 8.4 Cal. are developed, while in the formation of 127.86 grams 
 of hydrogen iodide from solid iodine and hydrogen 6.0 Cal. are ab- 
 sorbed, but with iodine vajior and hydrogen only 1.5 Cal. (4.5 Cal. being 
 required for the vaporization of the given amount of iodine). 
 
THERMO-CHEMICAL DEPORTMENT OF THE HALOGENS. 
 
 67 
 
 This may be expressed according to the method of J. Thomsen, as follows : 
 
 (H,F1) = -f37.6 Cal. ; (H,C1) == -f 22.0 Cal. ; (H,Br) = +8.4 Cal. ; 
 
 (H,I)=-i.5Cal. 
 
 The first three reactions, in which heat is liberated, are exotherDiic^ while the heat- 
 absorbing combination of iodine with hydrogen represents an endother 77 iic reaction (see 
 p. 30). The energy-content of hydrogen fluoride, chloride and bromide is less, and that 
 of hydrogen iodide greater than^ that of their components. 
 
 The quantity of heat disengaged in a combination must not be regarded as a measure 
 of the chemical affinity. As the elements with few exceptions ; see argon, helium, mer- 
 cury, and cadmium — do not exist as free atoms, but as molecules these require a definite 
 quantity of heat to decompose them into atoms before they can enter into chemical reac- 
 tion. This necessitates a definite amount of work (addition of energy). The union of 
 chlorine with hydrogen proceeds according to the molecular equation (p. 76) : 
 
 HH-f C1C1 = 2HC1 o 
 
 The heat here disengaged (2 X 22.0 =r 44.0 Cal.) is the algebraic sum of the follow- 
 ing unknown thermal values : (i) — x Cal., required for the decomposition of the hydro- 
 gen molecule into free atoms ; (2) — y Cal., consumed in the decomposition of the chlorine 
 molecule; (3) -fz Cal., liberated in the formation of hydrogen chloride from the free 
 chlorine and hydrogen atoms ; hence z — x — y = 44.0, i. e . , the thermal value is deduced 
 from the three unknown values. 
 
 It is very probable that the union of the free atoms always occurs with heat-disengage- 
 ment, and the heat-absorption, observed in many chemical changes, is to be credited to 
 the decompositions to which attention has been directed. In the formation of hydrogen 
 iodide, for example, the sum of — x and — y is very probably greater than z. 
 
 The greater the heat developed in a reaction, the more energetically 
 and the more readily will it occur, and in general, the resulting com- 
 pounds will be the more stable (compare p. 65, Behavior of the 
 Halogen Hydrides). The energetic reactions, those which are accompa- 
 nied by very appreciable evolution of heat, must be viewed as transitions 
 of systems from a state of comparative instability to one of greater per- 
 manency. The opposite occurs if the chemical change only takes place 
 upon the addition of external energy. The compound then produced 
 passes very readily into the original and more stable system. In this 
 sense the principle of greatest heat-development is true (p. 31). 
 
 In this way it can be understood from the thermal values of the halogen 
 hydrides why it is that iodine is displaced by the other halogens, bromine 
 by chlorine and fluorine and chlorine by fluorine from their hydrides and 
 their metallic derivatives — corresponding to the following thermo-chem- 
 ical equations : 
 
 HI + Cl =r HCl + I . . . (+28 Cal.) 
 
 ( — 6.0) (22.0) 
 
 IIBr-|-Cl=:HCl + Br . . . (-f 13.6 Cal.) 
 
 (8.4) (22.0) 
 
 HCl -f Fl= IIFl + Cl . . .(+15.6 Cal.) 
 
 (22.0) (37.6) 
 
 The thermo-chemical sign of a reaction is obtained by deducting from the heat of 
 formation of the jjroducts that of those reacting. 
 
 The reactions do not proceed wholly in the sense indicated. They are limited by 
 accurately opposing forces. When bromine acts upon hydrochloric acid the transposition 
 takes place, if even in a very slight degree, according to the equation Br -f- HCl = 
 
68 
 
 INORCIANIC CHEMISTRY. 
 
 IIBr -(- Cl. And it must therefore be assumed from tliis equation that when chlorine 
 acts upon liydrol)romic acid a portion of the tnomine liberated will trans])ose itself with 
 the hydrogen chloride which has been formed : the reaction reverses itself, is retrogres- 
 sive. To what extent this hai)pens, to what degree of divisions these reactions, occurring 
 simultaneously but proceeding in opposite directions, will go, depends uj)on the nature 
 of the reacting bodies, the ratios of their quantities in unit volume, the temperature, the 
 time, the pressure ; and finally a state of ecpiilibrium will be developed. 'I'lie retrogres- 
 sive reaction very frecjiiently advances so slowly, conse(|uently amounts to so little, that 
 it may ajjpear as if it actually did not occur at all. 
 
 These facts demonstrate that Terthelot’s principle of the greatest heat development is 
 not universally correct although it often indicates the direction in which a transhnniatioji 
 occurs easily and completely. As an argument favoring this we have the different decorn- 
 ])osability of the gaseous halogen hydrides by oxygen, taking into consideration, of course, 
 the heat of formation of water. For acjueous vapor this is 57.2 Cal., while for the liquid 
 it is 68.3 Cal. The union of i gram of hydrogen with 8 grams of oxygen is attended 
 
 with a thermal value of — 28.6 Cal., which is greater than that of hydrogen chloride, 
 
 bromide or iodide but less than that of hydrogen fluoride. Consequently, oxygen should 
 displace chlorine, bromine and iodine but not fluorine from their re.spective hydrides, and 
 this takes place the more readily the greater the difference in the heat of formation. 
 
 In fact, we observed (p. 64) that when a flame, or some glowing substance, was brought 
 in contact with a mixture of hydrogen iodide and oxygen, all the iodine was separated in 
 the form of vapor, in accordance with the following equation : 
 
 2111 + O = H.p (vapor) H 2I . . . ( + 69.2 Cal.) 
 
 Oxygen also liberates bromine from hydrogen bromide at a temperature of about 500° 
 (neither hydrogen bromide nor water suffer di.ssociation at this temperature). Aqueous 
 vapor is also produced. Hydrogen chloride, however, is only partially decomposed by 
 oxygen even at higher temperatures. 
 
 Phirther, in accordance with this idea, in a mixture of chlorine, hydrogen and oxygen, 
 the hydrogen will first unite with chlorine and if any remain then with oxygen, although 
 the heat of formation of water (H2,0) = 57.2 Cal. is greater than that of hydrogen 
 chloride (Cl2,H2') =44.0 Cal. For a more exhaustive study of thermo-chemical relations 
 the student is referred to H. Jahn’s Grundsatze der Thermochemie (2 Aufl. Wien, 1892), 
 and also to the works of Ostwald and of Nernst (see pp. 49, 66). 
 
 COMPOUNDS OF THE HALOGENS WITH ONE 
 ANOTHER. 
 
 These compounds, formed by the union of the halogens with one 
 another, are very unstable, and it may be remarked here, that this is also 
 true of most derivatives obtained from elements which are chemically 
 similar. 
 
 When chlorine is conducted over dry iodine, the latter being in excess, 
 iodine monochloride results, and when the chlorine is in excess, iodine 
 trichloride is formed. 
 
 Iodine Monochloride — ICl — is a red crystalline mass, fusing at 24.7°, and distilling 
 a little above 100°. Water decomposes it easily, with formation of iodic acid, iodine, 
 atid hydrogen chlorid(;. If fused iodine chloride be allowed to .solidify .slowly at a low 
 temperature { — I(j°) a modification, melting at -| 14°, is produced. The latter, however, 
 reverts very readily, with heat evolution, to the higlier melting body. 
 
WEIGHT PROPORTIONS. 
 
 69 
 
 Iodine Trichloride — 10)3 — is formed upon mixing iodic acid with concentrated 
 hydrochloric acid, and by the action of jdiospliorus pentachloride upon iodic anhydride. 
 It crystallizes in long, yellow needles, and, when heated, suffers decomposition into iodine 
 chloride and chlorine (at ordinary pressure, the dissociation commences at 25°). It dis- 
 solves in a little water without alteration ; but large quantities cause partial decomposi- 
 tion, with formation of iodic and hydrochloric acids. 
 
 Iodine Bromide — IBr — obtained by the direct union of the elements, consists of 
 iodine-like crystals, fusing at about 36°. 
 
 Iodine Pentafluoride — IFl^ — is produced by the action of iodine upon silver fluoride, 
 and forms a colorless, strongly fuming liquid. 
 
 WEIGHT PROPORTIONS IN THE UNION OF THE ELEMENTS. 
 
 ATOMIC HYPOTHESIS. CHOICE OF ATOMIC WEIGHTS. 
 
 In the preceding pages several different and independent methods have 
 been described for the preparation of each of the halogen hydrides. But 
 it is immaterial which of these may be selected in making any of these 
 compounds, for if the product be carefully purified and then analyzed the 
 hydrogen and the halogen will always be present in a definite, unalterable 
 proportion by weight. The percentage composition of the pure halogen 
 hydrides will be found under all circumstances to be the following : 
 
 H 5.05 H 2.77 H 1.25 H 0.79 
 
 FI 94.95 Cl 97.23 Br 98.75 I 99.21 
 
 HFl 100.00 IICl 106.00 IIBr 100.00 HI 100.00 
 
 A regularity exists here, which prevails in all chemical compounds. 
 For it is not alone in the halogen hydrides, but without exception m 
 every chemical combination that the constituents occur in definite unalter- 
 able proportions by weight. This observation ascertained by experiment 
 and based u])on facts has been called the law of definite or constant pro- 
 portions. From the period in which the French chemist Louis Proust 
 victoriously defended it against the attack of his countryman Claude 
 Louis Berthollet, the great theorist, in a remarkable controversy waged 
 from 1799-1807, down to the present no fact has been observed which 
 contradicts the law. 
 
 It will be advisable for later considerations that the quantities of the 
 halogens be calculated from the numbers given above for the percentage 
 composition of the haloid acids, which combine with a definite quantity 
 by weight of hydrogen, the constituent common to these four compounds. 
 For reasons to be discussed later hydrogen is no longer taken as unit, as 
 has been done in the past, and it will be better, therefore, to compare the 
 quantities by weight of the elements with one another, which are capable 
 of uniting with i.oi ])arts by weight of hydrogen. In this manner we 
 arrive at the following numbers: 
 
 II I.OI II I.OI II I.OI II I.OI 
 
 FI 19. Cl 35.45 Br 79.96 I 126.85 
 
 IICl 36.46 IIBr 80.97 
 
 III 127.86 
 
 1 1 FI 20.01 
 
70 
 
 INORGANIC CUKMISTRY. 
 
 This simple recalciilalion arfords a dearer insiglit into the ratios by 
 weight according to whicli the halogens unite with hydrogen. We thus 
 discover that 19 parts of lluorine, 35.45 parts of chlorine, 79.96 jarts of 
 bromine, and 126.85 P-'ii'ts of iodine, inasmuch as they are capable of 
 uniting with i oi parts of hydrogen, are eipial or e(pii valent to one 
 another. These numbers answer not only for the derivatives of the halo- 
 gens with hydrogen, but we find them in the compounds of the halogens 
 with one another, and in their derivatives with other elements, 'rims in 
 iodine monochloride the chlorine and iodine are present in proportions 
 by weight 35.45 : 1 26.85 iodine monobromide the ratio of bromine 
 
 and iodine is 79.96: 126.85. grams of sodium be converted into 
 
 fluoride, chloride, bromide, and iodide the (quantities of the halogens 
 required for this purpose will again be in the ratio of 19 : 35.45 : 79.96: 
 126.85. same occurs if we substitute potassium, calcium, magne- 
 
 sium, zinc, silver, etc., for sodium. Thus 19 q)arts by weight of fluorine 
 combine with the following weights of the metals: 23.05 i)arts sodium, 
 39.15 quarts potassium, 32.7 quarts zinc, 31.8 quarts copper, 100. 15 q^iarts mer- 
 cury, — and 35.45 parts chlorine, 79.96 q)arts bromine, and 126.85 
 iodine combine with exactly the same quantities by weight of these 
 metals. 
 
 Let us take another example. On bringing copper into a solution of a 
 mercuric salt the former dissolves, while mercury seq)arates out; indeed, 
 31.8 parts of copq^er displace 100. 15 q:)arts of mercury. If zinc be brought 
 into the copper solution thus obtained, it will dissolve, while coqjper 
 seq^arates — and 32.7 quarts of zinc seq^arate 31.8 quarts of copper. Further- 
 more, zinc disqdaces the hydrogen in acids; from all of them 32.7 parts 
 of zinc seq^arate i.oi quarts of hydrogen. In all these reactions we ob- 
 serve the elements apq^earing in the same quantities by weight. 
 
 There is a net of q^erfectly definite proportions, by weight, connecting 
 all these bodies with one another, and also the reactions which occur 
 between them, i.oi parts of hydrogen combine with 35.45 qiarts of 
 chlorine, and this quantity of the latter with 23.05 parts of sodium, 31.8 
 parts of copper, 32.7 parts of zinc. 32.7 qiarts of the latter rnetal precipi- 
 tate 31.8 parts of copper and 100.15 parts of mercury, from their salt solu- 
 tions, and these quantities of the two metals are caqiable of uniting with 
 the quantities of fluorine, bromine, and iodine which, like the 35.45 parts 
 of chlorine, combine with 23.05 qiartsof sodium or i.oi parts of hydrogen, 
 etc. Proceeding in this way with i.oi qiarts by weight of hydrogen we 
 obtain a number for each element which may be called its co 7 ?ibinin!:!; 
 7 ueight. It will be discovered that the atomic weight of oxygen (O = 16) 
 is the basis of all these number ratios. What has been written may be 
 summarized thus: The elements combine with one anothe)' in the ratio of 
 their combining weights. 
 
 A series of very im])ortant facts, however, compels us to accord this law 
 a broader meaning. It often transq)ircs that two elements combine with 
 one another not only in one pro|)ortion by weight, as in the case of the 
 halogens and hydrogen, but in several ratios, 'rims, there are two com- 
 I)ounds of chlorine with iodine — the monochloride and the trichloride. 
 The first of these always contains 35.45 qxarts by weight of chlorine to 
 
WEIGHT PROPORTIONS. 
 
 71 
 
 126.85 parts by weight of iodine, and the second 3 X 35-45 = 106.35 
 parts of chlorine to the same quantity of iodine. Two compounds of 
 hydrogen and oxygen are known : water and hydrogen peroxide. In 
 water there are always 8.00 parts by weight of oxygen to i.oi parts by 
 weight of hydrogen, while in the peroxide to i.oi parts of hydrogen 
 there are 8.00 X 2 = 16 parts of oxygen. 35.45 parts of chlorine com- 
 bine not only with 31.8 parts of copper and 100.15 parts of mercury, but 
 also with 63.6 parts of copper and 200.3 parts of mercury. Oxygen 
 forms five distinct compounds with nitrogen with the following propor- 
 tions by weight, in which 8, the combining weight of oxygen (see above), 
 is made the basis of comparison : 
 
 Nitrogen. Oxygen. 
 
 Nitrous Oxide, 14.04 parts 8 parts 
 
 Nitric Oxide, 14.04 “ 16 “ =2X8.00 
 
 Nitrous Anhydride, 14.04 “ 24 “ =3X8.00 
 
 Nitrogen Dioxide, 14.04 “ 32 “ =4X8.00 
 
 Nitric Anhydride, 14.04 “ 40 “ = 5 X 8.00 
 
 Proust observed that two elements could combine with each other in 
 different proportions by weight and that in so doing the composition 
 changed by definite increments. The underlying law, however, was first 
 recognized and propounded (evidently as the result of atomic considera- 
 tions) by John Dalton, and definitely established on a scientific basis 
 through the labors of J. J. Berzelius. It is the law of multiple propor- 
 tions : When two elements unite in several p 7 'oportio 7 is, the quantities of 
 the seco 7 td ele 77 te 7 it co 77 ibined with definite a 77 iounts of the first bear a si 77 iple 
 ratio 7 ial ratio to each other. The law of definite combination by weight 
 can be so expanded that it will at the same time include the law of con- 
 stant and also that of multiple proportions. Then it would read: The 
 ele 77 ients 07 ily unite m the ratio of their co 77 ibming weights or simple ratio 7 ial 
 77 iultiples of the sa 77 ie. 
 
 Compounds, therefore, contain their constituents either in the ratio of 
 their combining weights, or some simple rational multiple of these com- 
 bining weights — a simple, self-evident, analytical view of the law. 
 
 The elements of the doctrine of weight-combinations may be observed 
 and gathered from the investigations of the German chemists — Karl 
 Friedrich Wenzel and Jeremias Benjamin Richter upon the neutralization 
 of bases and acids, and the alternating transposition of salts. Following 
 the example of Richter this division of our science is even yet designated 
 stoichiometry (r« ffror/sJa, the constituents; [xirpov, measure), and the 
 laws just deduced are the stoichio 77 ietric laws. It was in this particular 
 direction that Berzelius, from 1808 forward, achieved so much by many 
 hundred accurate analyses which aided very materially in establishing 
 a foundation of irreproachable facts for the jireceding doctrine. 
 
 [Wenzel: Die T.elire von der Verwandtschaft, Dre.sden, 1777 ; Richter: Ueber die 
 neueren (jegen.stande der Cheniie ; lie.sonder.s iin 7, 8, and 9 Stuck, 1796-1798. Com- 
 pare also: Rerzeliii.s’ Lehrbiich der (Jhemie, 5 Aufl., Bd. iii, 1147; and Versuch, die 
 bestimmten und einfachen Verhaltnisse aufzufinden, etc., in Ostwald’s: Klassiker der 
 exakten Wi.ssenschaften, Nr. 35.] 
 
72 
 
 INORGANIC CIIRMIS’l'kV. 
 
 The statements tlnis far made liave l)een free from assumption and liave 
 been i)roved l)y exi)eiiment and analysis; they find malhemalieal ex- 
 pression in tlie law of chemical pro])orti(ms. JUit now speculation enters. 
 To ex[)lain the remarkable regulariiies in the ratios John Dalton took 
 refuge in the atomic hypothesis — “one of the greatest steps of which 
 chemistry availed itself in advancing to jierfection ” — berzelius. 
 
 As previously indicated in the introduction ([i. 25) Dalton assumed 
 that the elementary atoms iireferred uniting in the ratio of 1:1; 
 and whenever but one compound of two elements ajipeared Dalton re- 
 garded it as comiiosed of an atom of each of the two elements. If several 
 com])oiinds existed he viewed the first as consisting of A -f- T, the second 
 of A -|- 2B, the third of 2 A -J- B, and the fourth of A -j- 3B (New Sys- 
 tem of Chemical Philosophy, vol. 1(1808)). By this assumption the 
 facts, expressed in the stoichiometric laws, meet with an astonishingly 
 simple explanation. If hydrogen chloride contains one atom of hydrogen 
 to one atom of chlorine its composition must always be the same: law of 
 constant proportions. If an atom each of hydrogen, sodium, potassium, 
 silver, etc., unite with an atom of fluorine, chlorine, bromine, and iodine 
 to form the corres[)onding fluoride, chloride, bromide, and iodide, then 
 in all these compounds the fluorine, chlorine, bromine, and iodine must 
 appear in the same constant combination ratios, i. e., the quantities of 
 fluorine, chlorine, bromine, and iodine, which unite with a definite quan- 
 tity of another element, must always be to one another as their atomic 
 weights: law of combining weights. If iodine at one time unites with 
 one atom and again wdth three atoms of chlorine, the quantities of the 
 latter which iodine in one instance requires to yield the monochloride 
 and at another time the trichloride stand in the projiortion of i : 3 (Cl : 
 3CI): law of 7 ?iultiple proportions. 
 
 If two elements combine in but one proportion by weight then the re- 
 sulting compound, on the assumption of Dalton, contains one atom of each 
 element. This at once makes it possible to determine the ratio of the 
 atomic weights of these elements. Thus, in hydrogen chloride there are 
 97.23 percent, of chlorine to 2.77 per cent, of hydrogen, and on the as- 
 sumption of Dalton the compound consists of one atom of hydrogen and 
 one atom of chlorine, then these numbers possess a deejier meaning, for in 
 that case the atomic weights of the two elements will be as 2.77 : 97.23. 
 Similarly, the ratios of the atomic weights of the other elements may be 
 determined and the relative ato 7 fiic weights would re.sult just as soon as 
 some number is selected as the atomic weight for any one element and 
 the ])roportion numbers are referred to this number. The comparison 
 element should be one which forms analyzable derivatives with most of 
 the other elements. The number selected for its atomic weight is a 
 matter of ])racticability and of general agreement. These are the reasons 
 which were potent in the choice of oxygen as the standard. Its atomic 
 weight has been ])laced at t6. Our relative atomic weights, therefore, refer 
 — Compare Bericht der Kommission fiir die Festsetzung der 
 Atomgewichte (Landolt, Ostwald, Seul)ert), Ber. 31 (1898), 2761. 
 
 Hydrogen, the spocificnlly liglilcst cleiiuMif, was chosen by Dalton as the standard ; its 
 atomic weij^ht was taken as unit. 'I'lie disadvantage in this instance is tliat hydrogen 
 
GENERAL PROPERTIES OF GASES. 
 
 73 
 
 forms comparable derivatives, allowing of accurate analysis, with few elements. Conse- 
 quently the ratio of its atomic weight to that of other elements is usually determined with 
 the aid of oxygen. Since, however, the ratio of oxygen to hydrogen cannot be as sharply 
 determined as that of oxygen to the atomic weights of many other elements, errors which 
 are not justihable creep into the atomic weights by virtue of this recalculation. These 
 reasons led Berzelius, to whom we aie indebted for the first accurate atomic weight deter- 
 minations, to reject the Dalton unit. He chose, with analytical acuteness, oxygen as the 
 standard, setting its atomic weight at lOO (O = looj. 
 
 About the middle of the present century a return to the Dalton hydrogen unit occurred. 
 The movement was inaugurated mainly by the French, in order to distinguish new theo- 
 retical views at first sight from the older notions held by Berzelius and his students. The 
 standard has met with almost universal adoption from that period down to the present time. 
 
 In a recalculation of the atomic weights, conducted by L. Meyer and Carl Seubert 
 (1893) with all previous determinations as basis, the atomic ratio O : H =: 15.96 : i was 
 assumed as the most probable. Hitherto it has been regarded as the most sati.sfactory by 
 the majority of chemists and has been adopted in this and other text-books. Since, how- 
 ever, J. Thomsen (1895) and particularly E. W. Morley by a series of most admirable and 
 painstaking experiments have redetermined this ratio, the unsatisfactoriness of the hydro- 
 gen unit has become more apparent. All atomic weights determined in ratio to oxygen 
 are affected by the change in this proportion O : H. To avoid similar difficulties in the 
 future — for the ratio O : H cannot be regarded as definitely decided — chemists have taken 
 oxygen as the element of comparison, and in order that the new atomic numbers may be 
 as similar as possible to those previously used, the suggestion of Ostwald that the atomic 
 weight of oxygen be placed at 16 has been followed. The atomic weight of hydrogen 
 will then be H = 1.008, or for all practical purposes H — i.oi. Compare the historical 
 data in the Zeit. f. anorg. Chem. 14 (1897), 250, 256, and also the report of the Atomic 
 Weight Commission to which reference has already been made. 
 
 The relative atomic weights, determined in the manner indicated, are, 
 of course, only correct if our assumption as to the number of atoms in 
 union with one another is correct. The chemical analysis of water shows 
 it to consist in one hundred parts of 11.2 parts of hydrogen, and 88.8 
 parts of oxygen, hence 2.02 parts of hydrogen correspond to 16 parts of 
 oxygen. If, then, water contains one atom of oxygen in union with one 
 atom of hydrogen the atomic weight of hydrogen would be 2.02 (O = 16). 
 It may be possible, however, that water consists of two atoms of hydrogen 
 and one atom of oxygen or of one atom of hydrogen and two atoms of 
 oxygen, etc. In the first case the atomic weight of hydrogen would be 
 I.OI and in the second 4.04, etc., etc. 
 
 Analytical results afford nothing positive for the solution of this diffi- 
 culty. This was the condition in which the question relating to the 
 magnitude of the atomic weights existed fifty years ago. To establish 
 these correctly, various views were allowed to prevail, none, however, 
 with positive foundation. The question can only be solved upon a basis 
 of special views of the gaseous condition and new facts lying chiefly in 
 the domain of organic chemistry. 
 
 Nothing definite is known as to the actual magnitude of the atomic weights. This has 
 thus far been an unimportant matter for purely chemical consideration. 
 
 GENERAL PROPERTIES OF GASES. ATOMIC-MOLECULAR 
 
 THEORY. 
 
 Gases, i. <?., aeriform bodies, above their critical temperature, posse.ss 
 properties in pronounced degree which are inde])endent of their chem- 
 ical composition. As regards alterations of pressure and Icnqicrature a 
 
 7 
 
74 
 
 inor(;anic chemistry. 
 
 given volume, for examiile, one liter, of hydrogen deports itself in the 
 same regular manner as one liter of any other gas under tlie same condi- 
 tions. Here chemical composition no longer plays a role. Just as soon 
 as a substance i)asses into the gas condition it acquires the proiierties 
 expressed by the laws about to be discussed. 
 
 The relations between a volume of gas and the pressure under which it 
 exists are expressed, for like temperatures, in a law named after its first 
 discoverer, Hoyle (1662), although it often aiijiears under the name of 
 Mariotte (1679). reads: the volume occupied tfy a given mass of gas 
 under different pressures is inversely as the pressure. If the jiressure is 
 trebled the volume will be diminished to one-third; if the jiressure falls 
 to one-half, the volume will double, etc. If the gas volumes be re|)re- 
 sented by v, v' and the pressures by ]>, p', the law can be algebraically 
 expressed in the equation 
 
 or 
 
 p : p' = \ V 
 
 p : V = p' : v'' = Constant. 
 
 That is, the product of the pressure and volume of a given quantity of 
 gas, with constant temperature, is unalterable. 
 
 The second law was discovered in 1802, almost simultaneously, by Gay- 
 Lussac and by Dalton. It reads: the volume of any gas increases 
 or 0.003665 of its volume at 0°, for every degree centigrade which its 
 temperature rises under the same pressure. This number, the coefficient 
 of expansion of gases, is ordinarily represented by the letter a. If the 
 volume Vo of the gas at 0° was equal to i, then at 1° it is equal to 
 1.003665, at 273° equal to 2, and universally at t" 
 
 Vt = Vo Vo . 0.003665 . t = Vo (l + at). 
 
 By combining the two laws there results the important equation 
 P V = po Vo (i + at), 
 
 indicating the relations between the volume v, which a gas under pressure 
 p occupies at t“and the volume v^ at 0° under the pressure p^. We shall 
 return to this law when the measurement of gases and their variation from 
 the Boyle-Gay-Lussac law are considered. 
 
 A third law relates to the volume ratios according to which gases 
 combine chemically with one another. These were first investigated by 
 Gay-Lussac, partly in conjunction with A. v. Humboldt. Gay-Lussac 
 announced his observations as follows: The gases combine according to 
 simple volume ratios; the volume of a compound gas bears a simple ratio 
 to the volume of its constituents. In the union of hydrogen and a halo- 
 gen to form a halogen hydride it was found that one volume of hydrogen 
 and one volume of halogen combine to two volumes of the halogen 
 hydride. Hence, in this instance, the union took place without change 
 of volume; the halogen hydride occupied the space previously occupied 
 by its constituents: an important fact for the chemical atomic theory. 
 It is concluded from this and the previously mentioned facts, according to 
 which liydrogen and halogen unite with one another in the ratio of their 
 combining weights, that the ratio between the weights of eipial volumes 
 
GENERAL PROPERTIES OF GASES. 
 
 75 
 
 of these gases, i. e., of their densities, must be the same as that between 
 the combining weights. Each halogen has but one derivative with hydro- 
 gen. Hence, according to Dalton, it may be assumed that in the case 
 of the halogens the combining weight and atomic weight are the same. 
 Then their gas density would be proportional to their atomic weights. 
 Here, again, and also in similar deportments of the simple as well as 
 the compound gases, expressed in the law relating to gases, we are forced 
 to the assumption that in equal volumes of different gases there is an equal 
 number of atoms — smallest particles. Dalton and Berzelius both assumed 
 this conception, but soon abandoned it because it did not lead to a defi- 
 nite, well-defined distinction, justified by facts, between elementary 
 and compound gases. Equal volumes of hydrogen and of chlorine unite 
 without any change in volume to hydrogen chloride. If looo atoms are 
 present in a given volume of hydrogen, looo atoms of chlorine would be 
 required in an equal volume, and by their union there would result looo 
 particles of hydrogen chloride, which would occupy, as demonstrated by 
 experiment, the same volume as the chlorine and hydrogen together : 
 
 : 
 
 looo H -f- looo Cl = looo HCl. 
 
 I vol. I vol. 2 vols. 
 
 A volume of hydrogen chloride, therefore, would contain only half as 
 many particles as a like volume of the elementary gases — a conclusion 
 which contradicts everything known relative to the constant, physical 
 deportment of the gases (being independent of chemical constitution) 
 with reference to the laws of Boyle and Gay-Lussac. 
 
 To solve this contradiction has required almost a half century of chem- 
 ical investigation. It was especially the advances in the domain of 
 organic chemistry which led to the assumption that the elementary sub- 
 stances as well as the compound did not consist of a mass of free atoms 
 but of an aggregation of atom-groups — of molecules (p. 24). This had 
 been advocated (enunciated) by Avogadro (1811), then by Ampere 
 (1814), and subsequently by Dumas (1829), but it was only at the close 
 of the ’50’s that the theory, after being supported by the Parisian chem- 
 ist Gerhardt on a purely chemical basis, received general acceptance and 
 favor [Gerhardt, Lehrbuch der organ. Chemie, deutsche Ausgabe, 4 
 ('1857), 598, 627, etc.]. It is, consequently, necessary to distinguish 
 between atom and molecule {jjiolecula, mass-particles). It is not the 
 atoms but the molecules which, as a rule, exist in the gases as the smallest 
 constituent particles (exceptions will be discovered later; see also p. 55 
 on the Dissociation of the Halogens). It is obvious that the smallest par- 
 ticles of a compound body consist of several atoms. By further subdivi- 
 sion such a molecule breaks down into dissimilar constituents: a molecule 
 of hydrogen chloride is resolved into hydrogen and chlorine. Even the 
 elements, when free, consist, as a rule, of molecules, and the latter of sev- 
 eral, generally of two, atoms. Hence, the previously deduced rule that 
 in equal volumes of the elementary gases there is contained an equal num- 
 ber of atoms must be changed and amplified somewhat as follows: 
 
 The number of molecules in a unit volume of all gases is the same — 
 like pressure and like temperature being presupposed. Whde not entirely 
 correct, this law is very frequently designated as the law of Avogadro. 
 
76 
 
 INORGANIC CHEMISTRY. 
 
 The contradiction between the conclusions deduced from cliemistry 
 and from physics now vanishes. 
 
 'The [)rocess of tlie combination of hydrogen with chlorine (and the 
 other halogens) must be conceived, therefore, to be somewhat like the 
 following: one molecule of hydrogen, containing at least two atoms of 
 hydrogen, acts ui)on one molecule of chlorine, also composed of at least two 
 atoms of chlorine, and there result two molecules of hydrogen chloride : 
 
 II2 ( Cl, — 2IICI. 
 
 i. e., hydrogen chloride contains just as many molecules in an equal vol- 
 ume as hydrogen and chlorine. This is apparent from the folh)wing 
 representation, which attaches itself to the example given on j). 75 : 
 
 500 H, 
 
 I volume. 
 
 + 
 
 500 Cl, 
 
 I volume. 
 
 500 IICl 
 
 500 IICl 
 
 2 volumes. 
 
 In a similar manner two volumes of hydrogen (containing 2n molecules) 
 give with one volume of oxygen (containing n molecules) two volumes of 
 aqueous va])or ; consequently, 2fi molecules of water. In 2>^ molecules of the 
 latter (H^O) there are contained at least 271 atoms of oxygen ; therefore, 
 in 71 molecules of oxygen, 271 atoms of oxygen — or one oxygen molecule 
 co 7 isists of at least two ato 77 is (compare pp. 73, 80). 
 
 2 volumes. 
 
 nO, 
 
 yield 
 
 nH ,0 
 
 nH20 
 
 I volume. 2 volumes. 
 
 Let us take another analytical example. Two volumes (^2n mol.) of 
 ammonia gas break down, under the influence of the electric spark, into 
 one volume ft mol.) of nitrogen and three volumes of hydrogen (each 7 t 
 mol.). From this it is evident, as in the preceding example, that the 
 nitrogen molecule also is composed of at least two atoms : 
 
 nN.. 
 
 1 vulumu. 
 
 3 volumes. 
 
 2 volumes. 
 
GENERAL PROPERTIES OF GASES. 
 
 77 
 
 In the same way it may be shown that the phosphorus molecule consists 
 of at least four atoms of phosphorus (P4), etc., etc. 
 
 Many other facts corroborate the assumption that the molecules of the 
 elements consist of several atoms, for example, the existence of the allo- 
 tropic modifications of the elements (compare Ozone), the chemical reac- 
 tions (compare Hydrogen Peroxide), and the remarkable action of the 
 elements in the moment of their liberation. 
 
 We saw (p. 53) that the oxygen separated from water by chlorine 
 acted much more energetically than free oxygen. Other elements, espe- 
 cially hydrogen, behave similarly in the moment of fomnation — in statu 
 nascendi. As viewed by the atomic molecular theory, this may be very 
 
 Fig. 38. 
 
 Fig. 37. 
 
 Fig. 39. 
 
 easily explained. The free elements (their molecules) are compounds of 
 similar atoms whose chemical affinity has already been partly satisfied. 
 In the moment of their separation from compounds free atoms appear, 
 which, before they combine to molecules, must act more energetically. 
 
 The experiments illustrating and confirming what has been said upon 
 the volume relations of the gases will now be considered : 
 
 I. The concentrated acjueous solution of hydrochloric acid is decomposed by the action 
 of the galvanic current, and the chlorine and hydrogen collected ; these gases separate at 
 ojjjxjsite poles. The electrolysis may be made in an ordinary voltameter (Fig. 37). 
 Hofmann’s apj)aratus is better adapted to this }>urpose (Fig. 38 *). Two glass cylinders, 
 
 * I'ijf. 38 represents Uie volume relations the g-ases liberated in the electrolysis of water; 
 compare p. 92. 
 
78 
 
 INORGANIC CHEMISTRY. 
 
 provi(lo<l at llie lop vvitli .sto])-c()cks, are connected at the lower end with one another and 
 with a funnel tube ; the latter serves to lill the a])paratus with hijiiid ; and also, hy further 
 additions, to jness out the gases collected in the lubes. 'The platinum electrodes are fused 
 into the lower part of both tubes. I n another form of I lolinaim’s apparatus (l‘ig. yj) 
 the electrodes are introduced by means of caoutchouc stoi)i)ers. When the separating 
 gases (in this case chlorine) attack the platinum, carbon electrodes are substituted for the 
 latter. 
 
 To electrolyze hydrogen chloride, fill the aj)i)aratus with hydrochloric acid, of specific 
 gravity 1.125, to which has been added fifteen times its volume of calcium chloride of 
 1.36 si)ecilic gravity ; close the up])er sto])-cocks and connect the electrodes with the poles 
 of the battery, (bases separate in both tubes, and in equal iwlu/nes ; that se])aratcd at the 
 positive ])ole is chlorine ; the other combustible gas is hydrogen. [See Liipke, (Tund- 
 zilge der wissensch. Elektrochemie, llerlin, 1899; Lolhar Meyer, Her. 27 (1894), 850; 
 Haber and Grimberg, Z. f. anorg. Ch. 16 11898).] 
 
 d'liis experiment shows that hydrogen chloride decomposes into e(|ual volumes of 
 chlorine and hydrogen. 
 
 2. The production of liydrogen chloride by the union of equal volumes of hydrogen 
 and chlorine is shown in the next experiment. 
 
 Fill a cylindrical glass tube, provided with stop-cocks at both ends (hig. 40), with 
 equal volumes of chlorine and hydrogen. This is most conveniently done by conducting 
 the gaseous mixture obtained by the electrolysis of hydrochloric acid into the dry tube. 
 (The tube should be filled and kept in the dark, as the gases combine in daylight.) 
 
 When the tube is filled with the mixture, sunlight or magnesium light 
 is directed upon it, when chemical union ensues. On immersing the 
 lower end of the tube into water, and opening the lower stop-cock, the 
 water will rapidly fill the tube, as the hydrogen chloride that was pro- 
 duced dissolves ; all the hydrogen and all the chlorine have disappeared. 
 
 3. A modification of this experiment teaches us another important 
 fact which has reference to the ratio of the volume of the hydrogen 
 chloride to the volumes of its constituents. If the tube filled with 
 equal volumes of chlorine and of hydrogen be opened under mercury, 
 after the explosion, no diminution in volume will be detected, although 
 the mixture of chlorine and hydrogen has been changed to hydrogen 
 chloride. It follows from this that a mixture of equal volumes of 
 chlorine and hydrogen affords the same volume of hydrogen chloride, 
 or, as ordinarily expressed, one volinne of chlorine and one volume of 
 hydrogen yield tzvo volutnes of hydrogen chloride. 
 
 The following experiment confirms this conclusion : Into a bent tube 
 (Fig. 41), filled with mercury, conduct dry hydrogen chloride, and then 
 introduce in the bend of 
 the upper part a little piece 
 of metallic sodium. On 
 heating the latter with a 
 lamp, the hydrogen chloride 
 is decomi)osed, the chlorine 
 combines with the sodium 
 to form sodium chloride, 
 while hydrogen is set free. 
 
 Upon measuring the resid- 
 Fig. 40. liydrogen it will be 4 ** 
 
 found that its volume is 
 
 exactly the half of the volume of hydrogen chloride originally introduced. In the same 
 manner may be shown the fact that in two volumes of hydrogen bromide and hydrogen 
 iodide there is contained in each one volume of hydrogen. It follows further from the 
 (hmsities of bromine and iodine vaj^ors, that the (juantilies of these elements in gas form 
 combinitig with one volume of hydrogen also occupy one volume. Hence, one volume 
 of hydrogen and one volume of bromine 7 >afor yield /zoo zwlumes of hydrogen bromide, 
 and one zwlume of hydrogen and one volume of iodine vapor tzvo volumes of hydrogen 
 iodide. 
 
GENERAL PROPERTIES OF GASES. 
 
 79 
 
 All that has been j^-eviously stated may be summarized as follows: 
 
 1. A// bodies ewe coinposcd of atoms. 
 
 2 . The atoms unite to foriji the 7noiccuies of the simpte and compound bodies. 
 
 3. Equai votiwies of gases contain, at tike teniperatu7'e and under tike 
 pressure, an equat number of motecutes. 
 
 4. The gas densities, therefore, bear the sanie ratio to one another as the 
 motecuiar weights. 
 
 The gases argon and heiiuin, recentiy discovered in the atmosphere, are 
 exceptions. They co?isist of singie atoms at the ordinary temperature. 
 This is aiso true of the haiogens at etevated temperatures, and, so far as we 
 know, of metai vapors {inercury, cad 7 nium, zi 7 ic'). 
 
 Heretofore gas density has been referred to the hydrogen atom (H = 
 i), while the molecular weight was referred to the molecule H2 = 2, so 
 that the density (the specific gravity of the gas) was one-half of the molec- 
 ular weight. At present, the atomic weight of oxygen (16) being taken 
 as the standard in determining the atomic numbers of the other elements, 
 it appears proper to refer the density to the molecular weight of oxygen, 
 Og = 32, whe 7 i the vatues for density and 7 noiecuiar weight wiil coincide. 
 
 The following table contains the atomic and molecular weights of cer- 
 tain metalloids as well as the molecular weights of their hydrogen deriva- 
 tives. The molecular weights (gas densities) show how much a volume 
 of the respective gas weighs, if an equally large volume of oxygen, under 
 similar external conditions, be placed at 32 (p. 73) : 
 
 Atomic Weights. 
 
 Molecular Weights or Gas Densities. 
 
 H = 1. 01 
 
 U , 
 
 — 
 
 2.02 
 
 
 
 ci= 35.45 
 
 Cb 
 
 
 70.9 
 
 
 
 Br 79.96 
 
 Br^ 
 
 = 
 
 159.92 
 
 
 
 
 Br 
 
 
 79.96 
 
 above 
 
 1000° 
 
 I = 126.85 
 
 I2 
 
 = 
 
 253.70 
 
 
 
 
 I 
 
 = 
 
 126.85 
 
 above 
 
 1500° 
 
 
 HCl 
 
 = 
 
 36.46 
 
 
 
 HBr 
 
 
 80.97 
 
 
 
 
 HI 
 
 
 127.86 
 
 
 
 0 =r 16 
 
 O2 
 
 
 32 
 
 
 
 II 
 
 H ,0 
 
 N, 
 
 
 
 18.02 
 
 28.08 
 
 
 
 
 NH3 
 
 = 
 
 17.07 
 
 
 
 P = 31.0 
 
 
 = 
 
 124.00 
 
 
 
 
 PH3 
 
 
 34.03 
 
 
 
 As one liter of oxygen, under normal conditions (p. 44), weighs 1. 4291 grams, the 
 weight of one liter (L) of a gas of molecular weight M can be calculated from the equation 
 
 I, = = 0.04466 M. 
 
 32 
 
 The density D of the gas, referred to air as unit, follows from the equation 
 
 1 ) = -^- 
 28.95 
 
 because the density of air is 4291^^ ~ 28.95, referred to Oj = 32. 
 
So 
 
 INORGANIC CHEMISTRY. 
 
 The numbers deduced from tlicse equations vary somcwliat from tliose obtained by 
 direct observation, d'liis is due to the fact tliat wluit has been said is a simple but at the 
 same time not an absolutely correct expression of the dej)ortment of gases. W'e shall 
 return to this point when the measurement of gases is discussed. 
 
 by determining the gas density it is only possible to fix the molecular weight of an 
 element which exists in the gaseous condition, d'he magnitude of the atomic weight 
 remains uncertain. 'I'he chlorine molecule weighs from its gas density and the standard 
 selected 70.9 units, and consists of two atoms (Cl.^), if we suppose that the atomic weight 
 = 35-45- Its atomic weight could, however, be only the half (or another submultiple) 
 of 35.45 ; then its molecule would consist of four chlorine atoms (Cl^ — 70.90 when Cl is 
 made equal to 17.725), and the formula of hydrogen chloride would be IICl.^. As anal- 
 ysis and the vapor density determination of many, especially organic, derivatives show 
 that the smallest quantity of chlorine in a molecule of a gasifiable body must always be 
 expressed by 35.45, we can well accept this number as the atomic weight. 'I'hat the 
 maximum values thus derived have not been found too high, but correspond to the actual 
 relative atomic weights, follows from the agreement of these numbers with the atomic 
 numbers obtained from the specific heat of the elements. The complete certainty of their 
 correctness we reach by the law of periodicity, which is deduced from these numbers. 
 
 These convincing suppositions and conclusions drawn from actual re- 
 lations, form the atomic molecular doctrine, which is the foundation of 
 the chemistry of to-day. As this doctrine comjtletely explains the quan- 
 titative phenomena arising in the action of the chemical elements upon 
 one another, and as it has been repeatedly confirmed by phenomena of a 
 purely physical nature, and has thereby acquired a high degree of proba- 
 bility, it is only proper and correct that it be designated a theory. 
 
 ■ 
 
 OXYGEN GROUP. 
 
 In this group are included the elements oxygen, sulphur, selenium, and 
 tellurium. They are perfectly analogous in their chemical deportment. 
 One atom of each of these elements is capable of uniting with two atoms 
 of hydrogen. 
 
 1. OXYGEN. 
 
 Atom : O = 16. Molecule : O2 = 32. 
 
 Oxygen (oxygenium) is the most widely distributed element in nature. 
 It is found free in the air; in combination it exists in water. It is an 
 iin])ortant constituent of most of the mineral and organic substances. 
 
 It was discovered, almost simultaneously, by Priestley, in England, 
 i77<i', and Scheele, in Sweden, 1775. Lavoisier, in France, 1774-1781, 
 first exjilained the imi)ortant role attached to oxygen in processes of com- 
 bustion, of res])iration, and of oxidation. 
 
 Pj'eparation. — Heat red mercuric oxide, a compound of mercury with 
 oxygen, ima small glass retort; in this way the oxide is deconq)osed into 
 mercury and gaseous oxygen : 
 
 llg()=: Ilg + O. 
 
OXYGEN. 
 
 8l 
 
 The following method is commonly pursued in the chemical labora- 
 tory : Potassium chlorate, a compound of potassium, chlorine and oxy- 
 gen, is heated in a glass retort (Fig. 42) or flask, and thus decomposed 
 into solid potassium chloride and oxygen : 
 
 KCIO3 = KCl -f 3O.* 
 
 The evolution of the gas proceeds more regularly and requires a less 
 elevated temperature if the pulverized chlorate be mixed with ferric oxide 
 or manganese peroxide. (A small portion of the oxygen is ozonized ; see 
 p. 84). The liberated oxygen is collected over water. 
 
 McLeod (Jr. Chem. Soc., 55, 192) explains the mechanism of the action of man- 
 ganese peroxide on potassium perchlorate, when heated, as follows : First, the peroxide 
 acts on the chlorate producing potassium permanganate, chlorine, and oxygen, 2Mn02 + 
 2KCIO3 = K2Mn203 -f- CI2 -f- O2. Secondly, the permanganate is decomposed by heat : 
 K2Mn208 = K2Mn04 + Mn02 + O2, and in the third stage the change is probably : 
 K2Mn04 -|- CI2 = 2KCI -|- MnOg + 62. 
 
 Very pure oxygen may also be obtained by heating potassium bichro- 
 mate with concentrated sulphuric acid : 
 
 K2Cr20^ + SHaSO^ = Cr2(SOj3 + 2KHSO4 + 4H2O + 3O. 
 
 Besides these many other methods may be employed for the prepara- 
 tion of the gas: e. g., the ignition of manganese and barium peroxides; 
 the decomposition of sulphuric acid at a high heat ; the boiling of a solu- 
 tion of bleaching lime with a cobalt salt, etc. These methods, applied 
 technically, will be considered more fully later. 
 
 A very convenient laboratory method for the preparation of oxygen consists in allow- 
 ing dilute hydrochloric acid to act upon a mixture of barium peroxide (2 parts) and man- 
 ganese peroxide (l part). The gas is evolved at the ordinary temperatures. If the solid 
 ingredients are mixed with gypsum (l part) and a little water the mass can be moulded 
 into cubes, and the oxygen then be generated (like chlorine) in a Kipp apparatus (Ber. 
 20 (1887), 1585)* Kassner obtains a regular and continuous flow of the gas by pouring 
 water upon r, pulverized mixture of i part of barium peroxide and 2j4 parts of potassium 
 ferricyanide (red prussiate of potash) : 
 
 Ba 02 -f K3BaFe2(CN),2 + O 3 . 
 
 Elkan, in Berlin, furnishes oxygen under 100 atmospheres pressure in steel cylinders for 
 trade purposes. The gas is prepared by the method of Boussingault and Brin, by heat- 
 ing barium peroxide to about 700° under reduced pressure : 
 
 Ba02 = BaO -f O. 
 
 The resulting barium oxide is reconverted into peroxide at the same temperature but 
 higher pressure by the oxygen of the air. According to Kassner, calcium plumbate may 
 be similarly applied. It breaks down, under conditions which will be considered later, 
 into lead oxide, calcium oxide, and oxygen : 
 
 Ca2PbO^ BbO -f 2CaO -f O. 
 
 * The chemical equations used here and ])reviously are only intended to represent the 
 manner of the reaction, and to ex])ress the accom])anying relative (juan titles by weight. 
 It should not be forgotten that free atoms do not exist, but that they always occur com- 
 bined in molecules. Molecularly written the equation would be : 
 
 2KCIO3 = 2KCI -f 3O2. 
 
82 
 
 INORGANIC CIIFMISTRY. 
 
 A method wliich will jwohahly he of value later is that of Linde, hy which iiitropjen is 
 first evaporated from litjuid air, and the portions passinj^ over suhsecjuently are lound to 
 contain 75 per cent, of oxygen. 
 
 Properties . — Oxygen is ca colorless, odorless, tasteless gas. One liter 
 of oxygen at 0° C., and 760 niin. jjressiire, weighs 1.4291 grains (15.9 
 times more than one liter of hydrogen). Its sjtecific volume ecpials 699. 7 
 c.c. (}). 45). It is only slightly soluble in water; 100 volumes of the 
 latter dissolve 4.1 volumes of the gas at 0°, and 2.9 volumes at 15°. It 
 is more readily dissolved by absolute alcohol (28 volumes in 100 volumes). 
 
 The critical temiterature of oxygen is — 118°, and its critical jtressure 
 equals 50 atmosi)heres (]:>. 47). Idtiuid oxygen under a pressure of i 
 atmosphere boils at — 184°, and under 9 mm. j)ressure at — 225°. Its 
 specific gravity at — 118° equals 0.65, at — 139*^ it is 0.87, and 1. 124 at 
 — 184° (the boiling point at the ordinary pressure). Liquid oxygen has 
 a bright-blue color (Olszewski). 
 
 Oxygen combines with all the elements excepting fluorine, helium, and 
 
 Fig. 42. 
 
 argon. With most of them it unites directly, accompanied by the evolu- 
 tion of light and heat (p. 57). The burning of combustible bodies in 
 the air depends on their union with oxygen, which is present in the same 
 to the amount of 23 per cent. The phenomena of the respiration of ani- 
 mals are also influenced by the contact of the oxygen of the air — hence 
 the earlier designations of oxygen as inflammable air, and vital air. In 
 j)ure oxygen the phenomena of combustion proceed more energetically 
 than in air. Ignited charcoal or an ignited sliver inflames immediately 
 in the gas, and burns with a bright light. This test serves for the recog- 
 nition of pure oxygen. Sulphur and phosphorus ignited in the air burn in 
 it with an intense light. Even iron burns in the gas. To execute this 
 experiment, take a steel watch spring, previously ignited, attach a match 
 to the end, ignite the same, and then introduce the spring into a vessel 
 filled with oxygen gas. At once the match inflames and ignites the iron, 
 
OXYGEN. 
 
 83 
 
 which burns with an exceedingly intense light and emits sparks. (To 
 protect the vessel from the fusing globules of iron oxide, cover the bottom 
 with a layer of sand.) Iron will burn in any flame if a current of oxygen 
 be conducted into the same. 
 
 Oxygen combines with hydrogen to form water. The union occurs at 
 a red heat, by the electric spark or by the action of platinum sponge (p. 
 46). Palladium asbestos acts similarly. Of all the combustible gases 
 only hydrogen combines by its influence with oxygen at the ordinary 
 temperature. Hydrogen burns in oxygen with a flame; vice versa, oxy- 
 gen must also burn in hydrogen; this may be demonstrated in the same 
 manner as indicated under Hydrogen Chloride (p. 57). A mixture of 
 hydrogen and oxygen detonates violently ; most strongly if the propor- 
 tions are i volume of oxygen and 2 volumes of hydrogen ; such a mix- 
 ture is known as oxyhydrogen gas. The explosibility may be shown in 
 a harmless way by the following experiment: Fill a narrow-necked flask 
 of 150-250 c.c. capacity, over water, yi with hydrogen, and ^ with 
 oxygen; close the opening with a cork, then wrap the flask in a towel, 
 remove the cork and bring a flame near the opening. A violent explo- 
 sion ensues, generally with complete breaking of the flask. The tempera- 
 ture of ignition of oxyhydrogen gas is about 650° (V. Meyer and 
 A. Miinch, Ber. 26 (1893), 2421) (see p. 28). 
 
 Fig. 43. 
 
 The oxyhydrogen flame is only faintly luminous ; it possesses, how- 
 ever, a very high temperature, answering, therefore, for the melting of 
 substances which fuse with great difficulty, e. g., platinum. To get a 
 continuous oxyhydrogen flame, efflux tubes of peculiar construction are 
 employed (Fig. 43) ; through the outer tube, W, hydrogen is brought 
 from a gasometer; oxygen is conveyed through the inner, S, and the 
 mixture ignited at a. Such a flame impinging on a piece of burnt lime 
 or zircon makes the latter glow and emit an extremely bright light — 
 Dru 77 ii 7 iond' s Iwie light, Zirco 7 tiu 77 i light. 
 
 The union of oxygen with other substances is termed oxidatio 7 t. This 
 term, as well as the name Oxygenium (from and yewaM)), or acid 
 producer, suggested by Lavoisier, arises from the fact that acids are some- 
 times formed in oxidation. This the combustion experiments prove. 
 If the vessels, for instance, in which carbon, sulphur, and phosphorus 
 were burned, be shaken wu'th water, the latter will give an acid taste, and 
 redden blue litmus-paper. It was formerly thought that the formation 
 
84 
 
 inorc;anic chemisirv. 
 
 of acids is always conditioned by oxygen. We liave, however, already 
 noticed that the haloid acids — hydrochloric, hydrobronhc, and hydriodic 
 — contain no oxygen. Some of the elements yield acids by their union 
 with oxygen, or more correctly oxides, which form acids with water. Most 
 of these are the metalloids. Tims the following corresi)onding acids are 
 derived from the acid-forming oxides of suli)hur and phosphorus: 
 
 SO3 + II/) = II,S(),. 
 
 Sulphur Sulpliuric 
 
 trioxide. acid. 
 
 r,/)^ -f 1 1^0 = 2iiro3. 
 
 Phosphorus Mctaphosptioric 
 
 pentoxide. acid. 
 
 With oxygen the metals usually yield oxides, which form hydroxides 
 (hydrates) or bases with water : 
 
 K ,0 -f lIjO = 2KOII. 
 
 Potassium Potassium 
 
 oxide. hydroxide. 
 
 CaO -f II2O = Ca (011)2. 
 
 Calcium Calcium 
 
 oxide. hydroxide. 
 
 The salts are produced by the alternating action of acids and bases 
 (see p. 61). 
 
 Thirdly, there exist the so-called iiidiffereJit oxides, which yield neither 
 acids nor bases, with water, e. g. : 
 
 N2O NO Pb 02 . 
 
 Nitrous Nitric Lead 
 
 oxide. oxide. peroxide. 
 
 Oxidation is not only induced by free oxygen or bodies rich in it, but 
 frequently, also, by the halogens; in the latter case the halogens first 
 decompose the water with the elimination of oxygen, which then oxidizes 
 further (compare p. 53). 
 
 The opposite of oxidation, the removal of oxygen, is called reduction 
 (p. 64). Hydrogen {i?i statu nascendi'), and substances giving it off 
 easily (as hydriodic acid), have a reducing action. Most of the metallic 
 oxides are reduced at a red heat by hydrogen, with the formation of 
 water, e. g. : 
 
 CuO + H2 = Cu + II2O. 
 
 Copper oxide. Copper. 
 
 OZONE. 
 
 O3. 
 
 Ozone, discovered in 1840, by Schonbein, is, as will be shown later, a 
 peculiar modification of oxygen, characterized by a remarkable odor and 
 great reactivity, therefore it is called active oxygen. It is obtained from 
 oxygen in various ways; it is almost always ])roduced when oxygen 
 is liberated, or when it takes ])art in a reaction ; thus, in the decomposi- 
 tion of peroxides by concentrated sulphuric acid, or when thevare heated 
 in a current of oxygen to their decomposition-temperature [Brunk, Z. f. 
 anorg. (diem. 10 (i^^95), 222] ; in the electrolysis of water (at the positive 
 
OZONE. 
 
 85 
 
 pole), in the slow oxidation of moist phosphorus, in the combustion of 
 hydrocarbons, and in the action of the so called silent discharge in an 
 atmosphere of oxygen or air. In none of these instances is all of the 
 oxygen ever converted into ozone ; only a small portion — under most 
 favorable conditions 5-6 per cent. — suffers this change. 
 
 The following methods serve for the preparation of ozone : 
 
 1. Put several pieces of stick phosphorus into a spacious flask, cover them about 
 half with water, and allow them to stand for some hours. Or conduct oxygen over 
 pieces of phosphorus placed in a glass tube and moistened with water. Ozone is also 
 formed abundantly when a potassium bichromate solution is substituted for water. 
 [Leeds, Ann. Chem. ig8 (1879), 38.] 
 
 2. Pass the silent discharge through air or oxygen. For this purpose we may employ 
 a Siemen’s ozone tube (Fig. 44, Geissler form), which consists of two glass tubes a and 
 b fitting into one another. Each is coated on the inner side with tin foil or gold leaf. 
 The one coating is connected with the positive and the other with the negative pole of 
 an induction apparatus. Oxygen or air circulates between the two tubes in the direction 
 indicated by the arrows. 
 
 3. Gradually add barium peroxide in small portions (or potassium permanganate) to 
 cold sulphuric acid : 
 
 Ba02 + H2SO, = BaSO, + H2O + O. 
 
 The escaping oxygen is tolerably rich in ozone, and is collected over water. 
 
 Ozone possesses a highly penetrating, peculiar odor (hence its name 
 from to smell), which by prolonged respiration produces bad 
 
 results. In a thick layer, ozone shows a bluish color. If ozonized air be 
 subjected to powerful pressure (150 atmospheres) at a very low tempera- 
 ture, or if ozonized oxygen be conducted through a small tube cooled to 
 — 184° by boiling oxygen, the ozone will condense to a liquid with an 
 indigo-blue color. Liquid ozone, if preserved in a sealed tube, passes 
 into a blue gas, that can be again liquefied by chilling it with boiling 
 ethylene. Ozone is rather stable at the ordinary temperature ; and at 
 more elevated temperatures does not revert to ordinary oxygen as rapidly 
 as it was supposed. This was gathered from the experiments of Brunk, 
 according to whom ordinary oxygen is ozonized when it is conducted 
 over manganese peroxide heated to 400°. According to Mailfert’s experi- 
 ments ozone is fifteen times more soluble in water than oxygen. At the 
 ordinary temperature it takes up half its volume of the gas. In solution 
 the ozone gradually reverts to oxygen without formation of hydrogen 
 peroxide. Unlike ordinary oxygen, ozone, especially in a moist state, 
 oxidizes strongly at ordinary temperatures. Phosphorus, sulphur, and 
 arsenic are converted into phosphoric, sulphuric, and arsenic acids; 
 ammonia is changed to nitrous and nitric acids ; silver and lead are con- 
 verted into the corresponding peroxides; therefore paper moistened with 
 a lead salt is colored brown. Iodine is separated at once from potassium 
 iodide by it : 
 
 2KI -f II.2O -f O3 =: 2K0n + I2 + O2. 
 
 It also oxidizes all organic substances ; therefore the apparatus used in 
 its preparation must not be constructed of caoutchouc. Solutions of 
 dyestuffs, like indigo and litmus, are decolorized. Its ability to turn an 
 alcoholic solution of guaiacum tincture blue is very characteristic of ozone. 
 
86 
 
 INORGANIC CHEMISTRY. 
 
 For the detection of ozone the ordinary potassium iodide starch-paper (Schdnhein) 
 may be used. This is pre])ared by immersing wliite tissue-paper in a starch solution 
 mixed with j)otassium iodide. Tlie iodine which the ozone liberates from the potassium 
 iodide blues the starch-paper. I'he (piantity of ozone may be approximately determined 
 from the rapidity and the intensity of the coloration, 'rhallous hydroxide is a more 
 reliable reagent for ozone than the ])otassium iodide starch-paj)er. (iuaiacum tincture and 
 paper saturated with a lead acetate .solution may also l)e used to detect ozone ; the hrst 
 ac(|uire.s a blue color, the second becomes brown. Other substances al.so blue potassium 
 iodide starch-paper and guaiacum, e. g., chlorine, bromine, nitrogen dioxide, etc., etc. d'o 
 distinguish ozone from the.se, proceed as follows (Ilouzeau) : 'I'ake two strips of violet 
 litmus-i)aper, one of which is saturated with jjotassium iodide, and exj)ose it to the action 
 of the gas ; when ozone is present potassium hydroxide will be formed from the iodide, atid 
 color the violet litmus blue. 'Fhe second paper serves to show the ab.sence of ammonia. 
 
 'rhe preceding reactions of ozone are all produced by hydrogen peroxide, although 
 less ra])idly. The only test answering for the distitiction of very slight (piantities of 
 ozone from hydrogen peroxide is the blackening of a bright strip of silver by ozone. 
 
 Fig. 44. 
 
 Ozone is a variety of oxygen its molecules consist of three atoms: 
 
 3O2 yield 2O3. 
 
 3 vols. oxygen. 2 vols. ozone. 
 
 This is })roved by the following experiments : In ozonizing oxygen its volume dimin- 
 ishes ; upon heating (when ozone is again changed to oxygen), the original volume is repro- 
 duced ; when ozonized oxygen is brought in contact with oil of turpentine or oil of cinna- 
 mon, all the ozone is absorbed and the volume of the gas is diminished. Comj)ari!ig this 
 diminution, corresponding (o the ozone volume, with the expansion which an equal vol- 
 ume of ozonized oxygen suffers after the application of heat, we will find that the first 
 is twice as large as the latter ; this indicates that I volume of ozone yields i ^ volumes of 
 oxygen (.Sorc't). h'rom this it follows that the s|)ecific gravity of ozone must be I times 
 grcaU*r than that of ordinary oxygen, and that if the molecule of oxygen consists of two 
 atoms, the molecule of ozone must contain three atoms. 'I his conclusion is confirmed by 
 the specific gravity of ozone derived experimentally from the velocity of diffusion, and has 
 
OZONE. 87 
 
 been found to be approximately 48 (O2 = 32), corresponding to the molecular formula 
 O3 (Graham). 
 
 A diminution in the volume of the gas does not occur in the action of 
 ozone upon oxidizable bodies like potassium iodide and mercury, although 
 all the ozone disappears. It would appear from this that, in oxidizing, 
 ozone only acts with one atom of oxygen, while the other two atoms form 
 free oxygen, which occupies the same volume as the ozone : 
 
 O3 -|- 2K.I = O2 ”b K2O -j- l2« 
 
 I vol. I vol. 
 
 As a consequence of this behavior ozone is also called oxidized oxy- 
 gen; i. e., free oxygen (O2), which has combined with an additional 
 oxygen atom. 
 
 Thermo -cheviical Deportment . — Compared with ordinary oxygen, ozone is an endo- 
 thermic compound. Heat is absorbed in its formation from oxygen : 
 
 (02,0) -f 36.2 Cal. = O3. 
 
 This explains why ozone is produced with so much difficulty, and why the addition of 
 considerable energy is necessary. This may be applied directly in the form of heat or 
 electricity (electric sparks, silent electric discharge), or it may be withdrawn from the 
 heat of formation of other exothermic compounds which are produced at the same time, 
 e. g., the formation of ozone by the oxidation of phosphorus to phosphorous acid. 
 
 Being an endothermic derivative, we readily perceive why ozone is so unstable, and 
 why it changes so readily to ordinary oxygen. When this occurs the oxygen acts as if 
 in the moment of formation (O3 = O2 + O ; see p. 53), and this explains why ozone 
 acts more powerfully than ordinary oxygen. And to this must be added that in oxida- 
 tions performed by ozone there are 36.2 Cal. more set free than in oxidations with ordi- 
 nary oxygen. 
 
 We observe, therefore, that the elementary substance oxygen occurs in 
 free condition in two different forms — allotropic modifications — ordinary 
 oxygen (Og) and ozone (O3). We shall learn later that very frequently 
 substances of the same elementary composition possess different physical 
 and chemical properties; such bodies are called isomerides and the phe- 
 nomenon isomerism (lVo<r, like ; p.ipoq, part). The isomerism of the ele- 
 ments was called allotropy {aXkorpimoc;, differently formed) by Berzelius. 
 It is accounted for (as in the case of oxygen and sulphur) by the differ- 
 ent number of atoms in the molecule. Alterations in the molecular 
 energy relations are obviously of great importance. The allotropy of 
 oxygen confirms the conclusion drawn from the gas densities, that the 
 molecules of the elements are composed of atoms. 
 
 We have already seen that ozone is absorbed not only by turpentine 
 and cinnamon oil, but also by other ethereal oils. These bodies are, 
 however, very slowly oxidized; the oxygen is contained in them in a 
 peculiar condition. In this form it acts upon some bodies like free 
 ozone; in other instances, the oxidizing action is rendered possible only 
 by peculiar substances which carry the oxygen. Spongy platinum, ferrous 
 
88 
 
 INORGANIC CHEMISTRY. 
 
 sulphate, and the blood-corpuscles are examples of this class. Thus, old 
 turpentine oil, containing absorbed ozone, only acts on |Kiper saturated 
 with starch-paste and potassium iodide, when a few drops of a ferrous 
 sulphate solution have been added to it. 
 
 Since ozone is formed when electricity acts ui)on air, and, indeed, 
 probably, in all oxidation and combustion processes, — and, further, j^otas- 
 sium iodidestarch-paper is blued when exposed to the air, — it was believed 
 that ozone was a constant constituent of atmosj^heric air (i-io milli- 
 grams in loo liters of air) ; according to recent investigations it is, how- 
 ever, probable that the imagined ozone reactions are frecpiently produced 
 by hydrogen peroxide, which is very similar in reaction t(j ozone (p. 103), 
 and is almost constantly in the air (Schdne). [Her. 13 (1880), 1503; 
 Engler and his co-workers, ]k*r. 30 (1897), 1669; 31 (1898), 3046.] 
 
 Antozone, which was regarded as a third peculiar modification of 
 oxygen, has been proved in some cases to be hydrogen peroxide and in 
 others to be oxygen in statu nascendi. 
 
 Oxygen is taken up at the ordinary temperature by comparatively few bodies. The 
 alkali metals, finely divided metals, compounds of sulphurous acid, phosphorus, certain 
 organic derivatives, etc., are .such “ auto-oxidizable ” bodies. In their spontaneous oxi- 
 dation the very remarkable phenomenon has been noticed that other substances, present 
 at the same time and in themselves not oxidizable, are also oxidized. It was therefore 
 said that oxygen was rendered active by the “ auto-oxidizable ” .substances. More recent 
 experiments have proved that the oxidation was occasioned by intermediate products rich 
 in oxygen (peroxides) which readily parted with their oxygen and acted as oxygen 
 carriers. [See Engler and Wild, etc.] 
 
 COMPOUNDS OF OXYGEN WITH HYDROGEN. 
 
 1. WATER. 
 
 Molecule : H2O = 18.02. 
 
 Water, the product of the union of hydrogen and oxygen, is produced 
 in many chemical processes, e. g., in the formation of salts from bases and 
 acids (p. 61). Cavendish was the first (1781) to produce water syntheti- 
 cally by the combustion of hydrogen in oxygen. Lavoisier a little later 
 showed by analysis that water contained hydrogen and oxygen (p. 40). 
 The first determination, although far from accurate, of the quantitative 
 composition of water also originated with Lavoisier (1783). The weight 
 ratio of hydrogen and oxygen in it was first determined more correctly 
 by Berzelius and Dulong (1820). Gay-Lussac conjointly with Humboldt 
 (1805) had shown that water was produced by the union of two volumes 
 of hydrogen with one volume of oxygen (p. 74). 
 
 Physical Properties . — It is obtained chemically pure by the distillation 
 of naturally occurring water, which always contains other matter dissolved 
 in it. It apj)cars in all three states of aggregation : in the liquid, gaseous 
 (steam), anil s(;lid (ice, snow). In a thin layer water is colorless, while 
 
WATER. 
 
 89 
 
 it is blue in layers of greater thickness. When water is cooled it con- 
 tracts and attains its greatest density at +3.98° C. The weight of a 
 cubic centimeter at -1-4° is taken as the unit of weight (= i gram). By 
 further cooling the water expands — the opposite of most other bodies ; its 
 volume becomes greater, while the specific gravity decreases. 
 
 The following table gives the volume and specific gravity of water for 
 different temperatures referred to water at 4° (see the investigations of 
 Thiesen, Scheel and Diesselhorn, Zeit. f. Instrumentenkunde, 17 (1897), 
 
 332): 
 
 Temperature. 
 
 Specific Volume. 
 
 Specific Gravity. 
 
 0° 
 
 1. 0001 324 
 
 0.9998676 
 
 2° 
 
 1.0000320 
 
 0.9999680 
 
 4 ° 
 
 1. 0000000 
 
 1. 0000000 
 
 6° 
 
 1.0000320 
 
 0.9999680 
 
 8° 
 
 I.0001241 
 
 0.9998759 
 
 10° 
 
 1.0002730 
 
 0.9997271 
 
 12° 
 
 1.0004756 
 
 0.9995246 
 
 14° 
 
 1.0007292 
 
 0.9992713 
 
 16° 
 
 I.0010314 
 
 0.9989697 
 
 18° 
 
 1. 00 1 2000 
 
 0.9986220 
 
 20° 
 
 1.0017728 
 
 0.9982303 
 
 22° 
 
 1.0022083 
 
 0.9977966 
 
 26° 
 
 1.0032006 
 
 0.9968097 
 
 36° 
 
 1.0063297 
 
 0.9937101 
 
 38° 
 
 1.0070584 
 
 0.992991 1 
 
 50° 
 
 1. 01 200 
 
 0.98813 
 
 100° 
 
 1.04327 
 
 0.95863 
 
 By cooling, water solidifies to ice. The melting point of ice is taken 
 as the zero of the centigrade and Reaumur’s thermometric scales. The 
 solidification-temperature of water cannot be used for this purpose, 
 because water at rest can be chilled considerably below 0° without freez- 
 ing. The freezing point of ice, however, like all other solid bodies, is 
 constant at a definite pressure. As the latter rises the melting point of ice 
 falls. At a pressure of 1000 atmospheres ice melts at — 7° (Lord Kelvin). 
 Most other substances, unlike ice, melt with an increase in volume ; in 
 their case increase in pressure occasions a rise in the melting point 
 (Bunsen). 
 
 In the conversion of water into ice, a considerable expansion occurs : 
 TOO volumes of water at 0° yield 109 volumes of ice at 0° ; the specific 
 gravity of the latter is, therefore, 0.9173, referred to water at 4°. This 
 is why ice floats in water. Ice crystallizes in rhombohedral forms of the 
 hexagonal system as may be distinctly observed in snowflakes. 
 
 Different bodies require different quantities of heat to bring them to 
 the same temperature, 'fhe heat capacity of water is greater than that 
 of all other liquid or solid bodies. It is customary to take the quantity 
 of heat necessary to raise one part by weight of water from 15° to 
 16° C., as the unit of heat, or calorie (see p. 66). In the jiassage of a 
 8 
 
90 
 
 INORGANIC CHEMISTRY. 
 
 licjuid to the solid state lieat is always set free, while, on the other hand, 
 in the fusion of the solid heat is absorbed. 'I’he latent heat of fusion of 
 I kilogram of ice ecjuals 79 calories; this means that, for the melting of 
 the ice, a (juantity of heat is reejuired which is cajiable of raising 79 kilo- 
 grams of water from 15° to 16° (j), 92). 
 
 Water boils upon the a})plication of heat, and is converted into steam. 
 The boiling temperature, like that of all other licpiids, depends on the 
 pressure ; it is also influenced by the substances dissolved in it, although 
 the temperature of the vapors is constant (at a given pressure), d'he 
 tem})erature of the steam escaping from water at the ordinary pres- 
 sure of 760 mm. (45° latitude; sea-level) is 100° (= 80° Reaumur). At 
 680 mm. barometric pressure water boils at 96.9°, and at 800 mm. at 
 101.4°. 
 
 One volume of water at 100° C. yields 1696 volumes of vapor of the 
 same temperature. The specific gravity of steam equals 18.02 (03= 32) 
 or = 0.622 (air = i); see p. 79. One liter of aqueous vapor 
 
 weighs 0.590 gram at 100° under normal pressure. 
 
 The critical temperature of water (or its absolute boiling temperature, 
 p. 48) is +37o°j its critical volume 2.33, and its critical pressure 195.5 
 atmospheres, i. e., at 370° the tension of its vapor equals 195.5 atmos- 
 pheres, and above this temperature it can no longer exist as a liquid, but 
 only as a gas. 
 
 The vaporization of water, and of other liquids, occurs not only at the boiling point, 
 but also at lower temperatures. The tension of the vapor is measured by the height of 
 the column of mercury which holds it in equilibrio. 
 
 The following table gives the tension of aqueous vapor for various temperatures (ex- 
 pressed in mercury levels at 0°, 45° latitude, and sea-level) : 
 
 IMPERATURE. 
 
 Tension. 
 
 Temperature. 
 
 Tension. 
 
 —20° C. 
 
 0.93 mm. 
 
 20° C. 
 
 17.4 mm. 
 
 — 10° c. 
 
 2. 15 mm. 
 
 40° C. 
 
 54.9 mm. 
 
 0° C. 
 
 4. 6 mm. 
 
 60° C. 
 
 148.9 mm. 
 
 -f 10° c. 
 
 9. 1 mm. 
 
 80° C. 
 
 355.4 mm. 
 
 +15° C. 
 
 1 2. 7 mm. 
 
 100° C. 
 
 760.0 mm. 
 
 Moist gases, therefore, occupy a larger volume than those which are dry. (See Air : 
 measurement of gases). 
 
 The latent heat of the evaporation of a kilogram of water equals 536.4 
 heat units at 100° C. ; i. e., for the conversion of one kilogram of water 
 at 100° C. into vapor of the same temjierature a quantity of heat will be 
 absorbed cajtable of raising 536.4 parts of water from 15° to 16°. 
 
 Jn consequence of the evaporation of water, the gases separating 
 from an atpieous solution are always moist. To dry the same, conduct 
 them over substances which will take up the moisture, e. s^., calcium 
 chloride, stick ])otash, sulplmric acid, ])hosphoric anhydride (compare 
 j). 42). Many solids al)stract moisture from the air without chemically 
 uniting with it; to dry these let them stand in an enclosed space 
 over suli>huric acid, calcium chloride, or phosphorus pentoxide (desic- 
 cators). 
 
WATER. 
 
 9 
 
 Natural Wafers. — As water dissolves many solid, liquid, and gaseous 
 compounds, all naturally occurring waters contain foreign admixtures. 
 The purest natural water is rain- or snow-water; it contains about 3 per 
 cent, by volume of gases (oxygen, nitrogen, argon, and carbon dioxide), 
 and traces of solids (the ammonium salts of nitrous and nitric acids). 
 If water that has been standing exposed to the air be heated, the dis- 
 solved gases escape in bubbles. 
 
 River and spring waters contain solid constituents in widely varying 
 amounts. Water having much lime and gypsum present in it is ordinarily 
 known as hard, in distinction from soft water, which contains less lime 
 (see Calcium Carbonate). On boiling lime-waters, most of the impurity 
 deposits out. Spring water generally contains in addition larger quan- 
 tities of carbon dioxide, which impart a refreshing and enlivening taste 
 to it. Spring waters holding considerable quantities of solid constitu- 
 ents, or exhibiting special healing properties, are called mineral waters. 
 These are distinguished as saline waters (containing sodium chloride), 
 magnesian waters, sulphur waters (hydrogen sulphide), acidulated waters 
 (saturated with carbon dioxide), chalybeate waters (containing iron), 
 and others. 
 
 Sea-water contains 3.5 per cent, of salts, chiefly sodium chloride (2.7 
 per cent.). 
 
 To purify the natural waters they are filtered (for the removal of 
 mechanical admixtures) and, for chemical and pharmaceutical purposes, 
 distilled {distilled water) in apparatus of varying form. 
 
 Chemical Properties of Water. — Water is a neutral substance, i. e., it 
 possesses neither acid nor basic properties. As we have already observed 
 (p. 61), it forms bases with basic oxides and acids with acid- forming 
 oxides. 
 
 Despite the fact that the affinity of hydrogen for oxygen is may great, 
 water may be decomposed by many substances. At ordinary tempera- 
 tures, metals like potassium, sodium and calcium decompose it, with 
 liberation of hydrogen : 
 
 2H2O -f Kj = 2KOH -f H2. 
 
 Other metals do not decompose it, except at elevated temperatures. 
 Steam conducted over ignited iron gives its oxygen to the latter, form- 
 ing ferroso-ferric oxide, while hydrogen is set free (pp. 40, 94) : 
 
 3 Fe + 4H2O = FcgCh -f 4H2. 
 
 This is a reversible reaction. Many of these oxides (even sodium oxide, according 
 to Beketoff ) are again reduced by hydrogen at more elevated temperatures : 
 
 1^^304 + 411.! = Fe3 + 4H2O. 
 
 (See pp. 52, 68, 95.) 
 
 Chlorine decomposes water in the sunlight; the decomposition is more 
 rapid when the vapors are conducted through heated tubes (j). 52) : 
 
 2IT./) f 2CI2 = 4TTCI + O2. 
 
 Many chemical reactions are only comj)leted in the presence of moisture. Thus, the 
 metals are only oxidi/.cd at the ordinary temperature when both oxygen and moisture are 
 
92 
 
 INOKHANIC CHKMISTRY. 
 
 l)resent. Iron docs not rust in jicrfeclly dry air. Absolutely dry oxyjfcn docs not act 
 upon tlie sul)stanccs which it ordinarily attacks with great energy. I’liosplionis, carbon 
 and carbon monoxide do not burn, or at least but feebly, in perfccily dry oxygen. I try 
 liydrogen chloride does not turn blue litmus red, etc. [Dixon, l>er. 19 (1886), Kef. 157 ; 
 Baker, ibid. 18 (1885), Ref. 426; Lothar Meyer, I 5 er. 19 (1886), 1099; K. Otto, l>er. 
 26 (1893), II, 2050; Hughes, ibid, iv, 863; Veley, Her. 29 (1896), i, 577; also Out- 
 mann, Ann. Chem. 299 (1898), 267 ; see also p. loi.] 
 
 Electrolysis of Water. — The electric current, acting upon water acidu- 
 lated with sulpliuric acid, decomposes it ajiparently directly into its ele- 
 ments. Hydrogen collects at the negative pole — the kathode, while 
 oxygen ajtpears at the positive jtole — the anode. The volume of the 
 hydrogen is nearly twice that of the oxygen (pp. 76, 98). 
 
 The electrolytic decomposition of water is more complex than is ordinarily supposed, 
 as perfectly pure water is not ca{)able of conducting the current, and is conseciuently not 
 decomposed by it. It is rather the added sulphuric acid which suffers the decomposition 
 and by means of the water is always reformed (compare the Fdectrolysis of Salts). 
 Hydrogen and oxygen are merely the end-products of this change. In addition to 
 oxygen, about one per cent, of ozone is produced ; further, sulphur hejjtoxide (persid- 
 phuric acid) and liydrogen jieroxide are formed at the anode. Some hydrogen peroxide 
 is produced at the negative pole (the kathode) as the result of the union of nascent 
 hydrogen with the dissolved oxygen (p. loi). 
 
 Therfno-chemical Deportme 7 it. — Water is formed from its elements with 
 the liberation of much heat. 57.2 Cal. are disengaged in the union of 
 2 grams of hydrogen with 16 grams of oxygen to produce aqueous vajior 
 of 100° (H.^, O — vapor). In the condensation of the steam to water of 
 100° 9.63 Cal. are liberated (— 18 X o-SS^); this is the latent heat of 
 evaporation (see p. 89). And again, in the cooling of the water through 
 every 1° C., Cal. more escape; consequently in cooling from 100° 
 to 16° there would be a liberation of 84. ^ = 1.5 Cal. (see p 66). 
 
 So that in the production of the molecular weight (18 grams) of water of 
 16° temperature from its elements there is a total disengagement of 68.3 
 large calories : 
 
 (Hj, O — vapor) = 57.2 Cal. (Hg, O — liquid) = 68.3 Cal. 
 
 The decomposition of water at the ordinary temperature by a metal only takes place 
 if the heat of formation of the oxide is greater than that of the water. Sixteen grams of 
 oxygen in their union with hydrogen liberate 68.3 Cal., with sodium 100 Cal., and with 
 cop])er but 38 Cal. The decomposition of water by sodium is an exothermic reaction, 
 while that by copper is endothermic (see p. 94). 
 
 * Faraday introduced the following terms, which have been universally adopted : The 
 metal wires or j)lates by which the current enters and pas.ses from electrolytes are called 
 the electrodes {bbbr, way) ; the electrode by which the current enters is the and the 
 
 other through which the ])ositive electricity has its exit and by which the negative elec- 
 tricity enters is the kathode. That part of the electrolyte jjassing to the anode where it is 
 se])aratcd or deposited is the anion ( that which migrates upward — opposite to the current 
 of positive electricity) while the ])ortion going to tlie kathode and separating there is the 
 kation (migrnling downwards — with the current of positive electricity). Both are called 
 ions (/bora, wandering, migrating). As a rule the kations in a compound are replaceable 
 by hydrogen ; tlx.* anions are simple or compound halogenides (Cl, SO^) (compare the 
 F.Iectrolysis of Salts and b'.lect roly tic I )is.sociation). 
 
WATER. 
 
 93 
 
 Dissociation of Water — Water, like otlier clieiiiical compounds, is 
 broken down into its elements by heat. This was first observed upon 
 pouring molten platinum into water, when bubbles of oxyhydrogen gas 
 appeared (Grove). This decomposition of water was first ascribed to a 
 catalytic action of the platinum. Sainte-Claire Deville was the first to 
 carefully investigate and explain the decomjiosition phenomena induced 
 by heat, thus disclosing one of the most important chajiters of theoretical 
 chemistry. He proved that a decomi)osition (dissociation) like the 
 preceding did not take place suddenly, l)ut gradually; that it advanced 
 regularly with increasing temperature, and was limited by an opposing 
 combination-tendency on the part of the components. The temperature 
 at which the decomposition is half finished is usually designated as the 
 temperature of decomposition. 
 
 Deville illustrated the decomposition of water by the following experi- 
 ment: Pass aqueous vapor through a porous clay tube, a, cemented into 
 a wider non-permeable porcelain tube heated to a white heat in an oven 
 
 (Fig. 45 ).* The water suffers partial decomposition, the lighter hydrogen, 
 which passes through the inner tube into the porcelain tube more rapidly 
 than the oxygen, escapes through the gas tube b. The oxygen escapes 
 mainly through the inner tube at a. A part of the same diffuses simul- 
 taneously with the hydrogen and reunites with the latter. To avoid this, 
 conduct a stream of carbon dioxide through the wider porcelain tube; this 
 will carry out the hydrogen with it. The carbon dioxide will be absorbed 
 by the alkali solution in the collecting vessel, and oxyhydrogen gas be 
 found in the cylinder. The decomposition of the water commences at 
 about 1000°, and is half finished at about 2500°. The quantity of gas 
 liberated in equal periods rises successively with the temperature. 
 
 Many other compounds, as carbon dioxide, hydrogen chloride, iodine 
 (P- 55)) ammonium chloride, ])hosphorus pentachloride, etc., are simi- 
 
 * A tube of platinum may be well substitutecl for the porous clay tube ; at a red heat 
 it permits the passage of hydrogen, but not that of oxygen (p. 45), 
 
94 
 
 INORGANIC CHKMIS'IRY. 
 
 larly dissociaLcd by heal. These are all exothermic compounds, absorb- 
 ing energy in their decomposition, and are therefore decomposed but 
 gradually, dej)ending upon the amount of energy imparted to them. In 
 these instances heat ojiposes the affinity of the various components, so 
 that if the temperature be lowered there will occur a jiartial reunion of 
 the same. The siilitting up of the endothermic comjioiinds is entirely 
 different, e. g., that of ])otassium chlorate, KC 1 (),„ into potassium chloride 
 and oxygen, of ammonium nitrite, NH^NO^, into water and nitrogen, of 
 nitrogen chloride into chlorine and nitrogen, etc. Heat is set free in 
 the decomposition of these compounds. Any added or external heat 
 only incites or brings on the decomjiosition and overcomes the chemical 
 affinity. Under some conditions there are accompanying explosions; 
 there is no reunion of the components on lowering the temperature. 
 
 The ex])lanation of the dissociation plienoinena is found in tlie kinetic tlieory of gases 
 {Kivt/aii', motion). According to it the heat-motion of the molecules of a gas is motiem 
 in direct lines, jjrogressive and with uniform velocity. As the temperature rises the 
 velocity of this motion increases. Its energy for the molecules of different gases at any 
 given temperature is the same. Heavy molecules move corres[)ondingly slower than the 
 light molecules. The atoms, too, forming the molecule, have their own peculiar motions, 
 which become more energetic with the rise in temperature. By the peculiar motions of 
 the atoms — in which the center of gravity of the molecule is not concerned — the internal 
 arrangement of the molecule is affected ; it akso opposes the action of affinity. Therefore, 
 just as soon as, by augmented temperature, the centrifugally active energy of atomic 
 motion equals or exceeds the affinity, the breaking down of the molecule takes place. 
 Further, as a consequence of irregular collision, the molecules do not all possess the 
 .same velocity at a given temperature ; some move more rapidly, others .slower, than the 
 main portion ; the former are warmer than the latter. Only the sum of the existing forces 
 of ail the molecules is a constant quantity at every temperature. 
 
 The molecules first decomposed are those moving more rapidly and having a temperature 
 beyond the average temperature. Their number increases with the temperature. Hence, 
 it follows that dissociation also increases with rise of temperature. Compare in this 
 connection W. Nernst and A. Schonflies, Einfiihrung in die math. Behandlung der 
 Naturwissenschaften, 2 Aufl. (1898). 
 
 The dissociation of solids, which when heated develop gaseous ingre- 
 dients, is very instructive and remarkable — for example, the decomposi- 
 tion of calcium carbonate, CaCOg, into calcium oxide and carbon 
 dioxide, of sodium and potassium hydrides (K^H2) into their elements, 
 etc. These indicate that the dissociation is not only dependent upon the 
 temperature but also upon external pressure, and that for every tempera- 
 ture there is a corresponding definite tension of dissociation — a pressure 
 below which the decomposition will not occur. For further particulars 
 on this point see Potassium Hydride and Calcium Carbonate. 
 
 Dissociation, i. e., the partial decomposition increasing with the tem- 
 ])erature, ex])lains many chemical processes and jffienomena. Thus it 
 accounts for the abnormal vapor densities, which apparently contradicted 
 the law of tlie efjnal number of molecules being ])resent in equal volumes 
 of gases (j). 79) ; all variations from it are always due to the breaking down 
 of more complex molecules (see Sulphur, p. 106). The observed vapor 
 density affords a cine to the magnitude of the dissociation. The mass 
 action in reversible (inverse) chemical reactions is afforded a simple ex- 
 })lanation by dissociation. We have already said that iron raised to a red 
 
THE QUANTITATIVE COMPOSITION OF WATER. 95 
 
 lieat decomposed water with the separation of hydrogen and the produc- 
 tion of ferrous-ferric oxide : 
 
 3Fe + 4H2O = Fe30^ + 4H,. 
 
 On conducting hydrogen over ignited iron oxides the opposite process 
 occurs; the oxygen compound of the iron is reduced and water and iron 
 are formed : 
 
 Fep^ + 4H2 = 3 Fe + 4H2O. 
 
 In the first instance the excess of water acts. Some of its molecules are dissociated ; 
 oxygen combines with iron, while the liberated hydrogen is carried away by the excess 
 of steam. In the second case, we can suppose that some of the hydrogen molecules are 
 dissociated, the free hydrogen atoms withdraw oxygen from the iron oxide and form water 
 with it, which is removed by the excess of hydrogen, and thus prevented from acting on 
 the reduced iron. 
 
 If, however, iron and steam be heated in an enclosed space for every temperature 
 above 150°, where action begins, there will occur a state of equilibrium : ferrous- ferric 
 oxide, iron, water, and hydrogen will be present together in a ratio which does not alter 
 for any definite temperature, i. e., in a definite period of time as much iron will be pro- 
 duced from the ferrous-ferric oxide by the hydrogen as is oxidized by the water to ferrous- 
 ferric oxide. This may be expressed thus : 
 
 FegO, + 4H2 3 Fe -f 4H2O. 
 
 (See pp. 52, 68, 91.) 
 
 Similarly, hydrogen chloride is decomposed at a red heat by oxygen 
 with the formation of steam and chlorine gas, while in turn steam and 
 chlorine gas are transposed into hydrogen chloride and oxygen. As 
 hydrogen chloride is more stable than water and only dissociated at 
 high temperatures, its formation is the predominating process at a red 
 heat. Here, however, in the course of time, a state of equilibrium 
 appears for every given temperature and pressure. 
 
 The investigation of the dependence of this state of equilibrium upon temperature, 
 pressure, duration of action, the quantities, concentration, and solubility-proportions of the 
 reacting substances, has become of the greatest importance for theoretical chemistry, for 
 the reason that here for the first time chemical phenomena are capable of a mathematical 
 treatment. The fundamental investigations, both theoretical and experimental, made 
 upon the law of mass-action, which will not be entered upon here, are due to the Nor- 
 wegians C. M. Guldberg and P. Waage [Etudes sur les affinites chimiques, Progr. de 
 I’Universite Christiania, 1867. See also Jr. prakt. Ch. [2] 19 (1879), 69]. For study 
 use the works, referred to on p. 49, by Lothar Meyer ; also the Theoretical Chemistry 
 by W. Nernst, 2 Aufl., 1898. 
 
 THE QUANTITATIVE COMPOSITION OF WATER. 
 
 The weight and volume relations by which hydrogen and oxygen com- 
 bine to form water constitute the most important basis for the determina- 
 tion of the ratio of the atomic weights of these two elements to each other. 
 To determine this ratio with such accuracy that no question can longer 
 exist has been the purpose of innumerable investigations since the time of 
 Dalton. Numerous difficulties had to be surmounted. They were found 
 particularly when seeking to obtain the gases perfectly dry, in weighing 
 
96 
 
 INORGANIC CIlEMIS'l'kV. 
 
 ihcni, in measuring them, and in endeavoring to combine them, d’he 
 judgment, care, and untiring j)atience reviuisite for tlie solution of tin’s aj)- 
 parently simple problem may be gathered from the magnificent researches, 
 already referred to, of the American Edward W. Morley (pj). 44, 73). 
 
 d'he composition of water has been chiefly determined by its synthesis 
 which can be followed quantitatively in its details, as both oxygen and 
 hydrogen, as well as the water jiroduced from them, are weighed. 'J’he 
 hydrogen is liberated either by the electrolysis of dilute suliihuric acid or 
 sodium hydroxide (Thomsen, ]>. 42), and according to Reiser, is best 
 weighed as jialladium hydride (p. 46). 'I'he oxygen is obtained by heat- 
 ing potassium chlorate. Chemists have frefiuently been satisfied, instead 
 of weighing the three bodies, hydrogen, oxygen, and water, to merely 
 weigh two of them. This was true of Terzelius and Diilong (1819), Erd- 
 
 Fig. 46. 
 
 mann and Marchand (1842), and Dumas (1843), who proceeded according 
 to a method which has become classic: hydrogen was conducted over 
 ignited copper oxide, which was reduced to metallic copper, with the 
 formation of water: 
 
 CuO + H 2 = Cu -f H,0. 
 
 Cupric oxide. Copper. 
 
 Heat a weighed portion of cupric oxide (containing a definite amount 
 of oxygen) in a stream of ])ure, dry hydrogen, and weigh the quantity of 
 water obtained. The operation can be executed in the apparatus repre- 
 sented in Fig. 46. It was em[)loyed by Dumas. The hydrogen gener- 
 ated in the flask A is washed in B, and then dried in the tubes C, D, and 
 /f, which contain substances that will absorb water. The bulb tube F, 
 of difficultly fusible glass, contains a weighed amount of cupric oxide, and 
 is heated with a lamp. The water which forms collects in the bulb G, 
 and is com])letely absorbed in the tube H. Hydrogen is led over the 
 cupric oxide until it is reduced to red metallic copper, then allowed to 
 
THE MOLECULAR FORMULA OF WATER. 
 
 97 
 
 cool, when 7 ^ is weighed alone and G and iT" together. The loss in weight 
 of expresses the quantity of oxygen which has combined with hydrogen 
 to produce water. The increase in weight of and ZTgives the quantity 
 of water that was formed. The difference shows the amount of hydrogen 
 in water. 
 
 Cooke and Richards modified this method in that they conducted a 
 weighed amount of hydrogen, by means of a dry air current, over the 
 ignited copper oxide, and then determined the weight of the water 
 which resulted. 
 
 The composition of water by weight can also be ascertained from the 
 ratio of the specific gravities of the two gases and the volume ratio 
 according to which they combine. The most recent researches (Scott, 
 Leduc, Morley) indicate that 2.0027 volumes of hydrogen unite with i 
 volume of oxygen to form water. As one liter of oxygen weighs 1.4291 
 gram and one liter of hydrogen 0.08988 gram, their weight ratio in water 
 would be 
 
 I.4291 : 2.0027 • 0.08988 = 16 : 2.0153. 
 
 Thus we ascertain that in 100 parts of water, by weight, there are 
 
 1 1. 2 parts Hydrogen 
 88.8 “ Oxygen 
 
 100.0 “ Water. 
 
 Ostwald in his Lehrbuch der allgemeinen Chemie, 2 Aufl. (1891), i, 43, gives a 
 review of the history of these important investigations. See also the researches of 
 Morley and of Thomsen (p. 44). 
 
 THE MOLECULAR FORMULA OF WATER. ATOMIC WEIGHTS OF HYDROGEN 
 
 AND OXYGEN. 
 
 The molecular weight of water according to Avogadro’s law is 18.02 
 and that of hydrogen is 2.02 if the molecular weight of oxygen be placed 
 at 32. As to the number of atoms combined to a molecule we can con- 
 clude from what has been previously said that the molecules of hydro- 
 gen, oxygen, chlorine and hydrogen chloride contain at least two atoms 
 each (p. 75). The facts about to be presented make it probable that 
 these molecules do not consist of more than two atoms and that the 
 molecular formula H^O represents water (p. 73). 
 
 One volume of acpieous vapor contains only one-half as much oxygen 
 as an equal volume of oxygen gas, and in an equal volume of any gaseous 
 oxygen compound less oxygen has never been observed. Hence it may 
 fairly be concluded that there is only one atom of oxygen present in the 
 molecule of water, and that the oxygen molecule itself consists of two atoms. 
 
 Again, one volume of aqueous vapor contains just as much hydrogen as 
 an equal volume of hydrogen gas, and twice as much as hydrogen chlo- 
 ride gas. Since, therefore, in equal volumes of other gaseous bodies less 
 hydrogen and less chlorine, than in hydrogen chloride, have never been 
 observed, it maybe assumed that the molecule of hydrogen chloride con- 
 tains one atom each of hydrogen and chlorine, and further that the 
 9 
 
98 
 
 INORGANIC CHKMISTKY. 
 
 hydrogen molecule consists of two atoms of hydrogen, the water molecule 
 of one atom of oxygen and two atoms of hydrogen. In this manner 
 we arrive at the molecular and atomic weights given on p. 79; whereby 
 the atomic weight of hydrogen 1.008 is reduced to i.oi. 
 
 It seems practicable, for considerations like those given, to introduce several new 
 terms. As the sj)ecific volume Vs of a gas we have (iesignated that volume which 1 
 gram of the gas would occupy under normal conditions (p. 45). d'he quantity of a sub- 
 stance expressed by the molecular weight in grams slnndd he termed the g?'aiii-i?iolecule 
 or the mol. One mol of oxygen would then under normal conditions he 32 grams. 'I’he 
 volume Vm, which would he occupied by i mol of a gas of the molecular weight M under 
 normal conditions — the violectdar volmne — would he the sj)ecific volume multiplied by 
 the molecular weight : 
 
 Vm = Vs . M = 22.4 liters. 
 
 The mol-volume is the same for all gases : a new expression for the so-called Avo- 
 gadro law. It equals 22.4 liters ; i. e., 2.02 grams of hydrogen 70.9 grams of 
 
 chlorine (Cl.^), 36.46 grams of hydrogen chloride (IlCl), 32 grams of oxygen (02)> i^-02 
 grams of acjueous vapor occupy under normal conditions a volume of 22.4 liters. This 
 is only approximately correct because of the deviations of gases from the re(|uirements of 
 their laws. What has been said in the preceding lines may he expressed thus : .Since in 
 22.4 liters of a homogeneous gaseous compound never less than i.oi grams of hydrogen, 
 35.45 grams of chlorine, 16 grams of oxygen, etc., have been discovered, it may be con- 
 cluded that these express their atomic values. 
 
 After having thus derived the molecular formula of water, and the atomic 
 weights of oxygen and of hydrogen, we deduce the following conclusions; 
 
 I. Sixteen ytarts by weight of oxygen occupy the same volume as i.oi 
 
 parts by weight of hydrogen; since 16 partsof 
 the former unite with 2.02 parts of the lat- 
 ter in the production of water, one volume of 
 oxygen must combine with two volumes of 
 hydrogen. 
 
 2. In equal volumes we have an equal 
 number of molecules ; n molecules of oxygen 
 (Oj) unite therefore with 2n molecules of 
 hydrogen (H^) ; the same yield 2n mole- 
 cules of water; consequently two volumes of 
 aqueous vapor : 
 
 2nH2 -p n02 = 2nH20. 
 
 2 vols. I vol. 2 vols. 
 
 According to the above, two 7 'o/umes of hy- 
 drogen and one vohi 77 ie of oxygen co 7 iden 5 e i 7 i 
 their u 7 iio 7 i to tivo vohwies of aqueous vapor. 
 
 'These conclusions are confirmed by the 
 following exiteriments : 
 
 I. Wlien water is decomposed by the electric 
 current in a voltameter, or, more suitably, in Hof- 
 mann’s apparatus (Fig. 38, j). 77b it will be fouiul 
 that the volume of the separated hydrogen is double 
 that of the oxygen. This can also be proved syn- 
 thetically. Introduce I volume of oxygen and 2 
 volumes of hydrogen into a eudiometer lube fdlcd with mercury (see Air), and let the 
 
HYDROGEN PEROXIDE. 
 
 99 
 
 electric spark pass through the mixture. This will unite the two gases, a small quantity 
 of water forming at the same time ; all the gas has disappeared, and the tube is com- 
 pletely filled with mercury. In place of the eudiometer, the apparatus (devised by Hof- 
 mann) pictured in P'lg. 47 may be advantageously employed in this experiment (and 
 also in many others). It consists of a U-shaped glass tube, one limb of which, open 
 above, is provided below with an exit tube. The other limb really represents a eudi- 
 ometer ; it is divided into cubic centimeters. It has two platinum wires fused into the 
 upper end, and provided with a stop-cock to admit and let out the gases and thus test them. 
 P'ill the tube to the stop-coek with mercury, and run into the eudiometer limb i volume 
 of oxygen and 2 volumes of hydrogen. The side exit tube serves to run out the 
 mercury to the same level in both tubes, so that the gases are always measured under 
 the same atmospheric pressure, and thus their volumes are easily compared. 
 
 2. To determine the volume of the resulting water existing as aqueous vapor it is 
 only necessary, after the explosion, to convert it by heat into steam. The apparatus 
 (Fig. 48) will answer for this purpose. This is essentially the same as that pictured in 
 Fig. 47, with the eudiometer limb closed above and surrounded by a wider tube. 
 Through the latter conduct the vapors of some liquid boiling above 100° C. (toluene, 
 xylene or aniline). These, then, pass through the envelope B, and are again condensed 
 
 Fig. 48. 
 
 in the spiral tube C. The quantities of hydrogen and oxygen used are heated to the same 
 temperature, their volume noted, the explosion produced, and the volume of the result- 
 ing aqueous vapor determined. From this it is Wnd that the volume of hydrogen is ^ 
 of the volume of the gas mixture ; and 3 volumes of oxyhydrogen gas yield 2 volumes 
 of aqueous vapor. 
 
 2. HYDROGEN PEROXIDE. 
 
 H 2 O 2 = 34-02. 
 
 In addition to water, oxygen forms another compound with hydrogen, 
 known as hydrogen peroxide. It was discovered by Thenard in 1818. 
 It is produced by the action of dilute acids upon certain peroxides, such 
 as those of sodium, calcium and barium. It is usually obtained by the 
 action of hydrochloric acid upon barium peroxide : 
 
 BaOj -j- 2IICI = RaCh + HjOj. 
 
 Barium Barium 
 
 peroxide. chloride. 
 
lOO 
 
 INORGANIC CHEMISTRY. 
 
 Barium peroxide made to a paste with a little water (better, the hydrate — see Barium) 
 is introduced gradually, in small (juantities, into cold hydrcK'hloric acid, diluted with 3 
 volumes of water. Hydrogen peroxide and barium chloride result ; both are soluble in 
 water. To remove the second from the solution, add to the latter a .solution of silver sul- 
 phate as long as a precipitate is formed. Two insoluble compounds, barium sulphate and 
 silver chloride, are produced by this reagent : 
 
 BaCl, f Ag2SO, = BaSO, -f 2AgCl. 
 
 Remove the precipitate by filtration and concentrate the aqueous .solution. It now con- 
 tains only hydrogen peroxide. 
 
 In making the peroxide, carbon dioxide may be allowed to act on 
 barium peroxide susitendcd in water ; 
 
 Ba02 CO 2 + II 2 O BaCOg -f 
 
 The insoluble barium carbonate is filtered off and the filtrate concen- 
 trated. 
 
 Hydrogen jieroxide is most practically obtained by adding moist 
 barium hydrated peroxide (see llarium) to cold dilute sulphuric acid, d'he 
 reaction occurs according to the following equation : 
 
 BaOj + 112^^4 = BaSO^ -|- II 2 O 2 . 
 
 When the acid is almost neutralized, filter the solution, and from the 
 filtrate carefully precipitate the slight quantity of free sulphuric acid with 
 a dilute barium hydroxide solution, then concentrate the liquid. Dry 
 commercial peroxide of barium is not applicable for the above. A dilute 
 solution of hydrogen peroxide is very readily prepared, if sodium peroxide 
 (obtainable by fusing sodium in the air) is added to dilute tartaric acid. 
 
 A 45 per cent, solution can be readily made by evaporating the dilute aqueous hydrogen 
 peroxide solutions on a water-bath at a temperature not exceeding 70°. The loss will be 
 very slight. By extracting such a solution with ether and allowing the latter to evaporate 
 a more concentrated product can be obtained and finally an almost anhydrous peroxide 
 will remain. 
 
 The simplest means of preparing very pure hydrogen peroxide (99.7 per 
 cent.) consists in distilling the aqueous solution (for example, the commer- 
 cial 3 per cent, solutions) under reduced pressure. At 84-85°, under 
 68 mm. pressure, or at 69.2° and 26 mm. pressure, almost perfectly anhy- 
 drous hydrogen peroxide distils over [Wolffenstein, Ber. 27 (1894), ii, 
 3307; Spring, Zeit. f. anorg. Chem. 8 (1895), 424; Briihl, Ber. 28 
 
 (1895), 2847-] 
 
 Besides these decompositions of metallic peroxides, other methods exist for preparing 
 hydrogen peroxide (in small quantity). Thus, it ari.ses frequently in slow oxidations, when 
 ozone is also ])roduced. If phosphorus, covered with water, be allowed to oxidize in the air, 
 hydrogen peroxide will be found in the water, and the surrounding air will contain ozone 
 fp. 85). Or, if a flask filled with air be shaken with zinc and water or dilute sulphuric 
 acid, hydrogen peroxide will be ])rodiiccd. It is destroyed again by the prolonged action of 
 the zinc. If zinc amalgam, in (he presence of milk of lime, be .shaken with caustic potash, 
 calcium ])eroxidc, OaO,^, will be at once ])recii)itated. Copper, lead, iron, and other heavy 
 metals do the same when agitated with more or less dilute sulphuric acid, and we find the 
 
HYDROGEN PEROXIDE. 
 
 lOI 
 
 same result by the oxidation of many organic substances, e. g. , pyrogallic acid and tannin on 
 exposure to the air. In combustions, if the flame be cooled suddenly, we have formed, very 
 often, slight quantities of hydrogen peroxide (and ozone), e. g., in bringing a hydrogen or 
 carbon monoxide flame in contact with water (Traube, Her. 26 (1893), ii, 1471, 14761. 
 The explanation offered for this formation of hydrogen peroxide (and ozone) is, that in 
 the oxidations, the oxygen molecules are torn asunder, and the nascent oxygen atoms 
 oxidize the water to a slight degree to hydrogen peroxide, and oxygen to ozone. The 
 rare occurrence of ozone is due either to its difficult formation, or to the fact that it is 
 readily decomposed by the reacting bodies (zinc, etc.). This is also the case wilh hydro- 
 gen peroxide. The appearance of hydrogen peroxide in the oxidation of plu)Sj)horus 
 seems to prove that it can be formed by the oxidation of wate 7 \ This seems to be con- 
 firmed by its production on shaking turpentine oil with water and air, or if ozone be con- 
 ducted into ether, and the ozonized product shaken with water. It appears probable, 
 however, that in some oxidation reactions, the formation of the hydrogen peroxide is a 
 consequence of the reduction of oxygen (Hoppe-Seyler and Traube). It may, for ex- 
 ample, be assumed that when zinc (lead, iron) is shaken with air and water (or dilute 
 sulphuric acid), the latter is decomposed in such a manner that the hydroxyl group com- 
 bines with the zinc to hydroxide, and the liberated hydrogen then yields hydrogen 
 peroxide with oxygen : 
 
 Zn + 2OHH + 02 = Zn(0H)2 + H2O2. 
 
 Zn + H2O2 = Zn (011)2. 
 
 A confirmation of this supposition is found in the electrolysis of water, where we discover 
 hydrogen peroxide appearing at the negative pole (where hydrogen is found) if air or 
 oxygen be conducted through the solution, 2H + 02= H2O2. It is verified, too, in 
 the production of hydrogen peroxide upon shaking palladium hydride with water and air 
 (Traube) : 
 
 Pd4H2 + 02 = 4Pd + H 2 O 2 . 
 
 The excess of palladium hydride further decomposes the peroxide which was formed : 
 
 Pd,H 2 + H2O2 = 4 Pd + 2H2O. 
 
 In all these examples we can explain the formation of the peroxide by the action of 
 nascent hydrogen upon oxygen. It is, however, not true that hydrogen peroxide is 
 formed only by the reduction of molecular oxygen (see above). An evidence of this is 
 the fact that hydrogen peroxide is produced at the anode in the electrolysis of sulphuric 
 acid ; its appearance here is due to the decomposition of the persulphuric acid (H2S20g) 
 (Richarz). 
 
 The production of hydrogen peroxide in oxidations has led to the assumption that all 
 oxidations are conditioned by the transitory formation of hydrogen peroxide — Traube’ s oxi- 
 dation theory. It was supposed that the proof of this could be found in the circumstance 
 that various oxidations, e. g., the union of carbon monoxide with oxygen to form carbon 
 dioxide, could only occur with ease in the presence of aqueous vapor (Dixon) : CO + 2OIIH 
 -p O2 = CO2 + II2O + 112^^2- More careful investigations have, however, demonstrated 
 that the presence of moisture is not absolutely e.ssential in oxidations. Carbon monoxide 
 and oxygen also combine to carbon dioxide when perfectly dry if the temperature be suf- 
 ficiently high. Their union in the presence of moisture is due solely to the fact that the 
 following transpositions, CO + H2O = CO2 + Ikj and 2II2 + O2 = 2H2O, take place 
 more readily and at a lower temperature than the direct union of carbon monoxide with 
 oxygen : 2CO + 02 = 2CO2 (Lothar Meyer ; see p. 92). 
 
 Hydrogen peroxide, concentrated as much as possible under diminished 
 pressure, is a colorless, syrupy liquid which is blue in color in a thick layer. 
 It does not congeal as readily as water. Its specific gravity at 15°, referred 
 to water at the same temperature, is 1.49. It vaporizes on exposure to the 
 air. It produces a burning sensation and white spots when applied to the 
 
102 
 
 INORGANIC CHEMISTRY. 
 
 skin, ll sometimes cx])lo(ies spontaneously with great violence. Its 
 aqueous, rather concentrated solutions react acid and have a bitter, 
 astringent taste. Hydrogen peroxide is not nearly so sensitive to heat 
 as was tormerly siq)posed, provided that all alkaline-reacting compounds 
 and chemically active solids are absent. It is miscible in all proportions 
 with water. Mixtures of the conqiosition H./).^ -f H/) and -j- 2 H 2 O 
 solidify below — 20°. Upon warming solutions containing more than 
 40 per cent, of hydrogen j)eroxide, the latter volatilizes in large amounts 
 with the aqueous vapor and for the most part without decomposition 
 (Wolffenstein, p. 100). 
 
 In concentrated solutions, it is very unstable, and easily decomposed 
 with liberation of oxygen ; in more dilute acidulated solutions it may be 
 preserved longer. According to Spring, substances such as ether and 
 alcohol, which reduce the surface tension of the liipiid, tend to make its 
 solution more durable, whereas bodies like caustic potash, etc., which 
 increase the surface tension, hasten the breaking down of the hydrogen 
 peroxide. In consequence of its ready decomposition, hydrogen peroxide 
 oxidizes powerfully, since oxygen appears statu nascendi {\). 53). It 
 converts selenium, chromium, and arsenic into their corresponding acids; 
 sulphides are changed to sulphates (PbS to PbSOi) ; from lead acetate 
 solutions the peroxide is precipitated, but is again decolorized by the 
 excess of hydrogen peroxide. Organic dyestuffs are decolorized and 
 decomposed. From hydrogen sulphide, sulphur, from hydrogen chloride 
 and iodide, chlorine and iodine are set free : 
 
 H2O2 + 2HI = 2H2O + Ij. 
 
 Sulphurous acid is oxidized to sulphuric acid : 
 
 H 2 SO 3 + 11 , 0 , = H^SO, + H^O. 
 
 Thus hydrogen peroxide acts in a manner analogous to ozone ; in both 
 there exists a loosely combined atom of oxygen, which can readily be 
 transferred to other bodies. 
 
 Hydrogen peroxide acts very slowly upon a neutral potassium iodide 
 solution, while ozone separates iodine at once ; but if platinum-black, fer- 
 rous sulphate, or blood-corpuscles (see p. 88), be added to the solution, 
 iodine immediately separates, and colors added starch-paste a deep 
 blue. 
 
 In all these cases the action of hydrogen peroxide is oxidizing. Some 
 substances, on the other hand, are reduced by it, oxygen separating at 
 the same time ; this is true of certain unstable oxides, peroxides, and the 
 highest oxidations of some metals, like MuaO;, and CrO.,. Thus, argen- 
 tic, mercuric, and gold oxides are reduced to a metallic state with an 
 energetic evolution of oxygen : 
 
 Ag.p 4 II.A = 2Ag + II2O -f O2. 
 
 Lead peroxide is changed to lead oxide : 
 
 PhO^ -I- U,0, = PbO -f 1^0 -p O,. 
 
 In the jiresence of acids, the solution of potassium permanganate is 
 
REACTIONS FOR THE DETECTION OF HYDROGEN PEROXIDE. IO3 
 
 decolorized and changed to a manganous salt (see below). In the same 
 way chromic acid and its salts are altered to chromic oxide : 
 
 2Cr03 -f- 3H2O2 = Cr203 + 3H2O + 3O2. 
 
 Ozone and hydrogen peroxide gradually decompose into water and 
 oxygen : 
 
 O3 + 11.^02 = 02 + H2O Oj. 
 
 Chlorine in aqueous solution is oxidized to hypochlorous acid by 
 hydrogen peroxide, CI2 -f- H2O2 = 2HOCI, but is again reduced by an 
 excess of the latter : 
 
 HCIO + H2O2 = HCl + H ,0 -f O2. 
 
 Finally, hydrogen peroxide may be decomposed into water and oxygen 
 by many bodies, especially when the latter exist in a divided condition ; 
 they are not in the least altered. Gold, platinum, silver, manganese per- 
 oxide, and carbon act in this way. Such reactions, in which fhe 
 reacting substances undergo no perceptible changes, are designated cata- 
 lytic (xaraXooj, I open) (compare p. 102). In many cases these may be 
 explained by the previous formation of intermediate products, which 
 subsequently react upon one another. Thus, we can suppose that in the 
 action of silver and gold upon hydrogen peroxide oxides first result, but 
 these are afterward reduced by it in the manner mentioned above (see 
 p. 87). 
 
 REACTIONS FOR THE DETECTION OF HYDROGEN PEROXIDE. 
 
 Hydrogen peroxide decomposes potassium iodide very slowly ; in the presence of ferrous 
 sulphate, however, iodine separates at once, and is recognized by the blue color it yields 
 with starch-paste. In the same way guaiacum tincture, in the presence of ferrous sul- 
 phate, is immediately colored blue, and an indigo solution is decolorized. The most 
 characteristic test for the peroxide is the following : introduce hydrogen peroxide into a 
 chromic acid solution, add a little ether and shake thoroughly ; ths supernatant ethereal 
 layer will be colored blue (compare Chromic Acid). 
 
 A solution of titanic acid in sulphuric acid (diluted strongly with water), is also a deli- 
 cate reagent ; it gives an orange-yellow color with traces of hydrogen peroxide. This 
 reaction can also be applied in the presence of persulphuric acid. 
 
 Hydrogen peroxide is determined quantitatively by oxidation with potassium perman- 
 ganate (see Manganese). The latter is added to the solution, acidified with sulphuric 
 acid, until a permanent coloration occurs. The reaction proceeds according to the 
 equation : 
 
 2KMn04 -|- 3H2SO4 -j- 5H2O2 = 2MnS04 -f- K2SO4 -f- SH^O -(- 5O2. 
 
 Or the liquid to be examined (rain-water) for hydrogen peroxide is shaken in a stop- 
 pered glass with a 5 per cent, solution of pota.ssium iodide and some starch-paste, 
 allowed to stand several hours, and the iodine which separates is then determined color- 
 irnetrically (Schone). 
 
 Thermo-chemical Deportment. — The great reactivity ef hydrogen per- 
 oxide, its various modes of formation and its transpositions are fully 
 exi)lained by its thermal relations. Compared with water, it, like 
 ozone, is an endothermic compound, i. e., it contains more energy than 
 water : (H2G,0) = — 23 Cal. Therefore, its formation from the latter 
 requires the addition of energy. Its ])roduction by the oxidation of 
 
104 
 
 INORGANIC CHEMISTRY. 
 
 water is excei)lional, and occurs with (lifficulty. It loses energy (heat) 
 and readily changes to more stal)le water. Ij'ke ozone, its oxidations 
 proceed more energetically than those with free oxygen, because 23 
 Cal. more are disengaged. 
 
 The production of hydrogen peroxide by the transposition of barium 
 peroxide and hydrochloric acid proceeds with the liberation of heat: 
 
 BaO^ 4 - 2lICl,Aq = P>aCl2,Aq 4 H2O2 . . . 4-22.0 Cal. 
 
 Hydrogen peroxide is similarly formed from other peroxides, e. g., 
 potassium, calcium, and zinc i)eroxides. The superoxides or dioxides 
 of manganese and lead (Midland 1*1)02) not yield h\drogen i)erox- 
 ide with acids. This is due to the fact that they are differently con- 
 stituted chemically from the other j)eroxides. 
 
 Hydrogen peroxide occurs in slight (juantity in the air and is detected 
 in almost all rain-water and in snow — but not in natural dew and frost. 
 Its quantity varies from 0.5 to i milligram in a liter of rain. Its forma- 
 tion in the air is probably induced by the action of ozone upon ammonia, 
 whereby ammonium nitrite, hydrogen peroxide, and oxygen result (Carius). 
 
 Analysis shows that hydrogen |)eroxide consists of i.oi parts of hydrogen 
 and 16 parts of oxygen; its simj)lest formula would therefore be HO. 
 The difficult volatility of the comj)ound, and the reactions already 
 described, cause us to believe that the molecule of hydrogen ])eroxide 
 is more complicated, and is expressed by H2O2. It is supposed that the 
 peroxide is composed of two groups of OH, called hydroxyl ; these are 
 combined with each other : HO — OH. 
 
 2. SULPHUR. 
 
 Atom : S = 32.06. Molecule : S2 == 64.12 (above 1000° C.). 
 
 Sulphur is distributed throughout nature, both free and in a combined 
 state. In volcanic regions, like Sicily, it occurs free, and there it forms 
 vast deposits, mixed with gy])sum, calcite and marl. Its compounds 
 with the metals are known as blendes or glances. In combination with 
 oxygen and calcium it forms calcium sulphate, the widely distributed gyp- 
 sum. It is also ])resent in many substances of the vegetable and animal 
 kingdoms — e. g., in the albuminoid bodies. It is interesting to note that 
 many bacteriae and algae contain as much as one-fourth of their weight 
 of sul})hur. 
 
 d'o obtain sulphur, the natural product in Sicily is arranged in heaps, 
 covered with earth, and then melted, or it is distilled from earthen 
 retorts. To further ])urify this crude coanmercial product it is redistilled 
 from cast-iron retorts, and when in a molten condition is run into cylin- 
 drical forms — stick sulphur. If the suli)hur vapors are rapidly cooled 
 during distillation (which occurs by conducting them into a stone 
 chaml)er through which cold air circulates), they condense to a fine 
 yellow jiowder, known as jlowers of sulphur (Flores sulphuris). 
 
SULPHUR. 105 
 
 Sulphur may be obtained by heating the well-known pyrites (FeS^) 
 away from air contact. 
 
 Appreciable quantities of sulphur are obtained from the material used in purifying gas — 
 Laming’ s substance (ferric hydrate, lime and sawdust ; which finally contains iron sulphide 
 and sulphur), as well as in the LeBlanc soda process (see Soda). 
 
 Free sulphur exists in several allotropic modifications (see p. 87). 
 
 1. Ordinary octahedral or rhombic sulphur in nature in beauti- 
 
 ful, well-crystallized rhombic octahedra (p. 35). It is pale yellow, hard and 
 very brittle; on rubbing, it becomes negatively electrified. The specific 
 gravity of this variety equals 2.07. It dissolves with difficulty in alcohol 
 and ether; but is more readily soluble in hydrocarbons and ethereal oils. 
 The best solvents are sulphur monochloride (S2CI2) and carbon bisulphide 
 (CS2) ; 100 parts of the latter at 22° dissolve 46 parts of sulphur. By 
 slow evaporation of the solutions sulphur crystallizes in transparent, lus- 
 trous, rhombic octahedra, like those occurring in nature. It fuses 
 at 114.5° C. to a yellow, mobile liquid, which upon further heating be- 
 comes dark and thick, and at 250° is so viscid that it cannot be poured 
 from the vessel containing it. Above 300° it again becomes a thin 
 liquid, boils at 448°, and is converted into an orange-yellow vapor. 
 
 2. The prismatic or 77 ionoclinic sulphur is obtained from the rhombic 
 when the latter is heated to its point of fusion ; on cooling, it generally 
 assumes the monoclinic form (rhombic crystals separate at 90° from 
 sulphur which has been heated beyond the point of fusion). The mono- 
 clinic crystals are best obtained as follows: Fuse sulphur in a clay cruci- 
 ble, allow it to cool slowly until a crust appears on the surface ; break 
 this open near the side and pour out the liquid- portion. The walls of 
 the crucible will be covered with long, somewhat curved, transparent, 
 brownish-yellow needles, or prisms of the monoclinic system. The same 
 are obtained when a solution of sulphur in carbon bisulphide is heated to 
 100°, in a sealed tube, and then gradually allowed to cool; mono- 
 clinic crystals at first separate, and later, at lower temperatures, rhombic 
 octahedra. The monoclinic crystals separated from the solution are 
 almost colorless and perfectly transi)arent. 
 
 Prismatic or octahedral crystals may be obtained from a supersaturated benzene 
 solution of sulphur, by adding small fragments of the corresponding crystals to the 
 solution. 
 
 This form of sulphur has a lower specific gravity (=1.96) and fuses 
 above 1 20°. It is soluble in the same solvents as the rhombic variety. It 
 is very unstable; the transparent prisms and needles become opaque and 
 pale yellow at ordinary temperatures, and specifically heavier, and pass 
 over into an aggregate of rhombic octahedra retaining the external pris- 
 matic form. Stick sulphur deports itself similarly; the freshly moulded 
 sticks are comjiosed of monoclinic prisms, but in time their specific 
 gravity changes and they are converted into the rhombic modification. 
 
 3. So/t, plastic sulphur appesivs to consist of two modifications. It is 
 obtained when sulphur heated above 230° is poured in a thin stream 
 into water ; it then forms a soft, fusible mass, of a yellowish-brown color. 
 
INORGANIC CHEMISTRY. 
 
 lOG 
 
 and its specific gravity ccpials 1.92. In a few days it hardens, and is 
 converted into the rhombic variety. At 95° the conversion is instan- 
 taneous and accoini)anied by the evolution of considerable heat. It is 
 only partly soluble in carbon bisulphide, leaving an amorphous powder 
 undissolved — amorphous insoluble sulphur. As it reverts to the rhombic 
 at 100° molten suljdiur must be quickly chilled in order to obtain much 
 of it. It is also ])roduced when light acts upon dissolved or fused sul- 
 phur, and in the decomposition of the halogen-suljjhur compounds by 
 water. Flowers of sulphur, obtained by cooling sulphur vai)or quickly, 
 are for the most j)art insoluble in carbon bisulphide. 
 
 On adding hydrochloric acid to polysulphide solutions of potassium or 
 calcium, suli)hur sei)arates as a fine, white powder, known as 7 fiilk of sulphur 
 (Lac sulphuris) : 
 
 K2S5 + 2IICI 2KCI -f II^S -f 4S. 
 
 This is amorphous, soluble in carbon bisulphide, and gradually passes 
 into the rhombic form. 
 
 The existence of these various modifications of sulphur, like that of ordi- 
 nary oxygen and ozone, may be attributed to the fact that the sulphur 
 molecules do not, under all circumstances, consist of the same number 
 of atoms. This supposition is confirmed by the deportment of sulphur 
 vapor. The density of the latter at 500° has been found to equal 
 192 (O2 = 32). The vapor density steadily diminishes with increase of 
 temperature and becomes constant at 1000°, and equals 64. Since 
 the atomic weight of sulphur equals 32, it follows that at 1000° its 
 molecules consist of two atoms (S2 = 64 = 32 X 2). At lower temperatures 
 the sulphur vapors appear to contain molecules consisting of more than 
 two atoms; thus at 500°, where the vapor density equals 192, the molecule 
 consists of six atoms. According to this the hexatomic sulphur molecules 
 dissociate, on further heating, and break down into normal diatomic mole- 
 cules ; the dissociation begins at 700° and is complete at 1000° (see p. 93). 
 Since, therefore, the sulphur molecules in vapor form consist of two 
 atoms at very high temperatures and of more atoms at lower, we may 
 assume that the molecules in the liquid and solid condition are more 
 complicated, and that the various allotropic modifications are influenced 
 by the number of atoms contained in the molecules. Other solid metal- 
 loids — e. g., selenium, phosphorus, arsenic, carbon, and silicon — occur in 
 different modifications. As yet we have no means of ascertaining the 
 molecular size of the elements in liquid and solid conditions; there is 
 much, however, favoring the idea that when free they consist of complex 
 atomic groups. Another explanation of allotropy will be given in con- 
 nection with the different varieties of ])hosphorus. 
 
 Chemical J^roperlies . — In its chemical behavior sulphur is very similar 
 to oxygen. It unites directly with most of the elements. It ignites and 
 burns with a pale bluish flame, forming sulphur dioxide (SO.^) when 
 heated to 260° in the air. This union with oxygen occurs gradually 
 
HYDROGEN SULPHIDE. 
 
 107 
 
 even at lower temperatures (about 180°); in the dark it is accompanied 
 by a white phosphorescent flame. Nearly all the metals combine with it 
 to form sulphides. By rubbing mercury, flowers of sulphur and water 
 together, we obtain black mercury sulphide. A moist mixture of iron 
 filings and sulphur glows after a time. Copper and iron burn in sulphur 
 vapor. 
 
 The sulphides are analogous to the oxides, exhibit similar reactions, 
 and in the main possess a similar composition, as may be seen from the 
 following formulas: 
 
 H.^O, Water. 
 
 KOH, Potassium hydrate. 
 BaO, Barium oxide. 
 
 COj, Carbon dioxide. 
 
 K.^COj, Potassium carbonate. 
 
 HgS, Hydrogen sulphide. 
 
 KSH, Potassium sulphydrate. 
 
 BaS, Barium sulphide. 
 
 CS.^, Carbon bisulphide. 
 
 KjCSg, Potassium sulphocarbonate. 
 
 COMPOUNDS OF SULPHUR WITH HYDROGEN. 
 
 1. HYDROGEN SULPHIDE. 
 
 H2S = 34.08. 
 
 In nature hydrogen sulphide occurs principally in volcanic gases and 
 in the so-called sulphur waters (p. 91). It is always produced in the 
 decomposition of organic substances containing sulphur, particularly the 
 albuminoids by which the sulphates, e. g., gypsum, are reduced with the 
 formation of hydrogen sulphide. Hence the occurrence of the gas in 
 eggs, sewers, etc. It may be formed directly from its constituents, 
 although in small quantity, if hydrogen gas be conducted through boil- 
 ing sulphur, or if sulphur vapors, together with hydrogen, be conducted 
 over porous substances (pumice-stone, bricks) heated to 500°. Many 
 sulphides are reduced upon ignition in a stream of hydrogen, with sepa- 
 ration of hydrogen sulphide : 
 
 Ag^S + H, = 2Ag + H,S. 
 
 For its production acids are allowed to act upon sulphides. Ordi- 
 narily iron sulphide and diluted sulphuric acid are employed ; the action 
 occurs at ordinary temperatures : 
 
 FeS + H.SO^ = FeSO, -f H^S. 
 
 Iron Sulphuric Ferrous 
 sulphide. acid. sulphate. 
 
 The operation is performed either in a Kipp apparatus (p. 42) or in the 
 one pictured on p. 43. Hydrogen sulphide thus obtained contains ad- 
 mixed hydrogen, in consequence of metallic iron existing in the sulphide. 
 The pure gas is obtained by heating antimony sulphide with concentrated 
 hydrochloric acid : 
 
 Sb.Sj -f 6IIC1 = 2SbCl3 -P 3lI.,S. 
 
io8 
 
 INORGANIC CHEMISTRY. 
 
 Properties . — Hydrogen sulphide is a colorless gas, having an odor sim- 
 ilar to that of rotten eggs; inhaled in large (juantities it has a stupefying 
 effect, and is very ])(hsonoiis. At medium temperatures it eondenses 
 under a i)ressure of ry atmospheres (under ordinary pressure at — 74°) to 
 a colorless licpiid of specific gravity 0.9, which boils at — 63.5° under 760 
 mm. pressure, and at — 91° solidifies to a white crystalline mass. It is 
 1,18 times heavier than air. One volume of water dissolves 3 or 4 times 
 its volume of the gas ; the solution jiossesses all the jiroperties of gaseous 
 hydrogen suli)hide and is therefore called hydrogen sulphide water. 
 
 Ignited in the air the gas burns with a blue flame, water and suliihur 
 dioxide resulting : 
 
 II,S -f 3O = 11,0 + SO,. 
 
 With insufficient air access, or when the flame is cooled by the intro- 
 duction of a cold body, only hydrogen burns and sulphur sejiarates out in 
 a free condition. This behavior is utilized in the technical preparation 
 of sulphur from soda residues. In acpieous solution hydrogen sulphide is 
 similarly decomposed by the oxygen of the air at ordinary temperatures, 
 sulphur sejiarating as a fine powder : 
 
 II,S + O =r H,0 + S. 
 
 The halogens behave like oxygen ; the hydrides of the halogens are 
 formed with separation of sulphur : 
 
 H,S + I2 = 2HI 4- S. 
 
 This reaction, which occurs only in the presence of water, serves for the 
 production of hydrogen iodide (p. 63). 
 
 As hydrogen sulphide has a great affinity for oxygen, it withdraws the 
 latter from many of its compounds, and it therefore acts as a reducing 
 agent. Thus chromic, manganic, and nitric acids are reduced to lower 
 stages of oxidation. On pouring fuming nitric acid into a vessel filled 
 with the dry gas, the mixture will unite with a slight explosion. 
 
 Hydrogen sulphide possesses weak acid properties, reddens blue litmus- 
 ])aper, forms salt-like compounds with bases, and is, therefore, termed 
 hydrosulphuric acid. Nearly all the metals liberate hydrogen from it, 
 yielding metallic sulphides: 
 
 Pb -f 11,5 = PbS + II,. 
 
 Wi h the oxides and hydroxides of the metals hydrogen sulphide yields 
 sulphides and suli)hydrates : 
 
 KOI I + II,S = KSH -p 11,0. 
 
 Potassium 
 
 hydrosulphide. 
 
 PbO + II,S = PbS -f 11,0. 
 
 Lead 
 
 sulphide. 
 
 Sulphides, therefore, like the compounds of the halogens with the 
 metals, may be viewed as the salts of liydrosuliihuric acid. The sulphides 
 of almost all the heavy metals are insoluble in water and dilute acids; 
 
MOLECULAR FORMULA OF HYDROGEN SULPHIDE. I09 
 
 therefore, they are precipitated by hydrogen sulphide from solutions of 
 metallic salts: 
 
 CuSO, + H^S = CuS 4- H^SO,. 
 
 Copper Copper Sulphuric 
 
 sulphate. sulphide. acid. 
 
 The precipitates thus obtained are variously colored (copper sulphide, 
 black; cadmium sulphide, yellow; antimony sulphide, orange), and 
 answer for the characterization and recognition of the corresponding 
 metals. Pai)er saturated with a lead salt solution is at once blackened by 
 hydrogen sulphide, lead sulphide being formed— a delicate test for the 
 gas. 
 
 Thermo-chemical Deportment . — Hydrogen sulphide is a feebly exother- 
 mic compound. When hydrogen gas unites with amorphous sulphur to 
 form hydrogen sulphide 4.7 Cal. are developed. When the gas dissolves 
 in much water its heat of solution equals -{-4.6 Cal., so that the total 
 heat of formation of hydrogen sulphide in dilute aqueous solution is 
 9.3 Cal. : 
 
 (H2,S — gas) = 4.7 ; (H,S,Aq) == 4.6 ; (H,,S,Aq) = 9.3. 
 
 It is because of this low heat of formation that the gas is produced with 
 such difficulty from its elements, and it is for this reason that it is so readily 
 dissociated by heat into its elements. Its entire chemical deportment is 
 also accounted for by its heat of formation (p. 114). 
 
 MOLECULAR FORMULA OF HYDROGEN SULPHIDE. ATOMIC WEIGHT OF 
 
 SULPHUR. 
 
 Considerations similar to those which led us to adopt the molecular formula H2O for 
 water impel us to accept the formula H.^S for hydrogen sulphide. Its gas density, its 
 analysis, and the fact that in an equal volume of any sulphur-containing gaseous com- 
 pound there has never been observed less sulphur than is contained in hydrogen sulphide 
 speak in favor of the assumption. The atomic value of sulphur is therefore ^2.06 (com- 
 pare p. 97). 
 
 From the molecular formula H^S we further conclude that the hydrogen contained in 
 one volume of hydrogen sulphide would occupy in a free condition the same volume as 
 the latter : 
 
 nll.^S contains nH2. 
 
 I vol. I vol. 
 
 This conclusion is verified experimentally as follows : In a bent glass tube filled with mer- 
 cury (p. 78, Fig. 41), introduce dry hydrogen sulphide gas ; then in the bent portion place 
 a piece of tin, which is heated by a lamp. The sulphur of the hydrogen sulphide combines 
 with the metal to form solid tin sulphide, while hydrogen is set free ; its volume is 
 exactly equal to the volume of the employed hydrogen sulphide. The quantity of sul- 
 phur, 32 parts, in vapor form, at 1000°, when the density is 64 (p. 106) combined 
 with hydrogen (2 parts) would equal exactly half the volume of the hydrogen ; at 500°, 
 however, when the vapor density is three times as great, it will equal one-sixth of the 
 volume of the hydrogen. Written molecularly, we have : 
 
 At 500° : Sg 
 
 + 
 
 6112 = 
 
 6II2S. 
 
 I vol. 
 
 
 6 vols. 
 
 6 vols. 
 
 At 1000°, however : S2 
 
 
 2II2 = 
 
 2II2S. 
 
 I vol. 
 
 
 2 vols. 
 
 2 vols. 
 
I TO 
 
 INORGANIC CHEMISTRY. 
 
 HYDROGEN PERSULPHIDE. 
 
 Just as hydrogen ijeroxide, is formed by the action of acids 
 
 upon some peroxide, so may hydrogen persul])hide lie obtained from 
 metallic persulphides or polysulphides. Sodium and potassium each 
 form four iiolysulphides which are available for this jiurpose. With 
 sodium for exam])le there are, in addition to the ordinary suljjhide Na^S, 
 also the compounds Na,^S^, Na^S.,, Na,^S^, Na^S^. J'\;r (piite a while it was 
 thought that these polysulphides were decomposed by dilute acids simi- 
 larly to the decomposition of barium jieroxide by dilute sulphuric acid : 
 
 V,vi(\ + 1 1 , SO, = P,aSO, -f 
 
 that therefore each polysuljihide had a corresponding hydrogen polysul- 
 phide. This idea was enforced by the fact that the alkaloid strychnine 
 formed a crystalline product with H^S.,, and brucine one with 
 Rebs, however, demonstrated that when the different polysulphides are 
 decomposed it is probably always hydrogen pentasulphide, H^S^, which is 
 produced. It is a light-yellow colored, transjiarent mobile oil with a 
 peculiar odor. Its sjiecinc gravity equals 1.7 1 at 15°. When dry and 
 away from air contact it decomposes slowly. Water resolves it quickly 
 into hydrogen sulphide and sulphur, which separates. Heat induces the 
 same decomposition : 
 
 H2S5 = S, + H,S. 
 
 Rebs (see Ann. Chem. (1888) 246, 354) explains the formation of hydrogen penta- 
 sulphide from the alkaline bi-, tri-, and tetrasulphides in the following equations: 
 
 4Na2S, + 8 HC 1 8NaCl + 4H2S, ; 4H2S, = + H2S. 
 
 4Na2S3 + 8 HC 1 = 8NaCl + 4H2S3 ; 4H2S3 = 2II2S5 -f 2H2S. 
 
 4Na2S2 + 8 HC 1 = 8NaCl + 4H2S2 ; 4H3S2 = H2S5 + 3H2S. 
 
 COMPOUNDS OF SULPHUR WITH THE HALOGENS. 
 
 Sulphur and chlorine unite to form three compounds: SCI2, SC 1 „ and 
 S2CI2. It is only the last which meets with any practical application. 
 
 Sulphur Dichloride — SC^ — is produced when sulphur monochloride, S2CI2, is satu- 
 rated with chlorine at 6° to 10° : 
 
 S2CI2 + CI2 = 2SCI2. 
 
 The excess of chlorine is removed by conducting a stream of dry carbon dioxide 
 through it. 
 
 It is a dark-red colored lifpiid, with a specific gravity of 1.62; boils at 64°, with 
 j)artial decomposition into S2CI2 and ; the dissociation commences at ordinary tem- 
 peratures. 
 
 Sulphur Tetrachloride — SCI, — only exists at temperatures below 0° C. It is 
 formed by saturating S2CI2 with Cl at — 20° to — 22°. Tlie dissociation commences 
 at — 20°, and is comjdete at -[ 6°. It yields crystalline compounds with some chlo- 
 ridc.s — e, g.^ SnCl,, AsCIj, .SbCb. 
 
SULPHUR MONOCHLORIDE. 
 
 Ill 
 
 The most stable of the sulphur chlorides is 
 
 Sulphur Monochloride — S2CI2 — which is formed when chlorine is 
 conducted over molten sulphur contained in the flask C (Fig. 49). It 
 distils over and condenses in the receiver D ; the product is redistilled, 
 to obtain it pure. [Fig. 49, A: chlorine generator; .• wash-bottle ; 
 E : entrance for water intended to chill the vapors.] 
 
 Sulphur monochloride is a reddish-yellow liquid with a sharp odor, 
 provoking tears, having a specific gravity of 1.68, and boiling at 138°. 
 Its vapor density equals 135 (03 = 32), corresponding to the molecular 
 
 formula S2CI2. It fumes strongly in the air, and is decomposed by water 
 into sulphur dioxide, sulphur and hydrochloric acid : 
 
 2S2CI2 -f 2H2O == SO2 -f 4HCI + 3S. 
 
 Sulphur monochloride dissolves sulphur readily and serves in the vulcan- 
 ization of caoutchouc. 
 
 Bromine forms analogous compounds with sulphur. S.^Brj is a red liquid, boiling at 
 190-200°. Compounds of iodine and sulphur are not definitely known. 
 
II2 
 
 INORGANIC CHEMISTRY. 
 
 1 
 
 3. SELENIUM. 
 
 Atom : Sc = 79.1. Molecule: Scj 158.2 (at 1400°), 
 
 Tliis element is not very abundant in nalure, and is only found in 
 small (luantities, jirincijially in certain iron i>yritcs (in Sweden and 
 Jjohemia). Upon roasting this ore of iron, for the prejiaration of sul- 
 l)huric acid, selenium settles out in the chimney dust or in the deposit of 
 the lead chambers (comiiare Sulphuric Acid), and was found there by 
 Berzelius in the year 1817. It was called selenium (frskrjurj, moon) to 
 indicate its relation to the already known tellurium {Jcllus, earth). 
 
 Like sul{)hur, selenium forms different allotropic modifications. Avior- 
 phoiis selenium, obtained by the reduction of selenium dioxide (SeOJ by 
 means of suli)hur dioxide (SOp, is a reddish-brown powder, soluble in 
 carl)on bisulphide, and has a specific gravity of 4. 26. Selenium crystallizes 
 from carbon bisiiljihide in brownish-red crystals of specific gravity 4.5. 
 The solution of potassium selenide is brown-red, and when it is ex])osed to 
 the air, black leaf-like crystals of selenium (with si)ecific gravity 4.8) sepa- 
 rate. These are isomorphous with suljdmr and insoluble in carbon bisul- 
 ])hide. Upon suddenly cooling fused selenium it solidifies to an amorphous, 
 glassy, black mass, which is soluble in carbon bisulphide and hasa sjiecific 
 gravity of 4.28. When selenium (amorphous) is heated to 90-100°, its 
 temperature suddenly rises above 200° ; it is converted into a crystal- 
 line, dark-gray mass with a specific gravity of 4.8. It possesses metallic 
 luster, conducts electricity, and is insoluble in carbon bisulphide. The 
 crystalline, insoluble modification is obtained by slowly cooling the molten 
 selenium. 
 
 Selenium melts at 217°, and boils at about 660°, passing into a dark- 
 yellow vapor. The vapor density diminishes regularly with increasing 
 temperature (similar to sulphur), and becomes constant at 1400°. It 
 then equals 158; the molecule of selenium at 1400° consists of two 
 atoms (2 X 79- 1 = 158-2). 
 
 Selenium resembles sulphur very closely in its chemical behavior. It 
 burns in the air with a reddish-blue flame, forming selenium dioxide, and 
 emits a peculiar odor resembling rotten horse-radish. It dissolves with a 
 green color in concentrated sulphuric acid, and forms selenious and sul- 
 phurous acids : 
 
 Se -f- 2 H 2 S 0 ^ = SeOj T 2SO2 + 2H2O. 
 
 Hydrogen Selenide — H2Se — produced, like hydrogen sulphide, 
 from iron selenide and hydrochloric acid — is a colorless, disagreeably 
 smelling gas with poisonous action. In the air the aqueous solution 
 I ecomes turbid and free selenium separates. 
 
 Willi chlorine selenium forms SeCl, and Se2Cl2, perfectly analogous in general to the 
 sulphur cornjiounds hut differing fnim them in that selenium tetrachloride is a solid which 
 sublimes and does not begin to decompose until at about 200°. 
 
TETXURIUM. 
 
 II3 
 
 4. TELLURIUM. 
 
 Atom: Te = i27. Molecule: Te2 = 254 (at 1700°).* 
 
 Tellurium is of rare occurrence, either native or in combination with 
 metals. It is associated with gold and silver in sylvanite, with silver and 
 lead in altaite, and with bismuth in tetradymite. It is found principally 
 in Transylvania, Hungary, California, Virginia, Bolivia, Brazil, and in 
 the volcanic formations of the Liparian Islands. 
 
 The tellurium precipitated by sulphurous acid from a solution of tel- 
 lurous acid is a black ])ovvder of specific gravity 5.9. It is silver-white 
 when fused, of a perfect metallic luster, and conducts electricity and 
 heat. It crystallizes in rhombohedra, having a specific gravity 6.4. It 
 fuses at 452° and boils at 1390°. When heated in the air it burns, with 
 a bluish-gray flame, to tellurium dioxide (Te02). 
 
 The vapor density of tellurium at 1400-1700° has been discovered to 
 be about 254, corresponding to the molecular formula Te2. 
 
 Hydrogen Telluride — H2Te — formed by the action of hydrochloric 
 acid upon zinc telluride, is a colorless, very poisonous gas, with disagree- 
 able odor. 
 
 Two chlorides — TeC]2 and TeCfi, and two bromides, TeBr2 and TeBr^ — have been 
 formed. The tetrachloride boils at 380°. As its vapor density corresponds to the formula 
 TeCb the tellurium is quadrivalent. 
 
 SUMMARY OF THE ELEMENTS OF THE OXYGEN GROUP. 
 
 The elements oxygen, sulphur, selenium and tellurium form a natural 
 grouj) of chemically similar bodies. The similarity of the last three is 
 especially marked, while oxygen, ])ossessing the lowest atomic weight, 
 stands somewhat apart. Among the halogens, fluorine exhibits a similar 
 deportment; it departs somewhat from its analogues, chlorine, bromine 
 and iodine. Like the halogens the elements of the oxygen group present 
 a gradation in their properties corresponding to their atomic weights : 
 
 os Se Te 
 Atomic weights, 16 32.06 79.1 127. 
 
 With the increase in the atomic weight there occurs a simultaneous 
 condensation of substance, the volatility diminishes, while the specific 
 gravity and the points of fusion and boiling increase, as may be seen in 
 the following table. 
 
 *The atomic weight of tellurium was first made 128 and subsequently 125. But the 
 later development of the periodic system made it more probable that it was even lower 
 than the atomic weight of iodine (126.8), which view has since been confirmed exjieri- 
 mentally by Brauner. More recently results have been obtained which argue for the 
 higher atomic weight, d'e -- 127. (See Staudenmaier, Zeit. f. anorg. Chem, 10 (1895), 
 189.) 
 
 10 
 
in()R(;ani(: cukmistry. 
 
 114 
 
 
 Oxy(;kn. 
 
 Sui.i’inm. 
 
 SlU.KNIUM. 
 
 IT'.I.I.UKIUM, 
 
 Specific gravity, . . 
 
 1. 124 (at — 181°) 
 
 1.95-2.07 
 
 4. 2-4. 8 
 
 r,.4 
 
 Melting point, . . . 
 
 
 114-5° 
 
 217° 
 
 452° 
 
 Boiling j)oint, . . 
 
 — 181° 
 
 440° 
 
 600° 
 
 1 390° 
 
 Gas density, .... 
 
 32.00 
 
 64. 12 
 
 158.2 
 
 254 
 
 Oxygen is a difYicultly c()ercil)le gas, while the others are solids at 
 ordinary tem})eratures. We must, however, bear in mind that sulphur, 
 selenium and tellurium in a free state are probably composed of larger 
 complex atomic groiii)S (see j). 106). 
 
 Further, with rising atomic weight the metalloidal i)asses into a more 
 metallic character. Tellurium exhibits the physical properties of a 
 metal; even selenium possesses metallic properties in its crystalline 
 modification. In chemical deportment, however, the metalloidal char- 
 acter shows scarcely any alteration. 
 
 All four elements unite directly, at elevated temperatures, with hydrogen, to com- 
 pounds the composition of which is expressed by the general formula MH.^. At still 
 higher temperatures these derivatives are again resolved into their elements. The simi- 
 larity of water and the hydrides of sulphur, selenium and tellurium is restricted to their 
 formulas. They are entirely different in their chemical nature. Water is absolutely 
 necessary to life, while hydrogen sulphide, selenide and telluride are dangerous poisons. 
 It must also occasion surprise that oxygen, the only gaseous member of the group, forms 
 a derivative with hydrogen which is liquid at the ordinary temperature and a solid at 0°, 
 whereas the hydrides of the elements sulphur, selenium and tellurium, volatile with diffi- 
 culty, are also gaseous and are, comparatively speaking, condensed with more difficulty. 
 It was indeed observed with the halogens that the hydride of fluorine, the most difficult 
 to condense, possessed the highest boiling point, and we learned that the explanation for 
 this was that, in accordance with its vapor density, hydrogen fluoride did not have the 
 formula HP", but H^F2. Water also differs very much the7-mo-che77ncally from the hydrides 
 of the other elements of the group. It is a strongly exothermic body, while hydrogen 
 sulphide, on the other hand, formed from hydrogen and solid sulphur with the evolu- 
 tion of very little heat (H2,S = about 4 Cal.), and selenium and tellurium hydrides are 
 indeed endothermic bodies. 
 
 This is a gradation similar to that observed in the halogen hydrides (p. 65). 
 In accord with this we find that oxygen will displace these elements from hydrogen 
 sulphide, hydrogen selenide, and hydrogen telluride when in aqueous solution with the 
 formation of water. At higher temperatures and with an excess of oxygen the dioxides 
 (SO2, Se02) result. The halogens decompose them more readily than oxygen. 
 
 NITROGEN GROUP. 
 
 This group consists of nitrogen, ])hosphorus, arsenic, antimony, and 
 bismuth, d'lie last ])ossesses a decidedly metallic character. It does 
 not, like the other four elements, form gaseous derivatives with hydrogen, 
 d'he gases argon, helium, metargon, kry])ton, xenon, and neon, occurring 
 in the air, will l)e described in connection with nitrogen. 
 
NITROGEN. 
 
 
 1 . NITROGEN. 
 
 Atom : N = 14.04, Molecule : N2 = 28.08. 
 
 Nitrogen exists free in the air, four-fifths by volume are nitrogen and 
 one-fifth oxygen. In combination, it is chiefly found in the ammonium 
 and nitric acid compounds, as well as in many organic substances of the 
 animal and vegetable kingdoms. 
 
 Until 1894 it was thought possible to isolate nitrogen from the air 
 by depriving the latter of its second constituent — oxygen. This is 
 
 effected by bodies capable of ab- 
 sorbing oxygen without acting 
 upon the nitrogen, e. g , phos- 
 phorus, hepar, alkaline solutions 
 of pyrogallol and heated copper. 
 
 The experiment can be most 
 easily and simply performed in 
 the following manner: Several 
 pieces of phosphorus are placed 
 in a dish swimming on water, 
 ignited, and a glass bell-jar 
 placed over them (Fig. 50). In 
 a short time, when all the oxy- 
 gen is absorbed from the air, the 
 phosphorus will cease burning ; 
 the phosphorus pentoxide pro- 
 duced dissolves in the water, and the residual gas consists of almost pure 
 nitrogen : its volume will equal four-fifths of the air taken. Another 
 procedure consists in conducting air through a red-hot tube filled with 
 copper turnings; the copper unites with the oxygen and pure nitrogen 
 escapes. The portion of air remaining after these experiments was con- 
 sidered to be nitrogen until 1894 when Lord Rayleigh and W. Ramsay 
 proved to the universal surprise that so-called “ atmospheric nitrogen ” 
 was a mixture of nitrogen and argon. Lately, other very probably 
 elementary gases have also been discovered in it. 
 
 Pure nitrogefi can only be obtained from nitrogenous chemical com- 
 pounds. The following is a simple method to this end: Heat ammonium 
 nitrite in a small glass retort; this decomposes the salt directly into 
 water and nitrogen : 
 
 NH^N 02 = N2 T 2H2O. 
 
 In place of ammonium nitrite a mixture of potassium nitrite fKNO.^) and ammonium 
 chloride (NIbCl) may be u.sed ; upon warming, the.se salts yield, by double decomposi- 
 tion, potassium chloride and ammonium nitrite (KNO.^ + NH^Cl = NH^NOg + KCl), 
 which latter decomposes further. As potassium nitrite usually contains free alkali, some 
 potassium bichromate is added to neutralize the same. Practically, the .solution consists 
 of I part of potassium nitrite, i part of ammonium chloride, and i part of potassium 
 bichromate in 5 parts of water, and is then boiled ; to free the liberated nitrogen from 
 every trace of oxygen the gas is conducted over ignited copper. 
 
 The action of chlorine upon aqueous ammonia produces pure nitrogen. 
 Ammonia is a compound of nitrogen and hydrogen. While the chlorine 
 
 Fig, 50. 
 
INORGANIC CHFMISTRY. 
 
 I l6 
 
 conihinc's vvitli the hydrogen of tlie ammonia to liydroclilorie acid, nitro- 
 gen is lilierated. d'he hydrochloric acid unites witli the excess ol 
 ammonia to ammonium chloride (NH^Clj. 'I'he following e<juations 
 express the reactions : 
 
 and 
 
 2NII3 4 - 3CI2 = Nj -f GIICl 
 
 GHCl 4 6NII3 = 6NH,C1. 
 
 Animoiiiuin 
 
 cliloridc. 
 
 d'he apiiaratus jiictured on ]>. 50 will serve to carryout the experiment. 
 'The disengaged chlorine is conducted through a Woulff wash-bottle c(m- 
 taining ammonia water, the free nitrogen being collected over water. 
 
 In thi.s expernnent the greate.st care should he exerci.se(l tluit an excess of chlorinw is 
 not conducted into the solution, because its action ui)on the aminoniuin chl(;ride will 
 cause the formation of an exceedingly explosive body, nitrogen chloride, NCI3 (p. 133;, 
 which separates in oily drops. 
 
 Pt'operiies - — Nitrogen is a colorless, odorless, tasteless gas. One liter 
 of it weighs I *2507 grams at 0° and 760 mm. pressure. It is therefore 0.97 
 times as heavy as air. Its critical temjierature lies near — 146^^, and its 
 critical pressure equals 35 almosjdieres (p. 487. Liquid nitrogen is color- 
 less, boils under a jtressure of one atmosphere at — 194°, at — 225° under 
 a pressure of 4 mm., and has a specific gravity of 0.885 — ^94°)- 
 
 solidifies at — 214°. 
 
 In its chemical de])ortment it is extremely inert, combining directly 
 with but few elements, e. g., with oxygen and hydrogen under the in- 
 fluence of the electric S})ark ; with magnesium and other metals at more 
 elevated temperatures, and with lithium at the ordinary temperature. It 
 does not support combustion or respiration ; a burning candle is extin- 
 guished and animals are suffocated by it. This is not due to the activity 
 of the nitrogen, but to the absence of oxygen — a substance which cannot 
 be dispensed with in combustion and respiration. The presence of nitro- 
 gen in the air moderates the strong oxidizing property of the pure oxygen. 
 
 THE ATMOSPHERE. 
 
 The air, or the envelo])e encircling the earth, the atmosphere (ar/ioc — 
 vapor; frepaTpa — ball, sphere), consists principally of a mixture of nitrogen 
 and oxygen; it always contains, in addition, slight and variable quanti- 
 ties of aqueous vapor, carbon dioxide, and traces of other substances, as 
 accidental constituents (p. 123). Recently the gases argon, helium, met- 
 argo 7 i, neon, krypton, and xenon have been discovered in it. They occur 
 in small amounts, but are constant constituents. The pressure exerted 
 by the air is measured by a column of mercury which holds it in a 
 state of ecpiilibrium ; the height of the barometric column at the sea- 
 level and 0° C. e(]uals, upon an average, 760 millimeters. As i c.c. of 
 mercury weighs 13.6 grams, 76 c.c. will equal 1033.6 grams, and the last 
 
THE ATMOSPHERE. 
 
 IT7 
 
 number would indicate the pressure which the column of air exerts upon 
 one square centimeter of the earth’s surface. 
 
 One cubic centimeter of air weighs, according to recent experiments, 
 under normal conditions, 0.00129276 gram ; 1000 c.c., therefore, or one 
 liter, would weigh 1.29276 grams. As one liter of water weighs 1000 grams, 
 air is consequently 773 times lighter than it. Airis 14.4 times heavier than 
 hydrogen. Its density is 28.95 referred to O2 — 32. Its specific volume 
 is 773.4, i. e., I gram of air at 0° and 760 mm. pressure occupies 773.4 
 c.c. (pp. 45, 79). 
 
 The liquefactio 7 i of air — or at least a portion of it — can be accomplished 
 by the methods described on p. 48. Dewar and also Olszewsky chilled 
 air, by means of boiling ethylene, almost to its critical temperature 
 ( — 140°), and then liquefied it by a simultaneous, corresponding increase 
 of pressure (75 atmospheres). Another less expensive method, hence well 
 adapted for technical purposes, utilizes the great reduction in temperature 
 sustained by strongly compressed gases when suddenly released Dom 
 pressure. In the apparatus constructed almost simultaneously by Linde 
 (Munich), Tripler (New York) and Hampson (London) the gas under 
 slight pressure and cooled by sudden expansion is directed around a cur- 
 rent of gas passing in an opposite direction under high pressure, whereby 
 the temperature is eventually lowered below the critical temperature. 
 
 Liquid air is colorless. It is turbid, owing to the solid carbon dioxide 
 from which it can be freed by filtration. Its boiling point (according to 
 Dewar) is — 190°, at which temperature almost pure nitrogen is evolved. 
 The specific gravity of freshly prepared liquid air is 0.9951 referred to 
 water at 4°. It contains much more oxygen than the gas. Recently 
 prepared it holds as much as 54 percent, of oxygen by weight, while the 
 gas form has only 23 per cent, of oxygen by weight. Hence the name liquid 
 air is not an entirely correct designation. The oxygen-content of liquid 
 air increases by preservation. This may be accomplished by means of 
 Dewar bulbs or double-walled tin or wooden boxes the air space of which 
 is filled out with silk, etc. It will finally contain as much as 94 per cent, 
 by weight of oxygen. This gradual accumulation of the gas in the liquid 
 air is attributable to the fact that under the ordinary pressure nitrogen 
 boils at — 194°; oxygen, however, at — 184°. There is here the greatest 
 technical possibility, that by boiling out nitrogen from liquid air, almost 
 pure oxygen can be prepared on a large scale. 
 
 The effect of low temperatures upon the physical properties of bodies 
 and upon the course of chemical reactions can be well shown by means 
 of liquid air. Carbon dioxide and acetylene solidify in it. The solid 
 acetylene may be ignited; it then burns away like paraffin. Liquid air 
 immediately solidifies mercury and renders it malleable. Alcohol at 
 once forms drops in liquid air and soon becomes hard and crystalline. 
 The hand may be plunged for a short period into the liquid of — 190° 
 without experiencing any ill effects, because it is at once surrounded by 
 a protecting film. Dewar claims that the color of many bodies is changed 
 if they are immersed in liquid air. Red mercuric oxide, iodide and 
 sulphide (cinnabar) appear yellow in color, while the yellow-green nitrate 
 of uranium appears white. Mention has already been made that chemical 
 
ii8 
 
 INORGANIC CHEMISTRY. 
 
 transpositions do not occur, or at least very slowly, at the temperature of 
 lifpiid air (p. 28). 
 
 History. — It is well known that the air was formerly regarded as an element — a 
 simple substance. However, observations were made very early which argue for the 
 very opposite. As early as the ninth century certain chemists knew that the metals, when 
 heated in air — i.e., by calcination — increased in weight, and some of them correctly 
 attributed this to the absor[)tion of certain air particles l)y the hot metal. Because niter, 
 like air, accelerates combustion, 1 looke (seventeenth century) suspected that both contained 
 an ingredient .serving for combustion. It was especially the English chemists of the .seven- 
 teenth century who busied themselves with researches in this direction. Mayow in particular 
 deserves mention. lie contended that the glowing metal united with the “ spiritus nitro- 
 aerus” of the air. He demonstrated that by respiration as well as by combustion the 
 volume of air standing over water was diminished, and that the residuum — the 
 destroyed air — was no longer available for respiration or combustion. However, the.se 
 germs of the correet idea of combustion phenomena could not develop and grow under 
 the predominance of the peculiar views then extant. Becher, a German chemist, began 
 about 1700 to teach that the phenomenon of combustion originated in a peculiar volatile 
 and escaping earth or kind of sulphur. His pupil Stahl, about 1720, developed this 
 thought into a theory or doctrine, which held almost exelusive sway until near the closing 
 third of the century. Stahl called the substance, which escaped in combustion, 
 (combustible). He failed to note that which became the main subject of investigation 
 with the English chemists — that the burnt metal had increased in weight, and he 
 explained that all combustible bodies consisted of phlogiston and a non-combu.stible sub- 
 stance. Thus, sulphur consisted of phlogi.ston and sulphuric acid, iron of ferric oxide and 
 phlogiston, etc., etc. In the process of combustion the phlogiston escapes as flame and 
 the non-combustible portion remains. The reduction of a metallic oxide by another 
 metal or by a combustible body {e.g., carbon) depends, therefore, upon the passage of 
 phlogiston from l^ie reducing to the reduced body ; hence the terms — phlogisticated and 
 dephlogisticated. The first corresponds to reduction and the second to oxidation. In this 
 manner, because of the absence of more accurate experiments, the phenomena of combustion 
 could be explained with some probability and apparent certainty. And Stahl became the 
 founder of a chemical theory which lasted for more than a half century, and in its 
 decadence found its most ardent advocates among the best and most celebrated English 
 chemists of that period. Bayen (1774) showed that mercuric oxide was reduced without 
 the addition of phlogiston (i. e.^ by heating without the addition of carbon), and this led 
 Lavoisier to his experiments upon the absorption of air in the calcination of the metals. He 
 melted tin in a large, closed, air-tight flask. He had previously determined its weight 
 as well as that of the tin. When the latter had become coated with a thick layer of oxide 
 the apparatus was allowed to cool and was again weighed. Its weight was the same as 
 at first, but when the flask was opened, air rushed in and the weight increased. Hence 
 the tin, by calcination, had abstracted something from the air and had not, as required 
 by the phlogiston theory, given anything to it. About this same period Scheele and 
 Priestley discovered oxygen (p. 80), Rutherford (1742) had again described destroyed 
 air, i. e., nitrogen, and Scheele, by a series of excellent experiments, demonstrated the 
 composition of air and the difference between nitrogen (aer mephiticus, aer vitalis) and 
 carbon dioxide (aer aereum). But he believed in the phlogiston theory. It was 
 Lavoisier, during the .same period, who clearly indicated and showed experimentally the 
 role of oxygen in combustions and oxidations, and, shortly after, proved the elementary 
 nature of nitrogen. Lavoisier named nitrogen azote (from life, and a, privative), 
 from which we get the .symbol Az u.sed in France for nitrogen. Chaptal was the first to 
 give nitrogen the name nitrogenium (whence the symbol N), because it is a constituent 
 of niter (nitrum). How Engli.sh chemists have again extended our knowledge of the 
 constituents of the air to a most unexpected degree may be gathered from pp. I15-123. 
 Coirn)are Berzelius’ Lehrbuch der Chemie, 5 Auil. (1843), I, 140. 
 
 Quantitative Cottiposition of Air . — Its composition is expressed by the 
 quantity of oxygen, argon and nitrogen contained in it, as its remaining 
 admixtures are more or less accidental and variable. 
 
THE ATMOSPHERE. 
 
 II9 
 
 Boussingault and Dumas (181 1) determined the accurate weight com- 
 position of the air (nitrogen and argon were of course calculated as ‘‘at- 
 mospheric nitrogen ”) by the following experiment : A large balloon, V, 
 with a capacity of about 20 liters (Fig. 51), is connected with a porcelain 
 tube, a, b, filled with metallic copper. Balloon and tubes, closed by 
 stop-cocks, are previously emjitied and weighed apart. The bent tubes, 
 A, B, and C, contain caustic potash and sulphuric acid, and serve to free 
 the air undergoing analysis from aqueous vapor, carbon dioxide, and other 
 impurities. The porcelain tube, filled with copper, is heated to a red 
 heat, and by carefully opening the stop-cocks ?/, r, and r a slow current 
 of air is allowed to enter the empty balloon V. The impurities are given 
 up in the bent tubes, and all the oxygen absorbed by the ignited copper, 
 forming cupric oxide, so that only pure nitrogen enters V. Now close the 
 
 cocks and weigh the balloon and porcelain tube, a, b. The increase in 
 weight of the latter represents the quantity of oxygen in the air; the 
 increase in V the quantity of nitrogen. In this manner Dumas and 
 Boussingault found that in 100 parts by weight of air there are con- 
 tained : 
 
 Nitrogen, 7^-995 parts by weight. 
 
 Oxygen, 23.CX35 “ “ “ 
 
 Air, 100,000 “ “ “ 
 
 Asa liter of oxygen weighs 1.4291 grams and a liter of atmospheric 
 nitrogen 1.25 71 grams the volume composition of air would be: 
 
 Oxygen, 20.8 parts by volume. 
 
 Nitrogen, 79.2 “ “ “ 
 
 Air, loo.o “ “ “ 
 
 The density of air is 28.95 (P- 45 ) referred to oxygen etpial to 32. 
 
I 20 
 
 INORGANIC CHEMISTRY. 
 
 olf^ 
 
 Fig. 52. 
 
 the original 
 
 'Fhc voliiinc coiii])osition of air may 1)C directly fcjiind l)y means of the 
 al)sori)tiometer. 'I'he latter is a tube carefully graduated, and sealed at 
 one end. This is filled with mercury, and air allowed to enter; the 
 volume of the latter is determined by reading off 
 ^ the divisions on the tube. Now introduce into 
 
 jijl the tube, through the mercury, a i)latinum wire 
 
 i | having a ball of i)hos])horus attached to the end 
 
 (Fig. 52), or a ball of coke saturated with an 
 alkaline solution of jiyrogallic acid. The phos- 
 phorus absorbs the oxygen of the air, and only 
 nitrogen remains, the volume of which is read off 
 by the graduation. 
 
 The eudiometric affords greater accuracy. 
 
 It is dependent ii])on the combustion of the oxy- 
 gen with hydrogen in a eudiometer. Air and 
 hydrogen are introduced into the eudiometer, and 
 the electric sjiark then ])assed through the wires. 
 
 All the oxygen in the air combines with a ])ortion 
 of the hydrogen to form water. On cooling, the 
 aqueous vapor condenses and a contraction in 
 volume occurs. Assuming that we had taken 100 
 volumes of air and 50 volumes of hydrogen, and 
 that the residual volume of gas, after allowing for 
 all corrections (p. 121), equaled 87.15; then of 
 150 volumes of mixed gas, 62.85 volumes disap- 
 peared in the formation of water. As the latter results from the 
 union of i volume of oxygen and 2 volumes of hydrogen, the 
 TOO volumes of air employed in the analysis therefore contained 
 6 2^8^ = 20.95 volumes of oxygen. Hence air consists (accord- 
 ing to the determination of Regnault and Bunsen) of 
 
 79.05 volumes Nitrogen. 
 
 20.95 “ Oxygen. 
 
 100.00 “ Air. 
 
 A eudiometer {evbui, fair weather, and iiirpov, measure — for- 
 merly it was thought that there was some connection between 
 the quantity of oxygen in the air and the weather) is an absorp- piQ. 53. 
 tiometer, with two })latinum wires fused into its closed end. 
 
 'The i)assage of the electric spark from wire to wire causes the explosion 
 
 53 )-* 
 
 Numerous analyses show that the composition of the air everywhere on 
 the earth’s surface is constant. The most recent and exhaustive 
 researches of Kreusler, Hempel, Morley, Rayleigh, and Leduc indicate 
 it to be generally as follows: 
 
 * For furtlier .sliidy of gasometric methods consult Rob. Bunsen, (lasometrische Meth- 
 oden, 2 Aull., 1877; Wallher irem])el, Clasanalytische Melhoden, 2 Aufk, 1890; 
 Clemens Winkler, Lehrhuch der Ic^clinischen (lasanalyse, 2 Aull., 1892. 
 
THE ATMOSPHERE. 
 
 I 21 
 
 Nitrogen, 78.06 parts by volume ; 
 Oxygen, 21.00 “ “ “ 
 
 Argon, 0.94 “ “ “ 
 
 75.5 parts by weight. 
 23.2 “ “ “ 
 
 1.3 “ “ “ 
 
 100.00 
 
 100.0 “ “ 
 
 Measuring Gases . — The volume of gases is influenced by pressure, temperature, the 
 moisture contained in them, and to a slight degree by their chemical nature. To com- 
 pare statements of gas volumes, they must be recalculated to normal conditions, i. e., it 
 must be indicated what volume is occupied by the gas at 0° and 760 mm. barometric 
 pressure. The pressure exerted by the mercury column of a barometer at a definite tem- 
 perature depends not only on its height but also upon the intensity of gravity, which in 
 turn varies with the latitude and sea-level. Hence it has been agreed to refer the “nor- 
 mal conditions” to 45° geographical latitude and the sea-level. The normal volume 
 Vo (at 760 mm. and 0°) is calculated, according to Boyle-Gay-Lussac’s laws, from the 
 volume V, which a gas occupies at pressure p and t° by the equation : 
 
 pv = po Vo (l + at) 
 
 to be 
 
 V . p 
 
 760 (l -f at)‘ 
 
 This answers only for dry gases. Any quantity of gas occupies less space when it is dry 
 than when it is moist, because the tension of the aqueous vapor counteracts the atmos- 
 pheric pressure. 
 
 The moisture may be removed by introducing into the gas a ball of coke saturated 
 with sulphuric acid, which dries it. It is more convenient, however, to make the cor- 
 rection of the gas volume in the following manner : Water is brought in contact with the 
 gas to be measured, in order to perfectly saturate it with aqueous vapor ; the gas is then 
 measured and its normal volume calculated by the above formula, after deducting from 
 the observed pres.sure p the number of millimeters corresponding to the tension of the 
 aqueous vapor for the given temperature (p. 90). If the aqueous tension at t® be repre- 
 sented by s (mm. mercury) we finally reach the equation : 
 
 V == V • (P — s) 
 
 ° 760 (i + 0.003665 . t)‘ 
 
 Another form may be given the expression of the two laws relating to gases : 
 pv = Po Vo (i -f Ot), 
 
 if the temperature be counted not from the melting of ice forward but from the absolute 
 zero point of temperature, which can be developed from the following considerations. 
 According to Gay-Lussac the volume of a gas increases, for every degree of rise in tem- 
 
 perature, the 0.00367 or ^ of the volume it occupied at 0°. If the temperature be 
 
 lowered from 0° downward, and the law of Gay-Lussac continues to hold force, then the 
 volume of the gas at — 273° equals zero. This degree of temperature, which has been 
 approached to within 30°, answers for the zero point of ab.solute temperature. It is indi- 
 cated by the letter T. The relation between the degrees of ab.solute temperature and 
 those of ordinary centigrade degrees is expres.sed by t = T — 273. If this be introduced 
 into the preceding erjuation we obtain 
 
 pv ^ 
 
 Pn Vo 
 
 273 
 
 T, 
 
 or, if be made equal to R, then we have 
 ’273 i ’ 
 
 pv = R . T. 
 
 The value of R does not depend upon (he chemical composition of the gas, but .solely 
 upon the units of measure chosen ff)r p and v. d'be law of Avogadro can also be given 
 cxpres.sion through tliis ef|uation if, following Horstmann’s suggestion, consideration be 
 given to the volumes which molecular quantities of the gases occupy, — /. e. the volume 
 
 II 
 
22 
 
 INORGANIC CHEMISTRY. 
 
 occu})ic(l by one molecule or 2.02 f^rams of hydrogen, 32 grams of oxygen, 70.9 grams 
 of chlorine, 36.46 grams of hydrogen chloride. 'I'his volume is, Cf)n.se(iuently, 22.4 
 liters (p. 98). In this case R is the .same for all gases. If the volume be measured in 
 cubic centimeters, and the j)res.sure in grams j)cr .sfiuare centimeter, then, as v = 224(X), 
 p would ecpial 1033.6 (p. 116), T - 273, and R 84800. 
 
 'I'he entire expre.ssion for the three laws would then read 
 
 pv = 84800 T. 
 
 We shall see that this equation also posse.sses great value for solutions. 
 
 The I^oyle-Ciay-Lussac law is not an ab.solutely correct expre.ssion for the behavior of 
 gases. At very great pre.ssures they can be less compres.sed than wovdd correspond to the 
 letpiirements of the law. At low pre.ssures, on the other hand, gases, hydrogen excepted, 
 can be more strongly compressed than the law re(|uire.s. 'J'he.se variations justify the 
 theory propounded by the Hollander, van der Waals ( 1873). As the density of the gas 
 increases the attraction between its molecules becomes greater and the outward ))ressure 
 grows less; indeed, the diminution is inversely proj)nrtional to the .scjuare of the volume 
 of the gas-mass. Accordingly, the observed luessure, when com])ared with that demanded 
 
 l)y the law, is reduced about the value , in which a is a constant corresponding to the 
 
 attraction between the molecules — the cohesion of the gases. On the other hand, the 
 space remaining, with the greater density, for the motions of the molecules is le.ss than 
 the observed volume, becau.se we must deduct from the latter the .s{)ace actually occupied 
 by the molecules themselves. In calculation the ob.served pressure must therefore be 
 increased and the ob.served volume must be diminished. This is indicated in the equa- 
 tion of van der Waals, mentioned on p. 49 : 
 
 (P 4 - “0 (v-b) = RT. 
 or 
 
 (P + ■ vV) ('■ - b) = (I -b a) (I - b) ,* 
 
 which not only answers for gases but also for liquids. The critical data of a gas may 
 also be calculated from it. 
 
 From the great constancy of its composition air was supposed to be a 
 chemical compound, consisting of nitrogen and oxygen. This supposi- 
 tion is, however, opposed by the following circumstances: All chemical 
 compounds contain their constituents in atomic quantities, which is not 
 the case with air. In the mixing of nitrogen and oxygen to form air 
 there is neither disengagement nor absorption of heat, which is always 
 observed in chemical compounds. Further, the air absorbed by water or 
 other solvents jiossesses a composition different from the atmospheric; 
 this is due to the unequal solubilities of nitrogen and oxygen in water. 
 'The air expelled from water iqion application of heat consists of 34.9 
 volumes of oxygen and 65.1 volumes of nitrogen (Bunsen). These facts 
 indicate that air is not a chemical compound, but a mechanical mixture 
 of its two constituents (see Liquid Air). 
 
 I'he great con.staucy iu the comj^osition of our atmo.sjdicre is due chiefly to the fact 
 that there is a constantly renewed mixture produced by the unceasing air currents, by 
 wind.s and storms, rising of the warmer layers and the sinking of those which have 
 become cooler. The mutual diffusion of gases comes, therefore, into consideration. The 
 
 -x-As K 
 
 " ' , a.Hl v= I. 
 
 273 ‘ 
 
GASES RECENTLY DISCOVERED IN THE ATMOSPHERE. 
 
 123 
 
 gas molecules possess, as is now generally acknowledged, a direct, progressive, energetic 
 movement, and can therefore diffuse without limitation into space if by contact with other 
 molecules they are not deflected from or arrested in their course. This is the reason that 
 two gases (or liquids) in immediate contact mix generally with one another. Another 
 kind of diffusion need not be considered so far as concerns the constant 
 composition of the atmosphere, but may be given as a conclusion to what has already 
 been said. If two gases are separated by a permeable diaphragm or membrane, that one 
 possessing the lower density will traverse the septum the more rapidly. 
 
 The following experiment very clearly illustrates this : In the open end of an unglazed 
 clay cylinder (as used in galvanic elements) there is fixed a glass tube about one meter 
 long, its open end terminating in a dish containing water ; the cylinder and tube are filled 
 with air. Over the porous cylinder is placed a wider vessel filled with hydrogen. The 
 latter presses faster into the cylinder than the air escapes from it ; the air in the tube and 
 cylinder is displaced and rises in the water in bubbles. \\'hen the separation of gas 
 ceases, tube and cylinder are almost filled with pure hydrogen. On removing the larger 
 hydrogen vessel the gas will escape much more rapidly into the external air than the 
 latter can enter the cylinder ; the internal pressure will therefore be less than the external, 
 and water ascends in the glass tube. 
 
 In addition to nitrogen and oxygen, air constantly contains aqueous 
 vapor and carbon dioxide (CO2) in very small quantities. The presence 
 of the former can readily be recognized by the fact that cold bodies are 
 covered with dew in moist air. Its quantity depends on the temiterature 
 and corresponds to the vapor tension of water (see p. 90). One cubic meter 
 of air perfectly saturated with aqueous vapor contains 22.5 grams of 
 water at 25° C. ; on cooling to 0° 17.1 grams separate as rain. Gen- 
 erally the air contains only 50-70 per cent, of the quantity of vapor 
 necessary for complete saturation. The amount of moisture in it is either 
 determined according to physical methods (hygrometer), or directly by 
 weighing. To this end a definite quantity of air is conducted through a 
 tube filled with calcium chloride or sulphuric acid, and its increase in 
 weight determined. 
 
 To detect the carbon dioxide in the air, conduct a portion of the 
 latter through solutions of barium or calcium hydroxides, and a turbidity 
 will ensue. To determine its quantity, pass a definite and previously 
 dried amount of air through a weighed potassium hydrate tube, and ascer- 
 tain the increase in weight of the latter. Ten thousand parts by volume 
 of atmospheric air contain, ordinarily, from 3.0. to 4 parts by volume of 
 carbon dioxide. (See Ber. 30 (1897), 1450.) 
 
 Besides the four ingredients just mentioned, air usually contains small 
 quantities of ozone, hydrogen peroxide, ammonium salts (ammonium 
 nitrite), the newly discovered gases argon and heliiu?i, and probably 
 hydrogen. Finally, air contains microscopic germs of lower organisms ; 
 they are generally found in the lower air strata, and their presence influ- 
 ences the processes of decay and fermentation of organic substances. 
 
 GASES RECENTLY DISCOVERED IN THE ATMOSPHERE. 
 
 Lord Rayleigh in 1892, while engaged in an exhaustive research upon the density of 
 the elementary gases, incidentally observed that a liter of nitrogen isolated from the air 
 weighed 1.2571 grams under normal conditions, whereas a liter of the same gas pre- 
 pared from ammonia or nitric acid weighed 1.2507 grams. The determined purpose of 
 ascertaining the reason for this difference in the third decimal led Lord Rayleigh and 
 
124 
 
 INORGANIC CHEMISTRY. 
 
 W, Ramsay in 1894 to the brilliant and snr])risinfi; discovery of argon — a new conslitnont 
 of tlie air, the composition of wliich, after the masterly investigations of emiiK-nt chemists, 
 ai)peared to conceal from us no furtlier enigmas. Vet, in Marclj of 1895, as W. Ramsay 
 sought ft)r additional .sources of argon, lie found heliiim, a second ajiparently elementary 
 gas, which, through a jiortion of its spectrum, had been known as a constituent of the sun. 
 Prior to this it had not been certaiidy known to be pre.se nt on the earth. It is a very 
 subordinate constituent of the air, where it was first ob.served by Kayser. Ramsay 
 obtained it by heating minerals containing uranium. 
 
 The two gases to which reference has just been made differ from all other gases known 
 to us by the total absence of chemical activity. 'Phus far no other substance has been 
 made to react with them. Hence the name Argon, from iiv ipyov, without action. 'They 
 al.so differ from other gases in their real atomic structure ; their atoms have not com- 
 bined to molecules— a condition which heretofore has been ob.served oidy with other 
 substances at very high temperatures (ji. 79). 'Phis is evident from the relation of the 
 specific heat at constant jiressure to tliat at constant volume, which in the case of mon- 
 atomic gases eijuals 1.67, whereas with polyatomic gase.s — tho.se built uj) molecularly — 
 the value lies between i and 5 . Compare (). IT Meyer, 'Phe Kinetic Theory of Ga.ses 
 (2d ed., 1895-1899). 
 
 Recently, Ramsay availed himself of the remarkably developed refrigatory apparatus 
 and found four apparently elementary monatomic ga.ses in air, while seeking for the one 
 between argon and helium. He permitted large quantities of licjuid argon and liquid air 
 to boil away gradually, and then examined the more volatile portions and those portions 
 not so readily volatile and more difficultly condensed. Thus he di.scoverecl in the residue, 
 by the evaporation of 750 c.c. of liquid air to Jo c.c., the gas krypton {Kpv-iTTor, con- 
 cealed), which volatilizes with difficulty, has a density or atomic weight of 45 (O2 — 32), 
 and is characterized in its spectrum by a brilliant red, a yellow, and a green line, the 
 last lying close to that of the northern light ; similarly metargon and xenon [^evoq, 
 foreign), both of which are present in the heavier portions of the air ; and lastly the non- 
 condensible, light neon [veo^, new), rich in lines, with the density 20, and probably the 
 sought-for gas lying between argon and helium. No one of these four ga.ses has been 
 obtained in the pure state ; therefore only argon and helium will be more fully discussed. 
 (See Ramsay, Ber. 31 (1898), 31 ii.) 
 
 ARGON. 
 
 Atom : A = 40. 
 
 The air contains about 0.935 cent, by volume of argon, while “atmospheric nitro- 
 gen ” contains 1.183 per cent, by volume. We inhale daily about 20 liters of this enig- 
 matical gas which has remained concealed for so long a time. It is pre-sent in many 
 mineral waters and gases from springs, e.g., those of Bath, Cauteret, Voslau and Wild- 
 bad. It is liberated in small quantities, however, together with helium, upon heating 
 certain minerals (especially those containing uranium) — cleveite, broggerite, uraninite, 
 and has been obtained in the same manner from one meteorite. Frequently the gases 
 occluded in rock-salt contain argon. 
 
 'Pwo methods may be pursued in separating argon from the air : 
 
 1. Air — freed from oxygen, aqueous vapor and other admixtures — “atmospheric 
 nitrogen” — (p. 119) is conducted over red-hot magnesium filings, which ab.sorb the nitro- 
 gen, forming magnesium nitride : N2 -j- Mg.^ := Mg3N2 (p. ii6), while the argon is scarcely 
 acted upon. The nitrogen can be more quickly absorbed by lithium or by a heated mix- 
 ture of magnesium and calcium oxide or finely divided calcium : 
 
 Mg CaO = MgO -f Ca. 
 
 Ram.say employed this method. 
 
 2. Lord Rayleigh mixed “atmospheric nitrogen,” the mixture of nitrogen and argon, 
 with oxygen and allowed the induction spark to pass continuously through this mixture 
 in the presence of caustic ])ota.sh. Alkaline nitrite is formed in this way from the nitro- 
 gen, oxygen and alkali (sec Nitrous Acid, ])]X 205, 207), while the residue of argon and 
 oxygen is conducted over ignited co]q)er to remove the latter. Cavendi.sh, as early as 
 1785, observed that there remained .some “ atmo.s])heric nitrogen ” which was not con- 
 verted into alkaline nitrile ; he did not, however, pursue this observation further. 
 
COMPOUNDS OF NITROGEN WITH HYDROGEN. 
 
 125 
 
 The argon tlius prepared contains helium and neon as the lighter and inetargon, 
 krypton and xenon as tlie heavier admixtures. It is then liquefied by means of licpiid 
 air and purified by fractional distillation. Its density referred to C)^ 32 is 39.914 or in 
 
 round numbers 40 ; it is identical with its atomic weight. Under normal conditions a 
 liter of argon weighs 1.780 grams. At very low temperatures it congeals to an ice-like 
 mass, which melts at — 189.5° ^o a colorless liquid. It boils at — 185° ; its critical 
 temperature is — 121° and its critical pressure 50.6 atmospheres. Argon is approxi- 
 mately times as soluble in water as nitrogen : 100 volumes of water dissolve 4 
 
 volumes of argon at 12°. Hence argon accumulates in the gases of rain-water. The 
 spectrum distinguishes argon very certainly from nitrogen and other substances [com- 
 pare : Z. f. phys. Ch. 1895, 16, 344; Z. f. anorg. Ch. 18 (1898), 222 ; Chem. Central- 
 blatt 70 (1899), I, 469; also Mugdan : Argon and Helium, two new gaseous ele- 
 ments, Stuttgart, 1896]. 
 
 HELIUM. 
 
 Atom : He = 4. 
 
 Helium is as inactive chemically as argon. While the latter is widely distributed on 
 the earth, helium is one of the rarest of terrestial substances. It occurs, however, with 
 hydrogen in immense masses in the photosphere of the sun and other brilliant fixed 
 stars. Norman Lockyer detected it as early as 1868 by means of the spectroscope in the 
 chromosphere of the sun, and in 1869 he and Frankland named it Helium. In 1892 
 Palmieri observed its terrestial occurrence for the first time while studying spectroscopically 
 a substance which had been thrown out by Vesuvius. But it was first in March, 1895, 
 that W. Ramsay found helium while he was engaged in seeking for sources of argon. 
 Hitherto only a portion of the helium spectrum (line Dg) had been known. The gas 
 helium (together with hydrogen, carbon dioxide, nitrogen and probably also argon) is 
 evolved when certain rare minerals, usually consisting of salts of uranium, yttrium and 
 thorium (^. g., cleveite, uraninite, broggerite, monazite) are heated alone with dilute sul- 
 phuric acid (i : 8) or with sodium bisulphate. It was similarly obtained from a meteorite. 
 Kayser found it, in small quantities it is true, in the air and the later investigations of 
 Ramsay and Travers have confirmed this observation. Helium, like argon, occurs in the 
 gases from springs, e. g., in those from Wildbad, in that of Adano near Padua, as well 
 as in the gaseous exhalations of Tuscany. 
 
 Helium is monatomic (see above). Its atomic weight (its density) equals about 4. 
 Dewar has shown that it can be liquefied by means of liquid, vaporizing hydrogen, so 
 that at present all gases, with the exception of neon, can be liquefied. It is less soluble 
 in water than any other gas : loo volumes of water at 18° take up but 0.73 volume of 
 helium. It is as inactive as argon. Travers asserts that when a powerful electric dis- 
 charge is sent through a Pliicker tube (see Spectrum Analysis) filled with helium the latter 
 is absorbed by the platinum electrodes and is again liberated from the same on the appli- 
 cation of heat. This recalls the behavior of hydrogen and palladium ; it may be useful 
 in separating helium and argon. Five brilliant lines are prominent in the spectrum of 
 helium : one each in the red, the yellow, the green, the blue and the violet. The yellow 
 line Dg lies close to the two sodium lines, Dj and D^, toward the violet end of the 
 .spectrum. 
 
 COMPOUNDS OF NITROGEN WITH HYDROGEN. 
 
 Ammonia is the most important compound of nitrogen and hydrogen. 
 It has l^een known for tlie longest time. Since 1889 Th. Curtins has 
 added four otlier derivatives of these two elements to it : hydrazine, N^H^, 
 and hydrazoic acid, N.Jl, as well as the ammonium and hydrazine salts 
 of the latter: NJI, (= NHg + NgM)and NHI, (=N.^H, + NgH). See 
 Per. 29 ( 1896), 759. Oxvaminonia or hydroxylaminc, NHgO, discovered 
 by W. J.ossen in 1865, will be discussed after ammonia. 
 
26 
 
 INORGANIC CHEMISTRY. 
 
 1. AMMONIA. 
 
 Mulcculc : NUa :^ 17.07. 
 
 Ammonia occurs in tlie air in combination with some acids, in natural 
 waters and in the eartli, but always in small (juantities, Ihiestley first 
 studied it carefully and called it alkaline air. berthollet determined its 
 composition in 17CS5. d'lie formation of ammonia by the direct union 
 of nitrogen and hydrogen occurs under the inlluence of thesilent electric 
 discharge. Its comiiounds — ammonium salts — are freciuently ])roduced 
 under the most varying conditions. Thus ammonium nitrate is formed 
 by the action of the electric spark iijion moist air: 
 
 K, + O + 2ll,0 NII^NOg. 
 
 Aminoiiium 
 
 nitrate. 
 
 Ammonium nitrite, NH^NOj, is said to be formed in every combus- 
 tion in the air; and in the electrolysis of water. Further, ammonium 
 salts are ])roduced in the solution of many metals in nitric acid, in con- 
 setpience of a reduction of the acid by the liberated hydrogen : 
 
 HNO3 + 81 1 311,0 + NIIj. 
 
 The following conversion of nitrogen into ammonia deserves considera- 
 tion : Nitrogen unites with magnesium at a red heat to magnesium nitride, 
 MggN,, and the latter is energetically decomposed by water with the 
 evolution of ammonia: 
 
 MggN, + 311,0 = 3MgO + 2NII3 (p. 124). 
 
 Nitric acid in the form of salts, in alkaline solution, is reduced by 
 nascent hydrogen to ammonia (see Nitric Acid). 
 
 Ammonia is produced in large quantities in the decomposition and 
 dry distillation of nitrogenous organic substances. Even as late as the 
 last century the bulk of the ammonium chloride (the most important salt 
 technically), was obtained from camel’s dung or decayed urine. Its orig- 
 inal name Sa/ armoniacimi — Armenian salt — was confounded later with 
 the designation for Egyptian rock-salt — Sal ammoniacum. In the prepa- 
 ration of illuminating gas by the distillation of coal, ammonia appears as 
 a by-i)roduct and may be obtained by combining it with sulphuric or 
 hydrochloric acid. This method is used almost exclusively at present for 
 its production. 
 
 To ])re[)are ammonia heat a mixture of ammonium chloride and slaked 
 lime in a glass or iron flask : 
 
 2NIFCI f Ca(On), == CaCl, + 2 ll ,0 + 2NH3. 
 
 Aininoiiium Calcium 
 chloride. hydroxide. 
 
 d'he disengaged ammonia gas is collected over mercury, as it is readily 
 soluble in water (see ]>. 58, Fig. 33). For ])erfect drying conduct it 
 through a vessel filled with burnt lime (CaO). ('alcium chloride is not 
 applicable for this ])urpose, as it combines with the gas. In consequence 
 of its levity, ammonia, like hydrogen, may be collected by disi)lacing the 
 air in inverted ves.sels. 
 
AMMONIA. 
 
 127 
 
 Physical Propc 7 'ties . — Ammonia is a colorless gas with a suffocating, 
 characteristic odor. Its density is 0.59 (air = i). Under a pressure of 
 6.5 atmospheres (at 10° C.), or by cooling to — 40° C., under ordinary 
 pressure it condenses to a colorless mobile liquid with a specific gravity 
 of 0.623 at 0°, solidifies at — 85°, and melts again at — 75°. Liquid 
 ammonia has recently been introduced into commerce. 
 
 Ammonia gas may be condensed, just like chlorine. Take ammonium silver chloride 
 (AgCl . 2NH3), obtained by conducting ammonia over silver chloride, and enclose it in 
 a tube with a knee-shaped bend (p. 51, Fig. 32). 'I'he limb containing the compound is 
 now heated in a water-bath, while the other limb is cooled. The compound is decom 
 posed into silver chloride and ammonia, which condenses in the cooled limb (h'aradayj. 
 
 Ammonia gas dissolves very readily in water, with the liberation of 
 heat. One part of water at 0° and 760 mm. pressure absorbs 1146 volumes 
 (= 0.875 parts by weight); at 20°, 739 volumes (= 0.526 parts by 
 weight) of ammonia. At 16° and 760 mm. pressure 100 parts by volume 
 of water absorb 60 parts by weight of ammonia. The specific gravity of 
 a 34.95 per cent, solution at 15° is 0.882; henc e a liter of it contains 
 308.3 grams of ammonia. When a long glass tube, closed at one end 
 and filled with ammonia, has its open end placed in water, the latter 
 rushes up into the tube as it would into a vacuum ; a ])iece of ice melts 
 rapidly in the gas. The aqueous solution possesses all the properties of 
 the free gas, and is called Liquor ammonii caustici. d'he greater the 
 ammonia content the less will be the specific gravity of the solution. All 
 the gas escapes on the application of heat. 
 
 When the condensed liquid ammonia evaporates it absorbs a great amount of heat, 
 and answers, therefore, for the production arti- 
 ficially of cold and ice in Carre’s apparatus. 
 
 The simplest form of the latter is represented in 
 P'ig. 54. The iron cylinder A is filled about 
 half with a concentrated aqueous ammonia solu- 
 tion, and is connected, by means of the tubes 
 from b, with the conical vessel Aj in the middle 
 of which is the empty cylindrical space E. The 
 entire internal s])ace of A and A is hermetically 
 shut off. A is heated upon a charcoal fire until 
 the thermometer a, in it, indicates 130° C., 
 while A' is cooled with water. In this way the 
 gaseous ammonia is expelled from the aqueous ^ 
 solution in A^ passes through b, in which most 
 of the water runs back, and condenses to a 
 liquid in B, of the receiver F. The cylinder A 
 is removed from the fire, cooled with water and 
 the vessel A>, constructed of thin .sheet-metal 
 and filled with water, placed in the cavity A, FiG. 54. 
 
 which is surrounded with a poor conductor, 
 
 e. g., felt. The ammonia condensed in B evaporates, and is reab.sorbed by the water in 
 A. Ily this evaporation a large fjuantity of heat, withdrawn from A’ and its surroundings, 
 becomes latent ; the water in 1 ) freezes. 
 
 'I'he method of Carre for the artificial production of ice has acquired great apjdication 
 in the arts; recently, however, ice machines have been introduced, 'i'hese are driven 
 by liquid ammonia (Lindel, or by licjuid sulphur dioxide and carbon dioxide (Pictet) 
 (compare pp. 48, 117). The method of Windhausen, depending upon the expansion 
 of compressed air, is much used. 
 
128 
 
 INORGANIC CIIEMISIRY. 
 
 Chemical Properties . — A red heat or the continued action of the elec- 
 tric spark decomposes ammonia into nitrogen and liydrogen. On con- 
 ducting ammonia gas over heated sodium or potassium, tlie nitrogen 
 combines with these metals and hydrogen escapes : 
 
 NII3 -f- 3K NK3 -p 3li. 
 
 Magnesium unites, when heated in an atmosphere of ammonia, with 
 the nitrogen of the latter, giving a bright light: 
 
 3Mg -I- 2NII3 _ Mg3N3 }- 3II3. 
 
 Ammonia will not burn in the air; in oxygen, however, it burns with 
 a yellow flame : 
 
 2NIl3 + 3().= N3 + 3ll30; 
 
 ammonium nitrite and nitrogen dioxide are formed simultaneously. 
 When a mixture of ammonia and oxygen is ignited it burns with explo- 
 sion. 
 
 To show the combustion of ammonia in oxygen, proceed as follows: A 
 glass tube, through which ammonia is conducted, is brought with oxygen 
 into a vessel, bringing the opening of the latter near a flame at the moment 
 of the introduction of the glass tube. In contact with oxygen, the am- 
 monia gas ignites and continues to burn in it. 
 
 The following experiment (of Kraut) shows the combustion of ammonia 
 very conveniently. Place a somewhat concentrated ammonia solution in 
 a beaker glass; heat over a lamp, until there is an abundant disen- 
 gagement of gas, and then run in oxygen, by means of a tube dipped 
 into the liquid. Upon approaching the mixture with a flame, it ignites 
 with a slight explosion. The ignition may be induced without a flame, 
 by sinking a glowing platinum spiral into the mixture ; we then have a 
 number of slight explosions. The glass is filled at the same time with 
 white vapors of ammonium nitrite (NH^N02) ; later, when oxygen pre- 
 dominates, red vapors of nitrogen dioxide (NO^) and nitrous acid 
 appear. 
 
 If chlorine gas be conducted into the vessel with ammonia, it immedi- 
 ately ignites and continues to burn in the latter, with the production of 
 white fumes of ammonium chloride (NH^Cl). The chlorine combines 
 with the hydrogen of the ammonia, with separation of nitrogen, and 
 yields hydrochloric acid, which unites, with the excess of ammonia, to 
 form ammonium chloride : 
 
 NII3 -f 3CI = 3riCl + N 
 
 and 
 
 3NII3 ~p 3IICI 3 NII,C 1 . 
 
 Chlorine reacts similarly u] 3 on acpieous ammonia (p. ti6). 
 
 In gaseous form, as well as in solution, ammonia possesses strong basic 
 ])roj)erties ; it bines red litmus jiaper and neutralizes acids, forming salt- 
 like comj)onnds with them, which are very similar to the salts of the 
 alkalies — sodium and potassium. 'The following illustrates the similarity. 
 
QUANTITATIVE COMPOSITION OF AMMONIA. 
 
 I 29 
 
 NII3 
 
 + 
 
 IICl 
 
 = NII^Cl 
 
 Ammonium 
 
 chloride. 
 
 KCl. 
 
 Potassium 
 
 chloride. 
 
 2NH3 
 
 + 
 
 H2SO4 
 
 = (Nil,)., SO, 
 
 Ammonium 
 
 sulphate. 
 
 K3SO,. 
 
 Potassium 
 
 sulphate. 
 
 NHj 
 
 + 
 
 H3S 
 
 = NII,SH 
 
 Ammonium 
 
 sulphydrate. 
 
 KSH. 
 
 Potassium 
 
 sulphydrate. 
 
 In these ammonia derivatives NH^ plays the role of a metal. Hence, 
 to show its similarity to sodium, potassium and other metals, the group 
 (NH^) has been designated ammonium and its compounds, ammonium 
 salts. The latter, when acted on by strong bases, yield ammonia gas : 
 
 2NH,C1 + CaO = 2NH3 + CaCb + 
 
 The metallic character of the ammonium group is also confirmed by the 
 existence of the ammonium amalgam. Therefore, the ammonium deriva- 
 tives will be considered with the metals. 
 
 Thermo-chemical Deportjnent. — The heat of formation of ammonia from 
 hydrogen and nitrogen equals 12 Cal. When ammonia gas is dissolved 
 in much water 8.4 Cal. are set free, so that the heat of formation of 
 ammonia from its elements in dilute aqueous solution equals 20.4 Cal. : 
 
 (N,H3 - gas) = 12. (NH3,Aq) = 8.4. (N,H3,Aq) = 20.4. 
 
 The great heat of solution of gaseous ammonia explains why ice will melt in the 
 same {p. 127). 
 
 The explosibility of a mixture of ammonia and oxygen is accounted for by the follow- 
 ing great heat disengagement : 
 
 2NH3 + 30 = 3H3O + N, . . (+ 147.6 Cal.) 
 
 (24 Cal.) (3 X 57-2 Cal.) 
 
 The action of chlorine upon gaseous or aqueous ammonia is also very energetic : 
 
 NH3 gas + 3CI = 3HCI gas + N . . . (+ 54 Cal.) 
 
 (12 Cal.) (66 Cal.) 
 
 NIl3-dissolved + 3CI = 3TICl-dis.solved + N . . . (+ 97.5 Cal.) 
 
 (20.4 Cal.) (3X39-3 Cal.) 
 
 When there is an excess of ammonia the hydrochloric acid combines with it to form 
 ammonium chloride (NII3 -f- HCl = NII^Cl), and the heat disengagement is thereby 
 further increased. 
 
 QUANTITATIVE COMPOSITION OF AMMONIA. ATOMIC WEIGHT OF 
 
 NITROGEN. 
 
 The quantitative analysis of ammonia shows that it consists of i.oi 
 parts of hydrogen and 4.68 parts of nitrogen; hence we conclude that 
 the atomic weight of nitrogen is 4.68 or a multiple of it (see p. 69) : 
 
 II = I 
 N = 4.68 
 
 2II = 2.02 
 
 N = 9.34 
 
 NII2 - 11.36 
 
 3II = 3.03 
 
 N = 14.04 
 
 Nil 
 
 5.68 
 
 Nil, 
 
 17.07 
 
T30 
 
 INORGANIC CHKMISTRV. 
 
 As tlic density of ammonia ccjiials i 7.07 (O.^ — 32), its molecular weight 
 would e(}ual 17.07. In 17.07 parts of ammonia there are 3.03 parts, and, 
 therefore, three atoms of hydrogen. 'That the 14.04 partsol nitrogen united 
 with tliem corres])ond to one atom of nitrogen is a consc-(|uence, as never 
 less than 1^.04 parts of that element are present in the niotecnlar weije;ht of 
 any nitrogen derivative. The density of nitrogen ecjuals 28.08 ( Oj = 32 ) ; 
 therefore, the molecule of nitrogen consists of two atoms (N^). 'This is 
 also concluded from the volume ratios, as we shall soon see, occurring in 
 the decomposition of ammonia (comi)are p. 76). 
 
 From the molecular formulas NH., and N2 it follows, further, that i 
 volume of nitrogen and 3 volumes of hydrogen form 2 volumes of am- 
 monia gas, or that 2 volumes of ammonia decompose into 3 volumes of 
 hydrogen and i volume of nitrogen, corresponding to the molecular 
 equation : 
 
 N2 + 3H2 = 2NII3. 
 
 I vol. 3 vols. 2 vols. 
 
 The following exiieriments jirove these conclusions: 
 
 1. Decompose an acpieous ammonia solution, mixed with salt (NaCl) 
 to increase its power of conductivity, in a Hofmann’s apparatus (j). 77), 
 by the galvanic current. Hydrogen will sei)arate at the negative 
 and nitrogen at the positive ])ole; the former will have three times the 
 volume of the latter as soon as the solution is saturated with gases. 
 
 2. Pass electric (induction) sparks through dry ammonia gas contained 
 in a eudiometer, or the apj^aratus represented in Fig. 47 (p. 98). In this 
 way the ammonia is decomposed into nitrogen and hydrogen, the volume 
 of which is twice as large as that of the ammonia employed. That 3 
 volumes of hydrogen are present in the mixture for every volume of nitro- 
 gen is easily shown by the eudiometric method, by burning the hydrogen 
 with oxygen (p. 99). 
 
 2. HYDROXYLAMINE (OXYAMMONIA). 
 
 NH3O = NHoOH. 
 
 This compound was discovered (by Lossen in 1865), in the reduction 
 of ethyl nitrate by tin and hydrochloric acid, in the form of its salts and 
 in aqueous solution. Lobry de Bruyn first obtained it anhydrous and 
 in a solid form in 1891. It is produced, too, by the action of tin upon 
 dilute nitric acid, and by tin and hydrochloric acid upon all the oxygen 
 compounds of nitrogen. In all these reactions it is the hydrogen elimi- 
 nated by the tin which, in statu 7 iaseendi, reduces the nitric acid : 
 
 IINO3 -f 311.2 =3 NH3O -f 2H2O. 
 
 To prepare hydroxylamine treat ethyl nitrate f 120 grams) with granulated tin (400 
 grains) and hydrocddoric acid (800-1000 c.c. of specific gravity 1.19, mixed with three times 
 its volume of water). 'I'he metal should he completely di.s.solved. The .solution diluted 
 to twice its volume is treated with hydrogen .suli)hide to precipitate the tin. The filtrate 
 from llie tin sul|)hide is evaporated and the hydroxylamine hydrochloride, NII3O . IICl, 
 extracted from the residue with hot alcohol. 
 
 I lydroxylamiiie hydrochloride is most easily formed by the interaction of hydrochloric 
 acid and fulminating mercury (see Organic Chemistry). 
 
HYDROXYl. AMINE. 
 
 I3I 
 
 Ilydroxyhmiine suli)liate is technically prepared by healing potassium hydroxylamine- 
 disulphonate with water to 100-130°, when hydroxylainine sulphate and potassium sul- 
 phate are produced [see Raschig, Ann. Chem. 241 (1887), 161] : 
 
 NII 0 (S 03 K )2 H 2II2O = NIT3O. II^SO^ -f K^SCV 
 
 Potassium hydroxylamine-disulphonate is formed by the interaction of potassium nitrite 
 and acid potassium sulphite in the cold : 
 
 KNO.2 + 2KHSO3 = NH0(S03K)2 -f KOH. 
 
 To obtain solid hydroxylainine decompose the hydrochloride in methyl alcohol solution 
 by a corresponding quantity of sodium methylate ; hlter from the sodium chloride which 
 separates and distil the filtrate under greatly reduced pressure [Briihl, Ber. 26 (1893), 
 III, 2508; also Lobry de Bruyn, ibid. 27 (1894), i, 967]. 
 
 Hydroxylamine crystallizes in colorless, odorless needles. It melts at 
 33° and under 22 mm. pressure boils at 58°. Its specific gravity equals 
 1.235. rather stable at temperatures below 15°, but above that it 
 
 rapidly breaks down into nitrogen, nitrous oxide, nitrous acid and 
 ammonia. At about 130° and often at much lower temperatures the 
 decomposition is accompanied with explosion. It absorbs moisture 
 rapidly on exposure to the air. 
 
 Hydroxylamine is very similar to ammonia, and like it unites directi}- 
 with the acids to form salts : 
 
 NH3O + HCl = NH3O. HCl = NH3(0H)C1. 
 
 Hydroxylamine hydrochloride in distinction to ammonium chloride is 
 soluble in ale 'hoi. It passes into ammonium chloride when allowed to 
 stand exposed to the air. 
 
 On adding to the aqueous solution of the sulphate of hydroxylamine 
 sufficient barium hydroxide to remove all the sulphuric acid, an aqueous 
 solution of hydroxylamine is obtained, which, like the ammonia solution, 
 possesses strong basic properties, and colors red litmus-paper blue. The 
 solution is, however, very unstable, and readily decomposes into water, 
 ammonia, and nitrogen : 
 
 3NH3O NH3 + 3H2O + N2. 
 
 Upon the application of heat a portion of the hydroxylamine will be 
 carried over undecomposed along with the steam, but most of it is 
 decomposed. The hydroxylamine solution manifests a strong reducing 
 action ; it preci])itates metallic silver from silver nitrate, white mercurous 
 chloride, Hg2Cl2, from mercuric chloride, HgCl2, and cuprous oxide from 
 cupric salts. 
 
 Owing to its great similarity to ammonia and its various reactions, it is 
 supposed that hydroxylamine represents ammonia in which an hydrogen 
 atom is rejilaced by the hydroxyl grouji OH; it is therefore a compound 
 of the latter group with the amido grou]) : 
 
 NII3O = NII2. OH. 
 
 3. Diamide, or Hydrazine, N2TI4 ^ H2N . Nir2, a compound of two amido-groups 
 (NII2), was until 1889 only known in its numerous organic derivatives. Since then Cur- 
 tius and Ids co-workers have exhaustively investigated a large number of its inorganic 
 compounds, and in 1895 Bobry de Bruyn succeeded in getting the free diamide or hydra- 
 
132 
 
 INORGANIC CHEMISTRY. 
 
 ziiie unknown until then [Her, 28 (1895), 30 ^ 5 J- salts were first ma«le hy an indirect 
 
 method from certain orjranic ‘ nitrogen derivatives: hy the decomposition of dia/.o- or 
 triazo-acetic acid uj)on digesting them with water or mineral acids [Curtins and Jay, J. f. 
 prakt. Ch. [2] 39 (1889), 27J ; by heating amidoguanidine with sodium hy<lr(ixi<le 
 [ Thiele, Ann. Chem. 270 (1892), 31], etc. .See also the last edition of Richter’s 
 Organic Chemistry. Very recently diamide has been made fnmi inorganic compcnmds ; 
 by the reduction of nitric oxide-potassium sulphite with nascent hydrogen : 
 
 K2SO3. N2O2 -f 3II.2 = N.2II4. ll.p -f K^SO^ 
 
 [Duden, Ber. 27 (1894), 3498]. 
 
 Free hydrazine, like free hydroxylamine, is prepared by decomposing its hydrocliloride 
 with sodium methylate : 
 
 N2II,. IICl + CIIgONa = NaCl + CII3OII -f N^H,. 
 
 It may also be obtained by distilling its hydrate under reduced pressure with barium 
 oxide. It is a colorless liquid, fuming strongly in the air, boiling at 113.5° .solidify- 
 ing at 0°. It melts at 1.4°. Its specific gravity is about 1.003 23°. It is not explo- 
 sive and is stable even at 300°, Its hydrate, II./J, is a strong base, very much 
 
 like ammonia. It is ])roduced when the hydrazine salts are decompo.sed by powerful 
 bases. It is characterized by great reducing power. It will be discus.sed later along 
 with the diammonium salts in concluding the alkaline earth metals. 
 
 4. Hydrazoic Acid, or Azoimide, N3II = |1 /Nil, was also di.scovered by Cur- 
 
 tins in 1890 [Ber. 23 (1890), 3023]. It was, like the preceding body, first isolated 
 from organic derivatives, and indeed from tho.se of diamide [especially benzoyl azoimide 
 which upon digestion with sodium alcoholate yields sodium azoimide and ethyl benzoate : 
 
 CeH5CON3 + C^IipNa NaN3 + C.II.COOC^IIs, 
 
 and from diazohippuramide which is resolved by ammonia into ammonium azoimide, 
 hippuramide and ammonia : 
 
 C6H5CO-nh-ch2-co-nh-n=n-oh -f 2NH3 = 
 
 N3. NH^ -f H2O -h C6li5CO-NH-CH2-CO-Nil2. 
 
 [A description of these investigations can be found in the Jahrbuch der Chemie I (1891), 
 91, 227 ; II (1892), 229 ; also Thiele, Ann. Chem. 270 (1892), 53, and Richter’s Organic 
 Chemistry.] 
 
 The following are real inorganic methods for the preparation of hydrazoic acid or its 
 salts. 
 
 Curtius prepares a dilute aqueous solution of the acid by chilling with ice the red 
 vapors arising from heating arsenious oxide and nitric acid (see Nitrous Acid) and 
 adding the resulting blue liquid gradually, as long as gas is not evolved, to a cold solution' 
 of hydrazine. This is analogous to the formation of silver azoimide (Angeli) from silver 
 nitrite and hydrazine sulphate : 
 
 N==N 
 
 H2-N-NH2 -b NO2H 1= H 2 -N-N=N- 0 H + H2O = \ / -f 2H2O. 
 
 Nil 
 
 Wislicenus as.serts that .sodium azoimide is readily obtained by heating sodamide to 150- 
 240° in a current of nitrous oxide : 
 
 NIl2Na -f N.2O = NaNg + II.2O. 
 
 Ililute aiiueous .solutions of hydrazoic acid may be obtained by di.stilling the metallic 
 azoimidcs with dilute sul])lmric acid. By further distillation of the solution and by 
 means of calcium chloride the acid may be obtained anhydrous when it is a colorless, 
 
NITROGEN CHLORIDE. 
 
 133 
 
 mobile liquid with an unbearable odor. It boils at 37°, volatilizes with steam, 
 and is characterized by its exceptionally violent explosibility. In this respect it differs 
 from the haloid acids, especially hydrochloric acid, with which it possesses many points 
 of similarity. This behavior would indicate that hydrazoic acid is a strongly endothermic 
 body : 
 
 (N3,H) dissolved = —62 Cal. 
 
 Its alkali salts are not explosive. Some of the alkaline earth .salts explode when heated, 
 and some of the heavy metals w’hen touched or struck. These salts will be further 
 treated under the respective metals. [See Curtius and Rissom, Jr. prakt. Ch. 58 (1898;, 
 261 ; also Z. f. anorg. Ch. 17 (1898), 18.] 
 
 COMPOUNDS OF NITROGEN WITH THE HALOGENS. 
 
 NITROGEN CHLORIDE. 
 
 NCI3. 
 
 As we have seen, nitrogen is liberated when chlorine acts upon an 
 excess of ammonia (p. 116) ; when, however, the chlorine is in excess, it 
 acts upon the previously formed ammonium chloride to produce nitrogen 
 chloride : 
 
 NH,C 1 -f 3CI2 = NCI3 + 4HCI. 
 
 For the preparation of a small quantity of nitrogen chloride, dip a 
 flask filled with chlorine, open end down, into an aqueous ammonium 
 chloride solution warmed to 30°. The chlorine is absorbed, and heavy 
 oil drops separate, which are best collected in a small leaden dish. They 
 contain hydrogen in addition to nitrogen and chlorine. They yield 
 pure nitrogen chloride by a second treatment with chlorine [see Gatter- 
 mann, Ber. 21 (1888), 751]. A pentachloride of nitrogen, NCI5, is not 
 known. 
 
 Nitrogen chloride is an oily, yellow liquid, with a disagreeable odor; 
 its specific gravity equals 1.65. It does not solidify even at — 40°. It 
 is an extremely dangerous compound, as it decomposes by the slightest 
 contact with many substances, when the temperature exceeds 90° and by 
 the action of direct sunlight. Its decomposition is accompanied by an 
 extremely violent report. Its solution in benzene, ether and carbon 
 bisulphide can be handled with comparatively little danger; but it 
 rapidly decom])oses in them. Aqueous ammonia gradually decomposes 
 it into ammonium chloride and nitrogen : 
 
 NC]3 + 4NH3 = 3 NH,C 1 + N,. 
 
 It is converted into ammonium chloride and free chlorine by concen- 
 trated hydrochloric acid : 
 
 NCI3 -f 4 TIC 1 = NIbCl -f 3Cb. 
 
 This reaction is directly opposed to that by which nitrogen chloride is 
 formed. 
 
34 
 
 INORGANIC CIIF.MISTRY. 
 
 Nitrogen cliloride is (lccoini)osed l)y water jjartly into liypochI(j- 
 roiis acid [see Jkr. 27 (1894), i, 1017, and 30 (1897), ii, 1434, 1792; 
 III, 2642]. 
 
 'File formation and explosihility of nitrogen cliloride may be illustrated in a harmless 
 way as follows: Decompose a saturated ammonium chloride solution with the electric 
 current. Nitrogen chloride rising in small drops from the li(iuid will sejiaratc at the 
 ])ositive pole. Upon covering the surface of the solution with a thin layer of turpentine 
 oil, each drop will explode as it comes in contact with the latter. 
 
 Nitrogen Iodide. Uiion adding aininonimn liydroxide, ora mix- 
 ture of ammonium chloride and caustic soda, to a soltition of iodine in 
 atjneotis potassium iodide, a brownish-ldack powder separates. Its com- 
 position closely approximates the formula, NIII^. Its formation by 
 means of ammonium chloride and caustic soda is rejiresented in the 
 equation : 
 
 4I + NIbCl 4- sNaOII = NIII2 + 3II2O 4- NaCl j 2NaI. 
 
 [Com]iare Raschig, Ann. Chem. 230 (1885), 212.] 
 
 When the conditions are slightly changed a very similar compound 
 separates. Its formula is N2l3H3(= Nil, -f NI3). Protracted washing 
 with water decomposes it into ammonia and nitrogen tri-iodide, NI,. 
 
 Nitrogen Di-iodide and Nitrogen Tri-iodide, NHIj andNl,, are, 
 when dry, very explosive. The explosihility may be shown without 
 danger in the following manner: The precipitate is collected on a filter, 
 washed with water, the filter opened outand torn into small pieces, which 
 are then allowed to dry; upon the slightest disturbance these pieces 
 explode with a sharp report. 
 
 Szuhay claims that nitrogen iodide, NHI2, behaves like a feeble acid 
 and that its silver salt, NAgl2, explodes [Ber. 26 (1893), ii, 1933]. 
 
 Nitrogen iodide dissolves in dilute hydrochloric acid and decomposes 
 into ammonia and iodine monochloride : 
 
 NII2I + HCl = NH3 4- ICl. 
 
 H\ drogen sulphide and sulphurous acid convert it into ammonia and 
 hydrogen iodide. 
 
 'Fhe nitrogen iodide formed by digesting powdered iodine with ammo- 
 nia water manifests properties which are slightly different from those of the 
 ordinary iodide. It is only stable in the presence of ammonia. It some- 
 times explodes even when moist — if it be washed with water, or when 
 acted ui)on by hydrochloric acid. 
 
 TJie 7 'mo-cheniical Deporfmenf . — Nitrogen chloride and iodide are both strongly endo- 
 thermic comjKHinds ; considerable heat is absorbed in their production from the elements, 
 d'his ngain ])resup])oses conditions under which the nitrogen halides would be no longer 
 stable. ( ’onsc(|uently their heat of formations cannot be obtained by direct measurement 
 but must be calculated from other thermo-chemical relations. That for nitrogen chloride is 
 
 (N,Cl3) = -41.9 Cal. 
 
PHOSPHORUS. 
 
 135 
 
 2. PHOSPHORUS. 
 
 Atom: P = 3i.o. Molecule: P4= 124.0. 
 
 This element does not occur free in nature, because of its very great 
 affinity for oxygen. The phosphates, especially calcium phosphate, are 
 widely distributed. By the disintegration of the minerals containing 
 phosphates the latter pass into the soil, are absorbed by plants, and these 
 are consumed by animals. Hence phosphates occur in both plant 
 and animal ashes. The bones of animals consist chiefly of calcium 
 phosphate. 
 
 Brand, in Hamburg, and Kunkel, in Wittenberg (1670), obtained 
 phosphorus almost simultaneously by the ignition of evaporated urine. 
 In 1755, Scheele, in Sweden, showed that it could be obtained from 
 bones, in which Gahn had found calcium phosphate. Its name is derived 
 from its power of giving light in the dark — (puxTipopoq, light-bearer. 
 
 Phosphoru.s is prepared in France and in England almost exclusively in accordance 
 ■with a method first suggested in 1778 by Nicolas and Pelletier, which method was 
 subsequently improved by Fourcroy and Vauquelin. 
 
 The bones are burned, thereby destroying all organic admixtures, or they are first 
 treated with .superheated steam to remove the bone glue. The bone ashes consist princi- 
 pally of tertiary calcium phosphate Ca3(P04)2 (see Phosphoric Acid). They are dige.sted 
 with two-thirds of their weight of concentrated sulphuric acid, when the tertiary phos- 
 phate becomes primary calcium phosphate, and gypsum (calcium sulphate) is produced : 
 
 Ca3(P04)2 + 2H2SO4 = CaH,(POd2 + 2CaS04. 
 
 Tertiary Primary Calcium 
 
 calcium phosphate. calcium phosphate. sulphate. 
 
 The gypsum, which dissolves with difficulty in water, is separated from the readily 
 .soluble primary phosphate by filtration ; the solution is mixed with charcoal, evaporated 
 in leaden pans, and the residue raised to a red heat in clay retorts. This expels water 
 from the primary phosphate and the latter changes to calcium metaphosphate : 
 
 CaH,(POj2 = Ca(P 03 b + 211^0. 
 
 Calcium 
 
 metaphosphate. 
 
 The ignited residue is then raised to a white heat. The carbon partly reduces the 
 metaphosphate to ])hosphorus, by forming carbon monoxide with oxygen. The resulting 
 calcium oxide in like manner converts a portion of the metapho.sphate into tertiary phos- 
 phate, which is not further attacked by the carbon. The following equations show the 
 changes : 
 
 I. 2Ca(P03).2 + loC = P, f loCO + 2CaO 
 II. Ca(P03).2 4 - 2CaO Ca3(POb2 
 
 3Ca(P03).2 + loC = P4 -f- Ca3(P04)2 4 loCO. 
 
 Recently the great heat of the electric furnace has been apjdied in carrying out a 
 method suggested by Wohler, but which has not until the pre.sent been made operative : 
 the tertiary pho.sphate, mixed with .sand, is reduced by carbon. The silicic acid (sand) 
 acts like the .sulphuric acid, in that it liberates phosphoric acid when heated : 
 
 2Ca3(PO^)2 -|- loC -(- 6Si02 = 6CaSiO., -f- loCO 4 4 B 
 
 Mineral phosphates have been worked for phosphorus in a similar way. 
 
 The liberated phosphorus e.scapes in vapor form, and is collected and conden.sed under 
 water in receivers of peculiar con.struction. To remove mechanically admixed impurities 
 
T36 
 
 INORGANIC CHF-MISTRY. 
 
 the phosphorus is aji^ain distilled from retorts, melted under water, pressed through leather 
 or chamois, or through bone ash, and finally moulded into sticks. 
 
 Phosphorus may be obtained by the method of Rossel on a small scale, if 2.5 parts of 
 aluminium powder, 6 parts of sodium metaphosphate, and 2 parts of silica are carefully 
 heated in a glass tube in a current of hydrogen : 
 
 bNaPOj -f loAl -f 3 Si 02 = sNa^SiOj -f SAl^Oj -f 6P. 
 
 (Compare Chemiker-Zeitung, 1898, 237.) 
 
 The almost colorless })hos})horus obtained by distillation is a waxy, 
 transparent substance, of specific gravity 1.83 at 10° C. At ordi- 
 nary temperatures it is soft and tough; at o'^ it becomes brittle. It fuses 
 under water at 44.4° and boils at 287°. By the action of sunlight it 
 becomes yellow (hence called yellow phosphorus), and is coated with a 
 non-transparent, reddish-white layer. Phosphorus is insoluble in water, 
 slightly soluble in alcohol and ether, and very readily soluble in carbon 
 bisulphide. It crystallizes from the latter solution in forms of the isometric 
 system (rhombic dodecahedra). When exposed to moist air, it oxidizes to 
 phosphorous acid (HgPOg) and phosjihoric acid (H^POJ ; the white vajiors 
 which arise contain ammonium nitrite (NH^NOJ, ozone, and hydrogen 
 peroxide. The odor of phosphorus resembles that of ozone. It phosphor- 
 esces in the dark on exposure to air. It does this also in other gases, but only 
 in such as contain oxygen. It appears the phos])horescence is influenced by 
 the formation and combustion of the self-inflammable phosphine, as all 
 substances which destroy the latter prevent and put an end to the former. 
 It is noteworthy that in pure oxygen the oxidation of phosphorus begins 
 at 27°. If the oxygen be diluted by removal over an air pump or by the 
 addition of neutral gases, so that its quantity is not more than 40 per cent., 
 the absorption will be very energetic at 20°, but cease entirely at 7°. 
 
 The. 7 'ed phosphorus — discovered by v. Schrotter in 1845 — possesses 
 properties entirely different from those of the ordinary variety. It is a 
 reddish-brown powder, of specific gravity 2.19 ; is not amorphous as for- 
 merly believed but, according to the investigations of Retgers, it is 
 hexagonal (Z. f. anorg. Chem. v (1894), 211). It is insoluble in carbon 
 bisulphide, non-phosphorescent, does not alter in the air, and is, indeed, 
 very stable. While ordinary phosphorus is very poisonous, the red variety 
 is perfectly harmless. It does not fuse at a red heat, even when subjected 
 to strong pressure, and vaporizes very slowly (even at 100°), without 
 changing to the yellow variety; but when it is heated rapidly (above 
 260°) the vapors change to those of yellow phosphorus. 
 
 To prepare the red variety, yellow phosphorus is heated for some min- 
 utes to 300°, in closed, air-tight iron vessels; there is a partial conver- 
 sion at 250°. The resulting mass is then treated with carbon bisulphide 
 or sodium hydroxide, to withdraw the unaltered, ordinary phosphorus. 
 If iodine be added to the ordinary phosphorus, the change will occur 
 below 200°. 
 
 A third modification — f}iefallic phosphorus — is formed if the amorphous variety be 
 heatcfl in a glass tube, free from air, to 530°. Microscopic needles then sublime into the 
 upper, less healed, portion of the tube. It is more easily obtained if phos|)horus is 
 heated with lead, in a closed tube, to a red lieat. The molten metal dissolves the phos- 
 
PHOSPHORUS. 
 
 137 
 
 phorus and, on cooling, the latter separates in black, metallic, shining crystals (llittorf ). 
 Metallic phosphorus possesses the specific gravity 2.34, vaporizes with difficulty, and is 
 less active than the red variety. Perhaps it consists of a better crystallized form of the 
 red variety. 
 
 Two green lines characterize the spectrum of phosphorus. On con- 
 ducting hydrogen over a small piece of phosphorus, heated in a glass tube, 
 the escaping gas will burn with a bright green flame. When ordinary 
 phosphorus is distilled with water, some passes over with the steam and, 
 in the dark, phosphoresces. This procedure serves for the detection of 
 phosphorus in poisoning by this substance (Mitscherlich). 
 
 The density of phosphorus is referred to O2 = 32 and therefore its 
 molecular weight is 124. As the atomic weight of phosphorus is 31.0, 
 it follows that the molecule in the form of vapor consists of 4 atoms : 
 
 124.0. We saw that the sulphur molecule at 500° consists of 6 atoms 
 (Sg), and at 900° of 2 atoms (S2). Such a dissociation does not, however, 
 occur with phosphorus ; even at 1040° its vapor density remains unaltered, 
 although a partial dissociation does take place at a very intense heat. 
 
 When phosphorus is burned in oxygen or in air, it forms the pentoxide 
 (P2O5). The ordinary variety inflames at 40°, and also by gentle fric- 
 tion ; the red is not ignited below 260°. The first will burn with a 
 bright flame even under water. To this end heat pieces of phosphorus 
 in a flask with water, until they melt, and conduct a current of oxygen 
 through the water. Phosphorus combines very energetically with chlorine, . 
 bromine, and iodine at ordinary temperatures; on throwing a small piece 
 into a vessel containing dry chlorine gas it at once inflames. The red 
 only reacts with the halogens after applying heat. With most of the 
 metals phosphorus unites on warming, and precipitates some of them from 
 solutions of their salts. From a silver nitrate solution, it precipitates 
 silver and silver phosphide (AggP) ; this solution, therefore, answers as a 
 counterirritant in phosphorus burns. 
 
 The difference in deportment of the yellow and the red phosphorus is fully accounted 
 for by the circumstance, that when the red is produced from the yellow there follows a 
 considerable heat-disengagement : 
 
 P yellow = P red +19.2 Cal. 
 
 Hence, the red variety contains much less energy than the yellow. In its union with 
 other substances there will always be liberated 19.2 Cal. less, and the reaction conse- 
 f|uently will proceed more sluggishly and with less energy. It may be assumed, that in 
 the yellow phosphorus, owing to the very energetic motions of its atoms their mutual 
 affinity acts in a much reduced form ; the atoms are loosely combined with one another 
 and are therefore very reactive. In the pas.sage of the yellow into the red or black 
 variety the energy of the atom motions is converted into thermal energy, and the atoms 
 can attach themselves more firmly to one another than before, the molecule is more stable, 
 the reactivity is diminished. Formerly when heat was still regarded as a substance, such 
 relations as those above were well expres.sed by the terms calorons — bodies containing 
 but little heat, and calorids — bodies rich in heat. 
 
 12 
 
INORGANIC CHRMIS'IRY. 
 
 
 COMPOUNDS OF PHOSPHORUS WITH HYDROGEN. 
 
 Plla, RJI^, 
 
 The compounds of phosphorus with hydrogen can be jirejjared l)y the 
 acticni of nascent hydrogen upon ])liosp]u)rus, as, for exainjile, on gently 
 heating dilute suljihuric acid with zinc and ])hosphorus (coinj)arc ]>. 144). 
 I he usual course is to heat yellow jihosphorus with concentrated potassium 
 or sodium hydroxide, when si)ontaneously inflammable i)hosi)hine escai)es 
 and a salt of hypo})hosj)horous acid remains in solution. 
 
 The lil)erate(l jras mixed with air in a closed ve.ssel explodes violently ; hence, to make 
 it, proceed as follows : hill a small glass flask almost full of a(jueous pota.ssium hydroxide, 
 
 Fig. 55. 
 
 add a few pieces of phosphorus, and heat over a lamp (Fig. 55). When the liberation of 
 gas commences, and the air in the neck of the flask has been expelled, close the .same with 
 the cork of the delivery tube, the other end of which dips under warm water, to prevent 
 any obstruction arising in it from phosphorus that may be carried over and solidify by 
 cooling. Kach bubble rising from the liquid inflames in the air, and forms white cloud- 
 rings which ascend. 
 
 I he gas tints ])roduced consists of gaseous plios|)hine (PH.j) and hydro- 
 gen, with whicli is mixed a small (piantity of a licpiid substance 
 whose presence imparts the s|)ontaneous inflammability to the gas. On 
 ctmducting the latter through a cooled tube the P^H^ is condensed to a 
 lifluid, and the escaping gas no longer inflames sitontaneously. The 
 
COMPOUNDS OF PHOSPHORUS WITH HYDROOFN. 
 
 139 
 
 licpiid compound may be isolated in a similar manner if tlie gas is con- 
 ducted through alcohol or ether, which will absorb the cum})ound 
 
 Liquid Phosphine, P2H4> separated from the gas by cooling, is a colorless, strongly 
 refracting liquid of specific gravity 1.012, insoluble in water, and boiling at 57° [Gat- 
 termann and llausknecht, Ber. 23 (1890), 1174]. It inflames spontaneously in the air, 
 and burns with great brilliancy to phosphorus pentoxide and water. Its jjresence in com- 
 bustible gases, such as hydrogen, marsh gas, and gaseous phosj)hine, gives them their 
 spontaneous inflammability. In contact with .some compounds, like carbon and sulphur, 
 and by the action of sunlight and of strong hydrochloric acid, it decomposes into gaseous 
 and solid phosphine f 
 
 5P,H, = 6PH3-f lyi,. 
 
 Solid Phosphine, P4II2 (?) is a yellow powder, which inflames at 160° or by a blow. 
 It is produced in the decomposition of calcium phosphide by hydrochloric acid. 
 
 Gaseous Phosphine, PH3, may be formed, together with the liquid 
 and solid variety, in addition to the manner previously described, by 
 the action of water or hydrochloric acid upon the calcium phosphides, 
 Ca3P2 and Ca2P2 : 
 
 Ca3P2 -f 6 HC 1 = 3CaCl2 -j- 2PH3 ; 
 
 Ca2P2 + 4IICI = 2CaCl2 + P2H4 ; 
 
 5 P 2 ll 4 = 6 PH 3 -fP 4 ll 2 . 
 
 Further, by the ignition of phosphorous and hypophosphorous acids : 
 
 4H3PO3 = PH3 + 3H3PO, 
 
 Phosphorous Phosphoric 
 
 acid. acid. 
 
 and by decomposing phosphonium iodide with caustic potash (p. 140), 
 
 It is a colorless gas, with a disagreeable, garlic-like odor, and is 
 slightly soluble in alcohol. Its density is 34 03 referred to O2 = 32, 
 or 1.176 (air= i). It solidifies at — 133- 5° and boils at about — 85°. 
 When i)ure (free from P2HJ it ignites at 100°. Oxidizing agents con- 
 vert it again into the spontaneously inflammable variety, owing to the 
 production of PgH^. It is extremely poisonous. Phosphine is decom- 
 posed into phosiihorus and hydrogen when it is heated, or if it is exposed 
 to the action of the electric spark. When ignited in the air it burns 
 with a brightly luminous flame, disseminating at the same time a white 
 cloud of phosphorus pentoxide (P2O5) : 
 
 2PII3 + 4O2 = 3H2O -f P2O5. 
 
 When mixed with chlorine it explodes violently, with production of 
 phosphorus trichloride and hydrogen chloride: 
 
 PH3 4- 3CI2 - PCI3 + 3HCI. 
 
 Like ammonia, phosphine jiossesses faint alkaline properties, and com- 
 bines with hydrogen iodide and bromide to yield compounds similar to 
 ammonium chloride : 
 
 PII3 -f HI == PHJ. 
 
 It combines with hydrogen chloride at from — 30*^ to — 35°, or, at ordi- 
 
140 
 
 INORGANIC CHEMISTRY. 
 
 nary tem])eratiires, under a jiressure of 20 atmosplicrcs. 'Die groiiji PII^, 
 figuring in the role of a metal in tliese compounds, is analogous to 
 ammonium (p. 129), and is termed pJiosphoniuin. 
 
 Phosphonium Iodide, PHJ. Jt is best prepared by the decomposi- 
 tion of phosphorus di-iodide (ly^, p. 142), by a slight (quantity of water, or 
 by adding yellow jihosphonis (10 jiarts), and, after some hours, iodine (2 
 jiarts), to a saturated solution of hydriodic acid (22 jiarts). The bhpiid 
 becomes a solid mass, consisting of ])hosphonium iodide and phosjihoroiis 
 acid. Phosphonium iodide sublimes in colorless, shining, cube-like 
 rhombohedra and by this means may be obtained jiiire. It fumes in the 
 air, and, with water, decomposes into phosphine and hydrogen iodide. 
 When decomposed by potassium hydroxide it yields pure hydrogen 
 phosjihide, which is not spontaneously inflammable : 
 
 PI 1,1 + icon =3 KI + PH3 4- II/). 
 
 Phosphine is a feebly exothermic compound : 
 
 P yellow + 311 = PII3 + 1 1.6 Cal. 
 
 MOLECULAR FORMULA OF PHOSPHINE. ATOMIC WEIGHT OF 
 PHOSPHORUS. 
 
 The analysis of phosphine shows that it consists of i.oi parts of hydrogen and 10.33 
 parts of phosphorus. Were its molecular formula PH, the atomic weight of phosphorus 
 would be 10.33. The great analogy of phosphine to ammonia, and that of all the 
 phosphorus compounds to those of nitrogen, argues, however, for the formula PII3. 
 The atomic weight of phosphorus, therefore, is 31.00 (= 3 X ^0.33), and the molecu- 
 lar weight of the phosphine is 34.03 : 
 
 113= 3.03 
 P =31.00 
 
 PH3 = 34.03 
 
 This view is confirmed by the density. Direct experiment confirms this. Further, 
 from the formula PH3 it follows that 3 volumes of hydrogen are present in 2 volumes of 
 the gas : 
 
 2PH3 contain 3H2, 
 
 2 vols. 3 vols. 
 
 or in I volume there are volumes of hydrogen. On decomposing the gas in a 
 eudiometer, by means of electric sparks, it will be found that the volume increases one-J 
 half ; the gas consists, then, of pure hydrogen, while phosphorus .separates in a solid 
 condition. As the phosphorus molecule in the gaseous condition is composed of four 
 atoms (p. 137), the jdiosphorus (62.00 parts) sej)arated from 2 volumes of phosphine 
 will fill X volume when in the form of vapor; hence in 2 volumes of phosphine there 
 are present 3 volumes of hydrogen and X volume of phosphorus vapor. 
 
 Or, written molecularly : 
 
 4Pll3 = P4 -f 611.2. 
 
 1 vol, I vol. 6 vols. 
 
COMPOUNDS OF PHOSPHORUS WITH THE HALOGENS. 
 
 I4I 
 
 COMPOUNDS OF PHOSPHORUS WITH THE 
 HALOGENS. 
 
 Phosphorus combines directly with the halogens to yield compounds 
 of the types PX3 and PX^, in which X indicates an halogen atom. 
 
 Phosphorus Trichloride — Phosphorous Chloride — PCI3. Conduct 
 dry chlorine gas over phosphorus gently heated in the retort C (Fig. 49, 
 p. in). The phosphorus ignites in the stream of gas, and distils over as 
 trichloride, which is collected in the receiver Z>, and condensed. The 
 product is purified by a second distillation. It is a colorless liquid, con- 
 gealing at — 112°, boiling at 76°, and has a sharp, penetrating odor. 
 Its specific gravity equals 1.613 at 0°. It fumes strongly in the air, and 
 is decomposed by moisture into phosphorous and hydrochloric acids : 
 
 PCI3 + 3H2O = H3PO3 + 3HCI. 
 
 The vapor density of the trichloride is 137.35, corresponding to the 
 molecular formula PCI3. 
 
 Phosphorus Pentachloride — Phosphoric Chloride — PCI5. This is 
 produced by the action of an excess of chlorine upon the liquid trichlo- 
 ride. It is a solid, crystalline, yellowish-white compound. It fumes 
 strongly in the air and sublimes without melting when heated. It at the 
 same time sustains a partial decomposition into trichloride and chlorine. 
 Under increased pressure it melts at 148°. 
 
 In an atmosphere of phosphorus trichloride the vapor density of the pentachloride 
 has been found to be 208.25, corresponding to the molecular formula PCI5 = 208.25. 
 At increased temperatures the vapor density steadily diminishes, and a gradual de- 
 composition occurs — dissociation of the molecules of the pentachloride (PCI5) into 
 the molecules of the trichloride (PCI3), and chlorine (Cb). The decomposition temper- 
 ature, /.<?., the temperature at which the decomposition is half finished, lies at about 
 200° C. The dissociation is complete at 300°, and equals 104, the vapor density; i. e., 
 the vapor then fills a volume twice as large as at a lower temperature. The breaking 
 down of the pentachloride into the trichloride and chlorine, whereby 2 n molecules result 
 from n molecules, explains this : 
 
 PCb = PCI3 4- Cb. 
 
 I vol. I vol. I vol. 
 
 That such a decomposition of the penta- into trichloride and chlorine does really occur, 
 is y^roved, among other things, by the originally colorless vapor gradually assuming the 
 yellow color of chlorine as the temperature rises. The decomposition j:)roducts — the 
 trichloride and chlorine — may be separated from each other by diffusion. 
 
 Phosphorus pentachloride acts very energetically with water ; it yields, 
 depending upon the quantities of the reacting substances, phosphorus 
 oxychloride, POCI3, metaphosphoric acid, HPO3, or orthophosphoric 
 acid, H3pO^, and hydrochloric acid : 
 
 PC'b + 11,0 = POCI3 -f 2IICI 
 PCb + 3H,0 = HPO, + 5IICI 
 PCb 4 - 41130 = I^PO^ 4- 5HCI. 
 
 (See Phosphoric Acid.) 
 
142 
 
 INORGANIC CIIKMISTUY. 
 
 "I'he ])ioiiiiiie and iodine })h()sphonis compounds are jierfeclly analogous 
 to the chlorine derivatives. They are obtained by uniting ihe constitu- 
 ents in the ])roi)ortions by weight expressed by their formulas. As the 
 union is exceedingly energetic, it is best to ])roceed as follows: Dissolve 
 the phosphorus in carbon bisulj)hide, gradually add the calculated amount 
 of bromine or iodine, and then distil off the volatile solvent. 
 
 Phosphorus Tribromide, PHr.,, is a colorless liquid, boiling at 175°, and having a 
 specific gravity of 2.9 at 0°. 'I'lie pentabromide, PP>r^, formeil by the gradual acUlition 
 of 2Br to PBr-j, is a yellow, crystalline substance, which melts when heated, and breaks 
 down into the tribromide and bromine. Water decomposes both compounds, as it does 
 the corresponding chlorides. 
 
 Phosphorus Chlorbromide, PCbIjrj, is ])roduced by the union of PCI, with Br^ in 
 the cold. It is a yellowish-red mass, which decomjKises at 35° into PCI., and Br.^. 
 
 Phosphorus Tri-iodide, PI,, forms red crystals, melting at 55° and distils, with par- 
 tial decom|)osition, at a higher temperature. 'I'he so-called phosphorus iodide, P.^b 
 (corresiK)nding to P.^H,), crystallizes in beautiful orange-red needles or prisms, ami 
 fuses at 110°. Its vapor density at 265° and 90.7 mm. })ressure ecpials 569, correspond- 
 ing to the molecular weight P^b- A little water decomposes it into phosjihorous acid, 
 phosi)hine, and hydriodic acid. The last two bodies then form phosphoniiini iodide, 
 
 PIbl (p. 140). 
 
 The recently di.scovered Phosphorus Pentafluoride, PFl^, is interesting (Thorpe). 
 It results upon heating phosphorus trichloride or pentachloride with arsenic trifluoride, 
 AsFl, : 
 
 3PCb + 5ASFI3 = 3PFb 4 SAsCl,. 
 
 It is a colorless gas which fumes in moist air and is decomposed by water into phos- 
 phoric acid and hydrogen fluoride. Its density corresponds to the molecular formula 
 PFl^ =: 125.9. It may be liquefied at 16° under a pressure of 46 atmospheres, and 
 solidifies when the pressure is removed. 
 
 It is rather remarkable that although phosphorus pentiodide could not be obtained, 
 the stability of the compounds, PBr,, PCl^, PFl-, gradually increases with the diminution 
 of the atomic weight of the combined halogens. Phosphorus pentafluoride can be gasi- 
 fied without decomposition. 
 
 Thermo-chemical Deportfueiit. — While the halogen derivatives of nitrogen (like tho.se 
 of oxygen) are strongly endothermic, are produced wdth the absorption of much heat, and 
 are, in consequence, readily exploded (p. 134), those of phosphorus are exothermic. The 
 heat disengaged in the union of yellow phosphorus and chlorine (p. 137) corresponds to 
 the following symbols : 
 
 (P,Cl 3 ) = 75.5; (P,Cb) = io5. 
 
 In this we observe a transition to the halogen derivatives of the metals, all of whicli 
 are exothermic. In accordance with this we find that the heat of formation of the bro- 
 mides and iodides diniini.shes in regular succession : 
 
 (DCb) = 75-5 ; (DB‘3) = 45 ; (PJ 3 ) = io-9- 
 
 'riie great reactivity of all the.se derivatives with water is fully ex]>lained by the large 
 amount of heat .set free at the .same time — the decomposition of phosphorus tri- 
 chloride with .so much water that upon diluting the resulting .solution, no further heat 
 evolution takes place : 
 
 PCd, -f A(i — IbPOsAq 4- 3lIClAq . . . 4-65 Cal. 
 
ARSENIC. 
 
 143 
 
 3. ARSENIC. 
 
 Atom : As = 75. Molecule : As^ = 300, 
 
 Arsenic is a perfect analogue of phosphorus, but possesses a somewhat 
 metallic character. In its free state it is similar to metals. 
 
 Arsenic is found free in nature, although it occurs more frequently in 
 combination with sulphur (realgar, As.^S.^, orpiment, As^Sg), with oxygen 
 (arsenolite, As.^03), and with metals (mispickel, FeAsS, cobaltite, CoAsS). 
 To prepare it, heat mispickel with iron, and free arsenic will sublime, 
 while iron sulphide remains. Or, in the customary way of isolating metals 
 from their oxides, heat the trioxide (arsenolite) with charcoal : 
 
 2AS2O3 6 C = As^ -[- 6CO. 
 
 Arsenic appears in two modifications. Crystallized (hexagonal) arsenic 
 is obtained by the sublimation of ordinary arsenic. It forms a gray- 
 white, more or less metallic, crystalline mass, but may be changed into 
 acute rhombic octahedra. Its specific gravity equals 5.73. It is brittle, 
 and may be pulverized without difficulty. 
 
 The amorphous is microcrystalline (according to Retgers) and 
 
 probably isometric. It is formed along with the first variety when arsenic 
 is sublimed in a glass tube in a current of hydrogen (Bettendorff ) and 
 also upon heating arsine. It is black, with little luster, and possesses the 
 specific gravity 4.71. When heated to 360° it sets heat free and reverts 
 to the crystalline variety. 
 
 Away from air contact, and at the ordinary pressure, arsenic vapor- 
 izes at a dark-red heat (about 450°) without previously melting ; it will, 
 however, melt if heated under great pressure in a sealed tube. Its vapor 
 possesses a lemon-yellow color. The vapor density corresponds to the 
 molecular weight. As its atomic weight equals 75, it follows that the 
 molecule in the gaseous state consists, like that of phosphorus, of four 
 atoms. At a white heat (about 1700°) the density falls to one-half, 
 which is due to the fact that then the arsenic vapor consists largely of 
 diatomic molecules (Asj = 150). 
 
 Arsenic does not change in dry air. When heated in the air it inflames 
 at 180° and burns with a blue-colored flame, disseminating the garlic-like 
 odor of arsenic trioxide (As^Og). It combines directly with most ele- 
 ments. Powdered arsenic will inflame when thrown into chlorine gas. 
 It yields arsenides with the metals. 
 
 It is remarkable that arsenic, belonging to the nitrogen group and generally forming 
 compounds which in constitution are quite different from those of sulphur, should be 
 analogous to the latter in its metallic combinations. Thus the sulphides and arsenides 
 have similar formulas, are isomorphous, and in them sulphur and arsenic can mutually 
 replace each other in atomic ratios, e. g. : 
 
 FeS^, FeAs.^ and Fe(SAs). 
 
144 
 
 INORGANIC CUKMISTRY. 
 
 COMPOUNDS OF ARSENIC WITH HYDROGEN. 
 
 Arsine, AsH., = 78.03. Like nitrogen and i)h()sj)liorus, arsenic fiir- 
 nishes a gaseous compound containing three atoms of hydrogen. It is ol)- 
 tained i)ure by the action of dilute sulphuric acid or hydrochloric acid 
 upon an alloy of zinc and arsenic: 
 
 As 2 Zn 3 -f 6I1C1 3ZnC1.3 -f- 2 ASII 3 . 
 
 It also results in the action of nascent hydrogen (zinc and suliihuric acid), 
 upon many arsenic compounds, as, c. g.^ the trioxide : 
 
 AS3O3 ^-6U, = As,U,-\ 3II./). 
 
 Arsine, discovered by Scheele in 1755, is a colorless gas, of strong, 
 garlicky odor, and extremely poisonous action ; it may be condensed to 
 
 a liquid and even to a solid by cold. It melts at — ii3-5° and boils at 
 — 55°. Its density equals 78.03 (O2 = 32) or 2. 69 (air = i). It burns 
 with a bluish-white flame when ignited, and evolves white fumes of 
 arsenic trioxide : 
 
 2ASII3 + 3O2 = AS2O3 -f 3H.p. 
 
 It is decomj)osed at a dull-red heat or by the electric spark into 
 arsenic and hydrogen. On conducting the gas through a heated tube the 
 arsenic de|)osits behind the heated part as a metallic coating (arse/iic 
 mirror). On holding a cold object, e. g., a i)iece of porcelain, in the 
 ll.'ime of the gas, the arsenic forms a black deposit (arsenic sj)ots). In its 
 rhemical behavior arsine is very similar to phosi)hine; its basic proper- 
 ties are very slight, and it does not furnish any derivatives with the 
 halogens. 
 
COMPOUNDS OF ARSENIC WITH THE HALOGENS. 
 
 145 
 
 According to analysis, arsine consists of 1. 01 parts by weight of hydrogen and 25 parts 
 of arsenic. If, because of its analogy to phosphine, we ascribe it the formula AsH^, 
 then the atomic weight of arsenic would be 75 (= 3 X 25) and the molecular weight of 
 AsHg would equal 78.03. A determination of the density confirms this. The formula, 
 too, shows that 3 volumes of hydrogen are present in 2 volumes of ASH3 : 
 
 2ASH3 contain 3H2. 
 
 2 vols. 3 vols. 
 
 We can satisfy ourselves of this by decomposing the gas by electricity in a eudiometer 
 (see p. 120). 
 
 Marsh’s Method for the Detection of Arsenic . — The method described 
 for the preparation of arsine and the ease with which it is recognized 
 make it possible to detect arsenic in its compounds with great certainty. 
 This is a very important task because of the ready accessibility and fre- 
 quent use of poisonous arsenic derivatives. Hydrogen is generated in a 
 flask (Fig. 56, a) by the interaction of dilute sulphuric acid and zinc, 
 when the material presumed to contain arsenic is introduced through the 
 funnel tube b. The gas evolved, a mixture of hydrogen and arsine, is 
 then conducted through a calcium chloride drying tube (c) and escapes 
 through the tube of hard glass contracted at several points (d). Upon 
 igniting the escaping hydrogen (after all the air has been previously expelled 
 from the vessel, as otherwise oxyhydrogen gas will be present) it will 
 burn with a bluish-white flame, if arsenic be present, and at the same time 
 disseminate a white vapor. The dark arsenic spots are obtained by hold- 
 ing a cold porcelain dish in the flame. If the tubed be heated (as shown 
 in Fig. 56), an arsenic mirror will be formed upon the adjacent contrac- 
 tion. The slightest traces of arsenic may be detected by this method. 
 
 Besides the ordinary arsine, AsHg, we might expect the existence of As-^H^ and 
 AS4H2, corresponding to the liquid and solid phosphines (P2fl4 and The first is 
 
 not known ; its derivatives exist, and contain hydrocarbon groups instead of hydrogen. 
 An example of this class is cacodyl, As2(CH3)^ = (CH3)2As-As(CH3)2. Nitrogen 
 affords similar compounds — fCH3)2N-NH2 and (CH3)NH-NH2, derived from diamide 
 or hydrazine, N2H^ = H2N-NH2 (p. 131). 
 
 The solid arsine, As^H2, is obtained by the action of nascent hydrogen upon arsenic 
 compounds in the presence of nitric acid. It forms a reddish-browm powder, which 
 decomposes when heated. Retgers considers that the brown color of the arsenic spots 
 and mirror is due to solid arsine, soluble in caustic potash. 
 
 COMPOUNDS OF ARSENIC WITH THE HALOGENS. 
 
 These are perfectly analogous to the corresponding phosphorus com- 
 pounds, and are the result of the direct union of their constituents. The 
 fluoride, in union with potassium fluoride, is the only known representa- 
 tive of the compounds corresponding to the formula AsX^ (see p. 141). 
 The metallic character of arsenic is shown by the fact that arsenic chloride, 
 like other metallic chlorides, may be obtained by the action of hydro- 
 chloric acid upon the oxide : 
 
 AS2O3 -f 6 HC 1 = 2ASCI3 -f 3H2O. 
 
 3 
 
146 
 
 INORGANIC CHEMISTRY. 
 
 Arstnic chloride is evolved when a solution of the trioxide is boiled 
 with concentrated hydrochloric acid. 
 
 Arsenic Trichloride, AsCl,, is a colorless, oily licjuid, fuming in 
 the air, and having a S|)ecific gravity of 2.2. Jt solidifies at lower tem- 
 peratures, melts at — 18° and boils at 130°. Its vapor density corresponds 
 to the molecular formula AsCl, = 181.35. dissolves in a small quan- 
 tity of water without change, while much water converts it into the oxide 
 and hydrochloric acid 
 
 2ASCI3 + 3II2O = As/J, -f 6 IIC 1 . 
 
 Arsenic Tribromide, AsBr,, is a white crystalline mass, melting at 
 20°, and boilingat 220°. The Tri-iodide, Asl^, forms red crystals, which 
 melt at 146”. The Trifluoride, ASFI3, is a licpiid, fuming strongly in the 
 air; it boils at 63°. It results in the distillation of the trichloride or tri- 
 oxide with calcium fluoride and sulphuric acid. 
 
 4. ANTIMONY. 
 
 Atom : Sb = 120. 
 
 The metallic character exhibited by arsenic becomes more distinct with 
 antimony, which at the same time retains its complete analogy to the 
 metalloidal elements, arsenic and phosphorus. Antimony is a perfect 
 metal so far as its physical properties are concerned. 
 
 It (stibium) occurs in nature chiefly in union with sulphur, as stibnite, 
 ^^>283 (Japan, Hungary), and with sulphur and metals in many ores. It 
 is almost always accompanied by arsenic. To prepare antimony, stibnite 
 is roasted in a furnace, i. e., heated with air access, whereby the sulphur 
 burns, and antimony trioxide remains: 
 
 28^83 -f 9O2 = 28b203 -f- 68O2. 
 
 The residual oxide is ignited with carbon, which reduces it to metal 
 (general procedure for the separation of metals). Antimony may also 
 be obtained by heating its sulphide with iron, which combines with the 
 sulphur : 
 
 SbaSa + 3re = 28b + 3FeS. 
 
 The resulting commercial crude antimony is further purified in the 
 laboratory by fusing it with niter, whereby the admixed arsenic, sulphur, 
 and lead are removed. Chemically pure antimony is obtained by reduc- 
 ing the pure oxide. 
 
 It is a silver-white and very brilliant metal, of leafy crystalline struc- 
 ture ; specific gravity 6.71. Like arsenic it crystallizes in rhombohedra, 
 is very brittle, and may be easily broken. It fuses at 430^^, and vaporizes 
 between 1500 to 1700°. The density of its vapor shows that the anti- 
 mony molecule, unlike that of phosphorus, consists not of four atoms, but 
 rather, like arsenic at the same temperature, of two atoms [Riltz, Z. f. 
 ]fliys. Chem. ig (1896), 385]. It is not altered in the air at ordinary 
 temperatures, but when heated it burns with a blue flame, yielding white 
 
COMPOUNDS OF ANTIMONY WITH THE HALOGENS. 
 
 147 
 
 vapors of antimony oxide, SbjO,. Like phosphorus and arsenic it com- 
 bines directly with the halogens; powdered antimony inflames in chlo- 
 rine gas. It is insoluble in hydrochloric acid ; nitric acid oxidizes it 
 (depending upon its strength and the temperature) to antimony oxide or 
 antimonic acid. 
 
 Hydrogen Antimonide — SHbine — SbHg, is produced like arsine, 
 and is very similar to the latter. It has thus far only been obtained 
 mixed with hydrogen. It is a colorless gas of peculiar odor, and when 
 ignited burns with a greenish-white flame, disseminating white vapors of 
 antimony oxide, SbgOg. A red heat decomposes it into antimony and 
 hydrogen. In Marsh’s apparatus (Fig. 56, p. 144) it forms an antimony 
 mirror and spots. The mirror is distinguished from that of arsenic by 
 its black color, lack of luster, its insolubility in a solution of sodium 
 hypochlorite (NaClO), and by its slight volatility in a current of 
 hydrogen. 
 
 Arsine is also liberated from alkaline solutions in which hydrogen is 
 being generated, e. g., caustic potash and zinc. This is not the case 
 with stibine. 
 
 COMPOUNDS OF ANTIMONY WITH THE 
 HALOGENS. 
 
 Antimonous Chloride — Antimony Trichloride — SbClg, results from 
 the action of chlorine upon the metal or its sulphide (SbgSg) ; better by 
 the solution of the oxide or sulphide in strong hydrochloric acid : 
 
 SbgSj -f 6HC1 = 2SbCl3 -f 3H2S. 
 
 This solution is evaporated to dryness and the residue distilled. 
 
 It is a colorless, crystalline, soft mass {^Butyruni antimonii), melting at 
 73° and boiling at 223°. Its vapor density corresponds to the molecular 
 formula, SbClg — 226.35. the air it attracts water and deliquesces. 
 It dissolves unchanged in water acidified with hydrochloric acid. Much 
 water decomposes it; the solution becomes turbid and a white powder — 
 powder of Algaroth (so named in honor of an Italian physician, Victor 
 Algarotus, who used it as a medicine) — separates : 
 
 SbClg + H2O = SbOCl 4- 2HCI. 
 
 The composition of this powder varies with the conditions under which 
 it is formed, but generally corresponds to the formula 2(SbOCl) . SbgOg. 
 Pure Antimony Oxychloride, SbOCl, obtained by heating antimony 
 trichloride with alcohol, occurs in colorless crystals and is further decom- 
 posed by water into basic oxychlorides. 
 
 While the metallic chlorides are not decomposed by water at ordinary 
 temperatures, the ready decomposition of the halogen derivatives of 
 antimony shows that this element still possesses a partial metalloidal 
 character. 
 
148 
 
 INORGANIC CHEMISTRY. 
 
 Antimonic Chloride — Antimoiiy Pentachloride — SbCl^, results from 
 the action of an excess of chlorine ui)on antimony or the trichloride. It 
 is a yellowish licjuid which fumes in the air, becomes crystalline when 
 cold and melts at — 6°. Heat partly decomjioses it, like jihosphorus 
 pentachloride, into antimonous chloride and chlorine : 
 
 SbCl^ = SbCf, 4- Clj. 
 
 I vol. I vol. I vol. 
 
 It may be gasified without decomposition at 218° and 58 mm. pressure. 
 Its density corresponds to the formula SbCl^. 
 
 Water converts it into pyroantimonic acid (H^SbjO^), and hydrochloric 
 acid. It combines with one molecule of water, forming the crystallizable 
 hydrate SbCl^ . 11 ./^, melting at about 90°, and with four molecules to a 
 hard, crystalline mass, SbCl5.4H20 [Anschutz and Evans, Ann. Chem. 
 239 (1887), 285]. The metallic nature of antimony shows itself in the 
 formation of these hydrates. Neither of the two chlorides of non-metallic 
 phosphorus forms a hydrate ; both are immediately decomposed by a 
 little water with the production of hydrochloric acid. 
 
 Antimony Tribromide, SbBrg, is a white, crystalline substance, melting at 93° and 
 distilling at 275°. The Tri-iodide, Sb^,, is a red compound, crystallizing in three dis- 
 tinct forms. It melts at 166° and boils at 400°. The Pentiodide, Sblj, is dark brown 
 in color and melts at about 78°. 
 
 Antimony Trifluoride, SbFlg, obtained by the solution of antimony oxide in hydro- 
 fluoric acid, crystallizes in colorless rhombic pyramids. It deliquesces on exposure to the 
 air. It is not decomposed by cold water. Its compound with ammonium sulphate is 
 used as a mordant in dyeing. 
 
 Antimony Pentafluoride, SbFlg, is a gummy mass. It forms well-crystallizing 
 double salts with the salts of organic bases. 
 
 We must also include Bismuth, Bi = 208.5, the group of nitro- 
 gen, phosphorus, arsenic, and antimony. Its halogen derivatives resem- 
 ble those of arsenic and antimony in many respects, e. g., BiClg, Bilj, 
 BiOCl. Its metallic character, however, considerably exceeds its metal- 
 loidal. Thus, it does not unite with hydrogen, and bismuth oxide (Bi203), 
 similar in constitution to the acid-forming arsenious oxide, AS2O3, 
 possesses only basic properties. We will, therefore, consider bismuth 
 and its derivatives with the metals. 
 
 TABULATION OF THE ELEMENTS OF THE NITROGEN GROUP. 
 
 The elements belonging here — nitrogen, phosphorus, arsenic, anti- 
 mony — present similar graded differences in their physical and chemical 
 |)r()j)crtics, just like the elements of the chlorine and oxygen groups, and 
 this gradation is intimately connected with the atomic weights. As 
 the latter increase the substance condenses, the fusibility and vola- 
 tility decrease, and the similarity to the real metals becomes more prom- 
 inent. 
 
CARBON GROUP. 
 
 149 
 
 
 N 
 
 P 
 
 As 
 
 Sb 
 
 Atomic weight, 
 
 14.04 
 
 31.0 
 
 75 
 
 120 
 
 Specitic gravity, 
 
 0.9 (solid) 
 
 1.8-2. 1 
 
 4 - 7 - 5-7 
 
 66.7 
 
 Melting point, 
 
 — 214° (60 
 
 44 " 
 
 red-white heat 
 
 
 mm. ) 
 
 
 
 
 Vapor density (O^ = 32), . . 
 
 28.08 
 
 124.0 
 
 150 
 
 240 
 
 These elements do not resemble one another in chemical properties as 
 the halogens resemble one another, or as sulphur, selenium and tellurium. 
 Arsenic and antimony alone show chemical kinship. What is exceed- 
 ingly striking is the difference between the remarkable activity of phos- 
 phorus in its yellow variety and the sluggish nitrogen with which red 
 phosphorus may be compared. This difference manifests itself in many 
 of their derivatives as has been observed with their halides (NCI3, NI3 — 
 PCI3, PI3), and will again be seen in their oxygen compounds. 
 
 Excepting bismuth, which is perfectly metallic in its nature, the ele- 
 ments of this group form gaseous compounds with three atoms of hydrogen. 
 Ammonia (NH3) possesses strongly basic properties, and combines with 
 all acids to yield ammonium salts; phosphine (PH3) combines at the 
 ordinary temperature only with hydrogen bromide and hydrogen iodide 
 to form salt-like compounds. Arsine and stibine no longer show basic 
 properties. Arsenic and antimony, as well as the two preceding elements, 
 combine with the hydrocarbons CH3 and C2H3) and form com- 
 
 pounds which are analogous in constitution and similar in character to 
 the hydrides. These compounds [As(CH3)3 and Sb(CH3)3] will be de- 
 scribed in Organic Chemistry ; they possess basic properties and yield salts 
 corresponding to the ammonium salts. Compounds of arsenic and anti- 
 mony corresponding to hydroxylamine and hydrazoic acid are not known. 
 
 The oxygen derivatives of these elements exhibit a gradation similar to 
 that of the hydrogen compounds. With increase of atomic weight, cor- 
 responding to the addition of metallic character, the oxides which form 
 strong acids in the lower series acquire a more basic nature in the higher 
 series. 
 
 CARBON GROUP. 
 
 The two non-metals, carbon and silicon, and the metals, tin and ger- 
 manium, comprise this group. They unite with four atoms of hydrogen 
 or with four atoms of the halogens. 
 
 1. CARBON. 
 
 Atom : C = 12.00. 
 
 Carbon occurs free in nature as the diamond and graphite. It consti- 
 tutes the most important ingredient of all the so-called organic substances 
 originating from the animal and vegetable kingdoms, and is especially 
 
INORGANIC CHEMISTRY. 
 
 150 
 
 contained in the fossilized jiroducts arising from the slow decomposition 
 of vegetable matter — in i)eat, in lignite, in bituminous coal, and in an- 
 thracite. In combination with hydrogen it forms the so-called mineral 
 oils — petroleum and asjihaltum. It occurs, further, as carbon dioxide 
 (CO^) in the air and in many waters; and in the form of carbonates 
 (marble, calcite, dolomite) comprises many minerals and entire rock 
 formations. 
 
 It is found in different allotropic modifications when free ; these may 
 be referred to the three })rincii)al varieties — diamond, grajihite and amor- 
 phous carbon. In all these forms it is a solid, even at the highest tem- 
 peratures; non-fusible and only volatile at about 3500° in the electric 
 arc. This deportment can only be explained by the supi^osition that its 
 molecules are composed of a large number of carbon atoms (see j). 106). 
 All the modifications of carbon are quite stable, but not very reactive. 
 When burned all yield carbon dioxide. 
 
 1. The diamond occurs in alluvial soils in certain districts (in India, Brazil, and South 
 Africa) ; less frequently in itacoluniite, micaceous schist, and xanthophyllite. It was ob- 
 served recently in meteoric iron from Canon Diablo. It has great luster, strong power 
 of refraction, and the greatest hardness of all substances. It crystallizes in forms of the 
 regular system, which are mostly rhombic dodecahedra, rarely octahedra. Ordinarily, it is 
 perfectly colorless and transparent ; sometimes, however, it is colored by impurities. Its 
 specific gravity equals 3.5. It does not soften unless exposed to the most intense 
 heat — between the poles of a powerful galvanic battery. It is then converted into a 
 graphitic mass. When heated in oxygen gas to 700-800° it burns to carbon dioxide 
 [see Ber. 23 (1890), 2409]. It is scarcely attacked at all when acted upon by a mixture 
 of nitric acid and potassium chlorate. 
 
 The diamond has been made artificially. Molten iron dissolves carbon, which sepa- 
 rates on cooling, mostly in the form of graphite. If iron, at high temperature, be 
 saturated with carbon and then be quickly cooled so that its inner portions are subjected 
 to great pressure because of the contraction of the exterior, upon sudden chilling, the 
 carbon then crystallizes in the form of the diamond (Moissan). Molten olivine (mag- 
 nesium silicate) also dissolves carbon, which again separates on cooling in the form of 
 little diamond crystals. 
 
 2. Graphite {ypd<pEiv^ to write) is characterized by its oxidation to graphitic acid when 
 it is heated with a mixture of potassium chlorate and nitric acid. Like amorphous 
 carbon, it is oxidized to mellitic acid by an alkaline solution of potassium permanganate, 
 or when it is made the positive electrode in the electrolysis of alkaline solutions. (See 
 Richter’s Organic Chemistry.) Native graphite is found in the oldest rock formations, 
 and of especially good quality at Altai, in Siberia. It occurs, too, in considerable quan- 
 tities at many places in the United States. It is occasionally found crystallized in six-sided 
 forms, but usually as an amorphous, grayish-black, glistening, soft mass, used in the 
 manufacture of lead pencils. The specific gravity is 2.25. It conducts heat and elec- 
 tricity well. When away from air contact it is not altered even at the highest tempera- 
 tures. It is the only variety of carbon stable at high temperatures ; when heated the 
 other varieties pass into it. It usually burns when heated in an atmosphere of oxygen, 
 but with more difficulty than the diamond, forming carbon dioxide, and leaving about 
 2-5 ])er cent, of ash. To purify the poorer and more impure kinds of graphite, they 
 are ])iilverized and heated with a mixture of potassium chlorate and sulphuric acid ; the 
 {product is washed with water, and the residue ignited (Brody’s graphite). 
 
 Luzi claims two varieties of graphite: graphiUte and the true graphite. The first 
 does not swell up when moistened with red, fuming nitric acid, and is then strongly 
 ignited (Bassauer and Siberian graphite). 'I'he true graidiites do swell up under similar 
 tr(“atment (gra|)hite from (feylon, New York, 'ficonderoga) ; see Ber. 26 ( 1893), 890. 
 
 Urapliile may be obtained artificially by fusing amoridious carbon with iron ; when the 
 latter cools, a j)ortion of the dissolved carbon separates in hexagonal shining leaflets. 
 
COMPOUNDS OF CARBON WITH HYDROGEN. 
 
 15I 
 
 3. Amorphous carbon (a, privative, and /jopcl)//, form), is produced by the carbonization 
 of organic (containing carbon) substances, anci is found in a fossilized state. Nitric acid 
 and potassium chlorate convert it in the cold into brown substances soluble in water, d'he 
 purest amorphous carbon is soot, which is obtained by the imperfect combustion of resins 
 and oils (like turpentine) rich in carbon. Gas carbon, called metallic carbon, deposits 
 in the manufacture of gas in the retorts, and is very hard, possessing metallic luster, and 
 conducting electricity well ; hence its use in galvanic batteries. Coke, resulting from the 
 ignition of bituminous coal, forms a sintered mass, conducting heat and electricity well. 
 Charcoal is very porous, and can absorb many gases and vapors ; i volume of it con- 
 denses 90 volumes of ammonia, 55 volumes of hydrogen sulphide, and 9 volumes of 
 oxygen. At 1 00°, and under the air-pump, the absorbed gases are again liberated. 
 Charcoal will also take up many odorous substances and decaying matter ; hence it is 
 employed as a disinfectant. Anhnal charcoal- is obtained by the carbonization of animal 
 matter (bones, blood, etc.), and possesses the power of removing many coloring sub- 
 stances from their solutions ; hence it serves in the laboratory and in commerce for the 
 decolorization of dark solutions. 
 
 All these varieties of carbon contain smaller or larger quantities of nitrogen, hydrogen, 
 and mineral substances, which remain as ash after combustion. Hydrochloric acid will 
 withdraw almost all the mineral constituents. 
 
 The fossil coal varieties, bituminous coal, lignite and turf, are the products of a peculiar, 
 slow decay of wood fiber, which gradually separates oxygen and hydrogen, and enriches 
 itself in carbon. Fossil coal contains 90 per cent, and lignite 70 per cent, of carbon. 
 The fossil coal richest in carbon, the last product of the alteration, is anthracite. This 
 has lost all its organic structure, and contains 96-98 per cent, of carbon. 
 
 COMPOUNDS OF CARBON WITH HYDROGEN. 
 
 With hydrogen, carbon forms an unlimited number of compounds, into 
 which all other elements, especially oxygen and nitrogen, can enter. 
 The derivatives of carbon have been termed organic conipoimds, because 
 they were formerly obtained exclusively from vegetable and animal 
 organisms, and the idea was entertained that they were produced by the 
 influence of forces other than those forming the mineral substances. At 
 present, most carbon derivatives are prepared artificially from the ele- 
 ments by simple synthetic methods; we are aware that they do not differ 
 essentially from mineral substances. Hence the description of the carbon 
 compounds must be arranged in the general system of chemical bodies. 
 This, however, is not readily executed without sacrificing the review of a 
 defunct system. The derivatives of carbon are so numerous, and possess 
 so many peculiarities, that it appears necessary, from a practical stand- 
 point, to treat them apart from the other compounds, in a separate portion 
 of chemistry, which we, pursuing the old custom, term organic chemis- 
 try. We then designate the chemistry of all other bodies as Inorganic 
 Chemistry. Only the simplest carbon compounds will be considered here. 
 
 It is only under the influence of the electric arc that the direct union 
 of carbon and hydrogen can be effected ; the ])roduct is acetylene 
 (CJI.J. All other hydrocarbons are obtained indirectly in various ways. 
 
152 
 
 INORGANIC CHEMISTRY. 
 
 Methane — ATarsh Gas — CII^. "J'his simplest liydrocarlion is lliat 
 cliciiiical compound which is richest in hydrogen ; it contains 25.18 jier 
 cent. It is lormed in the decay of organic matter under water (in 
 swamps, hence called marsh gas, and in coal mines), and escapes in 
 large quantities in many regions of the earth (thus at I’.aku, on the Cas- 
 pian Sea, in Pennsylvania and in Ohio, where the gases esca])ing from 
 petroleum wells contain as much as 90-97 per cent, of methane). It is 
 produced in the dry distillation of coal, wood, i)etroleum and other 
 organic substances; hence is a constituent of illuminating gas. It may 
 be obtained synthetically by conducting vapors of carbon bisulphide and 
 hydrogen sulphide over ignited copper filings; 
 
 CS2 + 2 lI.,S + 8 Cu 4Cu.,S + CH^. 
 
 For its preparation, heat a mixture of sodium acetate with sodium 
 hydroxide : 
 
 CJbOjNa + NaOH ^ CII, -f- Na,C03. 
 
 Its formation from aluminium carbide by the action of water is rather 
 interesting. Aluminium hydroxide is also produced (Moissan): 
 
 Al.Cg + 6II.p == 2A1/)3 + 3CIb. 
 
 Methane is a colorless, odorless gas, insoluble in water. Its density 
 corres})onds to the molecular formula CH^. Methane, under great pres- 
 sure and at low temperatures (lielow — 82°, p. 48), is condensed to a 
 colorless liquid, which boils at — 164° under the ordinary pressure. Its 
 specific gravity is 0.415 at the boiling point. When ignited it burns 
 with a faintly luminous flame. It affords a violently explosive mixture 
 (fire-damp of the miners) with two volumes of oxygen (or ten volumes of 
 air) : 
 
 CH^ + 2O2 ^ CO2 -f 2H2O. 
 
 I vol. 2 vols. I vol. 2 vols. 
 
 MOLECULAR FORMULA OF METHANE. ATOMIC WEIGHT OF CARBON. 
 
 The quantitative analysis of methane shows that for every i.oi parts of hydrogen in it 
 there are 3 parts of carbon. Were the formula CH (analogous to hydrochloric acid) then 
 the atomic weight of carbon would be 3. If it corresponded to the formula of water 
 
 (II20) 
 
 or ammonia (NII3) then the atomic weight of carbon would equal 6, or 9 : 
 
 II 
 
 = I.OI 2II = 2.02 
 
 311 = 
 
 3 - 03 
 
 4II = 4.04 
 
 c 
 
 = 3 C = 6 
 
 C = 
 
 9 
 
 C = 12 
 
 ClI 
 
 = 4.01 CII2 = 8.02 
 
 CII3 = 
 
 12.03 
 
 CIb = 16.04 
 
 All these and many other formulas agree with the analytical results. The gas density, 
 however, decides in favor of the formula Clb, because a liter of methane weighs under 
 normal conditions 0.71464 gram and this would make the molecular weight, referred to 
 (), - 32 (.see p. 79), = .6.0. 
 
 d hat the atomic weight of carbon is really 12, is proven by the fact that of all its 
 in nunicrahle derivatives, not one contains in its molecule less than 12 farts, by 7 veight 
 {^therefore in the mol-volume (p. 98) not less than 12 grams'), of this element. It follows, 
 with certainty, from the i)criodic system of elements. 
 
ETHANE ETHYLENE. I53 
 
 From the formula CII^ it follows that in i volume of methane there are 2 volumes of 
 hydrogen : 
 
 CH^ contains 2li.^ 
 
 I vol. 2 vols 
 
 This is proved indirectly by the combustion of methane with oxygen in a eudiometer 
 (see p. 137)- Four atoms of hydrogen yield two molecules of water ; one atom of carbon 
 yields one molecule of carbon dioxide, CO2 Hence the volume relation in the combus- 
 tion of methane in oxygen is expressed by the equation ; 
 
 CH, + 2O2 =• CO2 + 21120 . 
 
 I vol. 2 vols. 1 vol. 2 vols. 
 
 In 2 volumes of aqueous vapor there are 2 volumes of hydrogen ; hence in I volume 
 of Cll^, there are 2 volumes of II2. The result of the eudiometric analysis conhrms 
 these conclusions 
 
 Ethane or Ethyl Hydride, C2Hg, is formed when hydrogen m 
 statu nascendi acts upon ethyl bromide : 
 
 C2H5Br + H2 =■ + HBr, 
 
 or by the action of potassium or sodium upon methyl iodide (p. 159) : 
 
 2CH3I -f Na2 = C2II6 + 2NaI. 
 
 This is a colorless gas, insoluble in water, and when ignited it burns 
 with a feebly luminous flame. Liquid ethane boils at — 89.5° under 735 
 mm. pressure. 
 
 Besides methane (Cbb) and ethane (C2Hg) there exists a long series of hydrocarbons 
 of the general formula CnH2n-l-2 g’, Cgllg, C5IIJ2. etc.), in which each mem- 
 
 ber differs from the preceding and next following by iC and 2H (CH2). Bodies 
 belonging to such a series, greatly alike in their chemical behavior, are termed honiolognes 
 {ouoiog, similar; Aoyo^, nature of the affair). In addition to this series of satm-ated 
 hydrocarbons others exist, with less hydrogen, and by the addition of the latter, pass 
 into the saturated^ and may, therefore, be termed jmsaturated. The first uusattirated 
 .series has the formula Cnll^,,, the second, the fonnula CnH2n— 2, etc. The lowest member 
 of the series Cnll2n is ethylene (see Chemical Structure, p, 168). 
 
 Ethylene, C2H^, is formed in the destructive distillation of wood, 
 bituminous coal, and many carbon compounds, hence is contained in 
 illuminating gas. It is most easily obtained by the action of sulphuric 
 acid upon alcohol, whereby the acid withdraws the elements of water from 
 the latter • 
 
 C2lTgO — 1120 =r C 2 IL. 
 
 Alcohol. Ethylene. 
 
 It is a colorless gas, of weak, ethereal odor. Its critical temperature 
 is its critical pressure, 51 atmospheres. Under a pres- 
 
 sure of I atmosphere, licpiid ethylene boils at — 102°, and when in a 
 vacuum at ■ — 150°. It solidifies at — t6o°. As it does not solidify on 
 va])orizing, it is well adapted for the liciuefaction of other gases (p. 48). 
 It burns with a bright, luminous flame, decomposing first into marsh gas 
 and free carbon : 
 
 €211, = CII, + C. 
 
 The methane then burns and heats the i)articles of carbon in the flame 
 
154 
 
 INORGANIC CHEMISTRY. 
 
 to incandescence ; tliese are then consumed to carl)on dioxide (CO.^) (see 
 l>er. 26 ( I ^93), III, 9). 
 
 'riie unsaliirated compound, ethylene, unites directly with two atoms 
 of chlorine and of bromine: 
 
 Cdb b Cl, C,TI,C1,. 
 
 The resulting compounds, and C,l l^Ilr.^, are oily liciuids; hence 
 
 the name olefiant ^as for ethylene. 
 
 The first member of the second unsaturated series is acetylene, 
 
 Acetylene is juoduccd in the dry distillation of many carbon com- 
 pounds, and is present in coal gas, to wh.ch it imparls a ])eculiar ]>ene- 
 trating odor. It is also formed in the incom])lete combustion of coal gas — 
 e. when the fiame of a Bunsen burner strikes back. Berthelot syn- 
 thesized acetylene by causing the electric spark to strike across from car- 
 bon points in an atmosphere of hydrogen : 
 
 C, + II, = C,H,. 
 
 No other hydrocarbon has as yet been directly built up from its ele- 
 ments. 
 
 The production of acetylene by the action of water or dilute acids 
 upon metallic carbides is very interesting, e. g., from calcium carbide: 
 
 CaC, + 211,0 = CjII, + Ca(OH),. 
 
 Pure acetylene is a colorless gas with a peculiar odor. It can be readily 
 condensed. Its critical temperature is 37°; its critical pressure equals 
 68 atmospheres. At 0° it may be liquefied by a pressure of 21.5 atmos- 
 pheres. Liquid acetylene boils at — 83°. When it is poured it changes in 
 part to a snow-like, combustible mass. At 19.5°, i kilogram of liquid acet- 
 ylene will yield 896 liters of gas under the ordinary atmospheric pressure. 
 At medium temperatures it dissolves in an equal volume of water; there- 
 fore, it must be collected over strong salt solutions. It is very soluble in 
 acetone, i volume of the latter dissolving 25 volumes of acetylene at the 
 ordinary temperature and pressure. 
 
 Being an unsaturated hydrocarbon it combines with chlorine; with 
 two or four atoms of the same : 
 
 C2H2CI2, C 2 H 2 Ch. 
 
 It is chemically very much like hydrazoic acid (p. 132). It possesses acid 
 l)roperties but these are not so pronounced as in the case of hydrazoic 
 acid ; acetylene and many of its metallic derivatives are explosive (p. 31). 
 
 As a result of o])servations made by Willson, an American, when 
 attempting to ])repare the alkaline earth metals by the reduction of their 
 oxides with carbon in the electric furnace, calcium carbide has been made 
 since 1894 on a large scale, for the purpose of producing acetylene by the 
 method employed by Moissan. As acetylene is now so readily accessible 
 it plays an imj)ortant role as an illuminant. It burns, on issuing from a 
 small aperture under a definite i)ressure, with a blinding white flame, 
 almost free from soot. As it is readily condensed it is especially adapted 
 for the illumination of vessels, trains, etc. 
 
THE NATURE OF FLAME. 
 
 155 
 
 Mixtures of acetylene with from 1.25 to 20 volumes of air are explosive; 
 a most powerful explosion will occur when the j^roportion is i volume of 
 acetylene and 12 volumes of air. The acetylene prejjared from commer- 
 cial calcium carbide usually contains hydrogen sulphide and phosphide. 
 
 Compare: Fr. Liebetanz, Calciumkarbid und Acetylentechnik, 2. 
 Aufl., Leipzig, 1899; F. Dommer, Calciumkarbid und Acetylen ; 
 deutsch von Landgraf, Miinchen und Leipzig, 1898. 
 
 The three hydrocarbons considered above, methane (CHJ, ethylene 
 (C2HJ, and in slight amount acetylene (C2H2), constitute, together with 
 hydrogen and carbon monoxide (CO), ordinary ilhwiinating gas, which is 
 produced in the dry distillation of bituminous coal, lignite, or wood. 
 The illuminating power is influenced by its quantity of ethylene and acet- 
 ylene (and their homologues). 
 
 Late investigations indicate that all hydrocarbons are broken down by the heat of the 
 flame into acetylene and carbon, which then burn and become luminous (see Z. f. anorg. 
 Ch. 9 (1895), 233). 
 
 THE NATURE OF FLAME. 
 
 We are aware that every chemical union which occurs in a gaseous 
 medium, and is accompanied by the evolution of light is designated com- 
 bustion. Some bodies, like sulphur and phosphorus, yield a flame when 
 burned in the air or in other gas; such sub- 
 stances are converted into gases or vapors at the 
 temperature of combustion. Pure carbon burns 
 without a flame, becomes incandescent, because 
 it is non-volatile. The carbon compounds, 
 wood, bituminous coal, and tallow, are, indeed, 
 not volatile, but burn with a flame because 
 under the influence of heat they develop com- 
 bustible gases. Flame is, therefore, nothing 
 more than a combustible gas heated to incan- 
 descence. We know, too, that hydrogen burns 
 in oxygen and chlorine, conversely, oxygen and 
 chlorine will burn in hydrogen (p. 57), and 
 that illuminating gas burns in the air, therefore 
 air (its oxygen) burns in the former. This may 
 be demonstrated in the same manner as in the 
 case of chlorine and hydrogen. 
 
 The relative combustibility and the so-called return of 
 the flame may V)e very plainly illustrated by means of the 
 following contrivance : An ordinary lamp chimney (Fig. 
 
 57) is closed at its lower end with a cork, through which FlG. 57. 
 
 two tubes enter ; the narrow lube, a, somewhat con- 
 tracted at its end, is connected with a gas stop-cock ; the other tube, h (best a cork-borcr \ 
 is about 5 mrn. wide, and communicates with the air. The gas issuing from the tube a 
 is ignited, and the chimney is then drojiped over the not too large flame ; it contituies to 
 burn along (|uietly, as sufficient air enters through the wide tube b. Upon increasing the 
 supply of gas, the flame becomes larger, the globe fills with illuminating gas, while the 
 
156 
 
 INORGANIC CHEMISTRY. 
 
 air is displaced. The p;as flame is extin^uislied, and an air flame appears upon the 
 wider tube, h, as the enlerinjj; air continues to burn, in the atmosphere of illuminatinfr 
 f^as. 'I'he exce.ss of tlie latter e-scapinj^ from the upper ])ortion of the ^lobe may be 
 if^nited, and we then have a gas flame above, while within the globe we have an air llame. 
 On again lessening the gas How the air flame will distribute itself, extend to the exit of 
 the tube and then the gas flame will appear upon the latter, while the flame above the 
 globe is extinguished. In this manner, we may repeat the return ])roeess of flames at 
 will. That the air actually burns in the air flame may be plainly proved if we introduce 
 a small gas flame from c, through the wide metallic tube the little flame will con- 
 tinue to burn in the air flame, but will be extingui.shed if it be introduced higher up 
 into the atmosphere of illuminating gas. 
 
 We say ordinarily that only those bodies are combustible which burn 
 in an atmosphere of oxygen or in air. If we imagine, however, an at- 
 mos])here of hydrogen, or illuminating gas, then bodies rich in oxygen 
 
 Fig. 58. 
 
 must be combustible in these. In fact, nitrates, chlorates, etc., burn in 
 an atmosphere of illuminating gas with the production of an oxygen 
 flame. This may be demonstrated as follows: An Argand lamp chimney 
 (Fig. 58) is closed at its lower end with a cork, bearing a gas-conducting 
 tube, d'he gas which escapes through the opening of the metal cover, 
 a, is ignited. Then the substance (potassium or barium chlorate, etc.) 
 is introduced into the flame on an iron spoon provided with a long 
 handle, heated to the temperature of decomposition (disengagement ot 
 oxygen), and the sjioon then jilunged through the opening into the gas 
 atmosjihere. d'he substance burns with a brilliant light, and gives a 
 cliaracteristic flame reaction. 
 
THE DENSITY OF THE GAS OF FLAMES. 
 
 157 
 
 The brilliancy or luminosity of a flame is influenced by the nature of 
 the substances contained in it, also by their temperature and density. In- 
 candescent gases shine very faintly per se ; this is especially true when 
 they are diluted. Thus hydrogen, ammonia and methane burn with a pale 
 flame. Even sulphur burns in the air with a slightly luminous flame. If, on 
 the contrary, sulphur or phosphine be permitted to burn in pure oxygen, 
 or arsenic and antimony in chlorine gas, an intense display of light fol- 
 lows. This depends on the fact that the flame is not diluted by the nitro- 
 gen of the air, and therefore develops a higher temperature. That the 
 density of the gas of flames exercises a great influence upon the lumi- 
 nosity is proved by the fact that hydrogen, compressed into a smaller 
 space with oxygen, burns with intense light. 
 
 A slightly luminous flame may be rendered intense by introducing solid 
 particles into it. For example, if hydrogen be passed through liquid 
 chromium oxychloride (Cr02Cl2) it burns with a bright, lumi- 
 nous flame, because the chromium oxychloride is changed by 
 the high temperature, with the absorption of hydrogen, into 
 water, hydrochloric acid, and solid, non-volatile chromium 
 oxide, whose particles are heated to incandescence by the 
 hydrogen flame. The illuminating power of the various 
 hydrocarbons and carbon compounds is similarly explained. 
 
 Marsh gas, CH^, and ethane, C2Hg, give a pale flame, be- 
 cause they burn directly to aqueous vapor and carbon dioxide ; 
 ethylene, on the contrary, burns with a bright, luminous 
 flame, because, by the temperature of combustion, it decom- 
 poses first into methane and carbon, whose particles glow in 
 the flame (see p. 153). 
 
 Let us consider the flame of an ordinary stearin candle : On 
 approaching the wick with a flame the stearin melts, is drawn 
 up by the fibers and converted into gaseous hydrocarbons, 
 which ignite, and by their chemical union with the oxygen 
 of the air produce the flame. Three zones are distinguish- 
 able in this flame. The unaltered gases exist in the inner non-volatile 
 zone a (Fig, 59); they cannot burn because of lack of air. If the 
 lower end of a thin glass tube be inserted here the gases will rise in it, 
 and may be ignited at the upper end. There is a partial combustion of 
 the gases in the middle, luminous part, y, g; ethylene, C2H4, breaks 
 down here into methane, CH^, and carbon, C ; the first burns completely, 
 while the carbon is heated to a white heat, because there is not sufficient 
 oxygen ])resent for its combustion. The presence of carbon particles in 
 the luminous part may be easily proved by placing a cold glass rod or a 
 wire in it ; it will at once be covered with soot. 
 
 In the outer, very feebly luminous and almost invisible mantle, bj c, d, 
 of the flame, which is completely surrounded by air, occurs the perfect 
 combustion of all the carbon to carbon dioxide. 
 
 A perfectly identical structure is possessed by the ordinary illuminating 
 gas flame. By bringing as much air or oxygen into it as is necessary for 
 the complete combustion of all the carbon, none of the latter separates 
 (see p. 158), and there is produced a faintly luminous but very hot flame. 
 
 Fig. 59. 
 
INORGANIC CHEMISTRY. 
 
 158 
 
 Upon this ])rincii)le is based tlie construction of the Bunsen burner, the 
 flame of which is employed in laboratories for heating and ignition. 
 Fig. 60 represents a form of the same. The upi)er tube, c, is screwed 
 into the lower portion, and in the figure is separated merely for the sake 
 of exi)lanation. The gas enters through the narrow ojiening, a, from the 
 side gas tube, and mixes with air in the tube c, which enters through 
 the openings of the ring, b. In this way we obtain a flame which is but 
 faintly luminous, although affording an intense heat. On closing the 
 openings in b the air is cut off, and the gas burns at the uiiper end of the 
 lube c with a bright, strongly smoking flame. The non-luminous flame 
 contains an excess of oxygen, and hence oxidizes — oxidizing It 
 is employed to effect oxidation reactions. Its temperature is about i 200°. 
 'I'he luminous flame, on the other hand, is reducing in its action (tem- 
 ])erature 1000° C.), and is designated the 7 'educing flame, because the 
 glowing carbon in it abstracts oxygen from many substances. 
 
 The non-lurninosity of tlie Bunsen burner flame, clue to addition of air, depends on a 
 more complete combustion of the separated carbon or of the yet undecomposed hydro- 
 carbons. 
 
 Another variety of non-luminosity of hydrocarbon flames is induced by the ad- 
 mixture of inactive gases, like nitrogen and carbon dioxide. By this means the 
 
 flame is enlarged and the combustion, as in the luminous flame, takes place only 
 in the outer cone ; further, the temperature is lowered, and probably does not 
 acquire the decomposition temperature of methane and 
 ethylene. The simple extension of an illuminating flame 
 upon a plate, will render it non-luminous, because then 
 the air comes in contact with a larger flame surface. On 
 heating a gas made non-luminous by the admixture of 
 nitrogen, and then letting it burn, its flame becomes lumi- 
 nous because the increased temperature can induce the 
 decomposition of methane. 
 
 rh-. : 
 
 
 irisfai 
 
 Fig. 60. 
 
 Fig. 61. 
 
 In rendering flame non-luminous by carbon dioxide, we must also consider that the 
 same is converted, by the particles of carbon, into carbon monoxide : 
 
 CO2 -f C = 2CO. 
 
 Indeed, but a few per cent, of carbon dioxide in a gas flame suffices to considerably 
 diminish its luminosity : 
 
 (',H, I ro., _ C’ll, + 2C(), 
 
 1 vul. I vul. I vul. 2 vols. 
 while the presence of nitrogen is far less detrimental. 
 
COMPOUNDS OF CARBON WITH THE HALOGENS. 1 59 
 
 Every substance requires a definite temperature for its ignition — tem- 
 perature of igjiition. When a substance is once ignited it generally burns 
 further, because additional particles are raised to the temperature of 
 ignition by the heat of combustion. By rapid cooling (<?. g., by the 
 introduction of a piece of metal into a small flame) every flame may be 
 extinguished. By holding a wire gauze over the opening of a gas lamp, 
 from which gas issues, and igniting the same above the wire (Fig. 6i), 
 the latter, being a good conductor of heat, cools the flame so much that 
 it is incapable of igniting the gas below the gauze. Upon this phe- 
 nomenon depends the construction of Davy’s safety lamp, which is used 
 in coal mines to avoid ignition of the fire-damp. It is an ordinary oil 
 lamp surrounded and shut off from the air by a metallic wire gauze. On 
 bringing a lighted lamp of this sort into an explosive mixture, or into a 
 combustible gas {e. g., into a large jar, in which ether is present), the 
 gas penetrating into the lamp will burn, but the combustion will not ex- 
 tend to the external gases. 
 
 COMPOUNDS OF CARBON WITH THE HALOGENS. 
 
 Carbon combines directly with fluorine alone of the halogens, forming 
 gaseous carbon tetrafluoride, CFl^. Compounds with all the halogens 
 result, however, by the action of the latter upon the hydrocarbons. We 
 have seen that chlorine and bromine act upon water, ammonia, hydrogen 
 sulphide, etc., in such manner as to unite with the hydrogen to form 
 hydrogen chloride, etc., while the other element is either set free or is 
 also combined with the chlorine (p. 128). Chlorine acts similarly upon 
 the hydrocarbons; here hydrogen is displaced, atom after atom, by 
 chlorine, forming hydrogen chloride and chlorine derivatives : 
 
 CH, + CI2 = CH3CI -f HCl, 
 etc + 2C1, = ciccq + 2HC1, 
 
 CH^ -p 3CI2 =: CHCI3 -f 3HCI, etc. 
 
 Such a process is termed substitution^ and the products substitution prod- 
 ucts. These halogen derivatives are distinguished from the halides of 
 other non-metals by their comparatively greater stability toward water 
 (see pp. 148, 1 61). In this way we obtain from methane, CH^, the 
 products, CH3CI, CH2CI2, CHCI3 (chloroform) and finally CCl^, carbon 
 tetrachloride, a colorless, ethereal liquid, boiling at 76°. 
 
 The compound hexachlorethane, obtained by the action of chlorine upon ethane, 
 
 CjHg, is a crystalline mass, melting and boiling at 186°. On conducting its vapors through 
 a red-hot tube they decompose into ethylene tetrachloride and chlorine : 
 
 QCl, = C,Ch -p cp. 
 
 Ethylene tetrachloride^ CjCp, is a liquid boiling at 122°. 
 
 P'luorine, bromine, and iodine yield similar compounds ; they will be treated more 
 exhaustively in Organic Chemistry. 
 
i6o 
 
 INORGANIC CHRMISTRY. 
 
 Tlie heals of formation of (lie previously mentioned hydrocarbons (from amorphous 
 carbon and hydrogen) are deduced from tlie heat of conil)ustif)n, and ef|ual : 
 
 (C,II,) = 2I.7; (C„ig = 28.6; (C„nj --2.7; (C 2 , II,) -47-8. 
 
 The a])sorption of heat in the formation of acetylene, f',!!,, and ethylene, is 
 
 explained by the fact that the solid carbon molecules must first be gasi(ie<I and sejiarated 
 into atoms in order to unite with hydrogen. The energy reijuisite for this eipials very 
 nearly 39.7 Cal. for 12 parts by weight of amorjihous carbon (76.4 Cal. for 24 jiarts) ; 
 then the preceding heats of formation must be increased about'this amount, in order to 
 express the true heat of formation (from gaseous carbon atoms) of the hydrocarbons. 
 
 Beginning with acetylene, C,!!,, we discover that its conversion into ethylene, C,!!^, 
 and into ethane, C,IIg, ensues with heat disengagement : 
 
 (c,ii„h.,) = 45 -i; ((’2^4,11,) = 25.9 
 
 In accordance witli this acetylene readily unites with liydrogen (/w statu naseendi, or 
 by action of platinum sponge), forming ethylene and elhane. 
 
 The heats of formation of the carbon chlorides approach those of the hydrogen deriva- 
 tives very closely ; 
 
 = 21.0; (C2,Cl^-gas) = —I. I 
 
 2 . SILICON. 
 
 Atom : Si = 28.4. 
 
 Silicon next to oxygen is the most widely distributed element in nature. 
 Owing to its affinity for the latter it does not occur in a free condition. 
 Combined with oxygen as silicon dioxide (Si02),and in the form of salts 
 of silicic acid (silicates) it comprises many minerals and almost all the 
 crystalline rocks. 
 
 It was first obtained in a free condition by Berzelius (1823) on heating 
 potassium silico-fluoride (K2SiFlg) with metallic potassium: 
 
 K,SiFle + 4K = 6 KF 1 + Si. 
 
 The ignited mass is treated with water, which dissolves the potassium 
 fluoride and leaves the silicon as a brown, non-lustrous, amorphous pow- 
 der. It can also be obtained by heating silicon fluoride (SiFl^) with 
 sodium, or very readily by igniting a mixture of quartz sand (silicon di- 
 oxide) with magnesium powder [see Cl. Winkler, Ber. 23 (1890), 2652]: 
 
 2SiO., -f 2Mg = Si -f 2MgO. 
 
 It burns to silicon dioxide (SiOg) when heated in the air. 
 
 Another modification — the crystalline silicon — discovered by St. Claire 
 Deville, is obtained by fusing sodium silico-fluoride with aluminium or 
 with sodium and zinc. The separated silicon dissolves in the molten 
 zinc, and on cooling, deposits out in crystals, which remain on dissolving 
 the zinc in hydrochloric acid. The crystalline product consists of black, 
 shining octahedra and needles, of specific gravity 2.49, and of very great 
 hardness. U])on ignition in the air or oxygen it is not oxidized; it is not 
 
SILICON. 
 
 l6l 
 
 attacked by acids. On boiling it with a sodium or potassium liydroxide 
 solution it dissolves, forming a silicate and liberating hydrogen : 
 
 Si + 4KOH = K,SiO, + 211,. 
 
 Heated in chlorine gas, silicon burns to the chloride. 
 
 Hydrogen Silicide, SiH^, the analogue of CH^, is produced like 
 arsine and stibine by dissolving an alloy of silicon and magnesium in 
 dilute hydrochloric acid : 
 
 Mg, Si + 4HCI = SiH^ + 2MgCl2. 
 
 Such an alloy can be obtained by heating finely pulverized quartz or 
 glass with 1^4 times its quantity of magnesium powder. The escaping 
 
 hydride contains admixed hydrogen, has a disagreeable odor, ignites 
 spontaneously in the air, and burns to the dioxide and water : 
 
 SiH^ + 20 , = SiO, + 2H2O. 
 
 Perfectly pure silicide, free from hydrogen, is obtained by heating the 
 compound, SiH(O.C2H5)3, which will be described in Organic Chemistry. 
 In the air at ordinary pressure it ignites only upon warming; if, however, 
 the gas, by diminution of pressure or by the addition of hydrogen, is 
 diluted, it becomes spontaneously combustible at the ordinary tempera- 
 ture. A red heat decomposes the hydride into amorphous silicon and 
 hydrogen. When mixed with chlorine it inflames and probably forms 
 substitution products similar to those of methane (CHJ. Pure hydrogen 
 silicide condenses to a liquid at — 1° and a pressure of 100 atmospheres. 
 
 Silicon Chloride, SiCl^, results from the action of chlorine upon 
 silicon, or magnesium silicide, or by conducting chlorine over an ignited 
 mixture of the dioxide and carbon (Fig. 62) : 
 
 SiO, + 2C -f- 2CI, = SiCh + 2CO. 
 
 14 
 
i 62 
 
 MORGAN 1C CHEMISTRY. 
 
 The niixtiirc is i)laccd in a porcelain tube, wliicli is licated lo a red 
 heat. Tlie chlorine generated in the flask is dried in the two calcium 
 chloride towers and the silicon tetrachloride is condensed in the receiver 
 chilled with ice-water. While carbon or chlorine does notact separately 
 upon the dioxide, when they act simultaneously the reaction is induced by 
 the mutually supporting affinities of carbon for oxygen and of chlorine 
 for silicon. 
 
 Silicon chloride is a colorless liquid, having a sj^ecific gravity of 1.5, 
 and boiling at 59°. It fumes in the air, and is decomposed by water 
 into silicic and hydrochloric acids: 
 
 SiCh + 4lbO = IbSiO, q- 4IICI. 
 
 This is a remarkable distinction between it and carbon tetrachloride. 
 
 Silicon chloride may he employed in determining the atomic weight of silicon. Analysis 
 shows that it contains 7.1 parts of silicon for every 35.45 parts of chlorine. .Siij)posing, 
 from the great analogy of the silicon compounds to tho.se of carbon, that its formula is SiCfi, 
 the atomic weight of the silicon would be 28.4 : 
 
 Si = 28.4 
 
 Cl, = 141.80 (4 X 35 - 45 )- 
 SiCl, — 170.20 
 
 This .supposition is confirmed by the vapor density of the chloride and other deriva- 
 tives of silicon. 
 
 Silicon Bromide, SiBr,, and Silicon Iodide, Sil„ are formed 
 in the same manner as the chloride. The first is a colorless liquid, of 
 s])ecific gravity 2.8, becoming solid at — 12° and boiling at +153°. 
 The iodide forms colorless octahedra, melting at 120° and boiling at 290°. 
 Like the chloride, both are decomposed by water. 
 
 Besides these compounds, which may be viewed as hydrogen silicide, in which all of the 
 hydrogen is replaced by halogens (see p. 159), others exist in which only a part of this 
 element is replaced. Thus, silicon chloroform ^ SiHClg, corresponds to the chloroform 
 (CHClg) derived from methane. It is produced by the action of phosphorus pentachlo- 
 ride, or antimony pentachloride, on hydrogen silicide : 
 
 SiH, + 3SbCl5 = SiHCb + 3SbCl3 + 3HCI ; 
 
 or upon heating silicon in dry hydrogen chloride gas : in this case a mixture of the tetra- 
 chloride and silicon chloroform results (.see also Ber. 22, 188). These compounds may 
 be separated by fractional distillation. Silicon chloroform is a colorless liquid, of specific 
 gravity 1.6, and boils at 35-37°. The vapor density corre.sponds to the molecular 
 formula .SillClj. It fumes in the air, and decomposes with water into silicic and hydro- 
 chloric acids. 
 
 The silicon bromoform, SillBrj, and iodoform, SiHIj, are very similar to chloroform ; 
 they corre.spond to the analogous carbon compounds. 
 
 The compounds Si^Clg and Si.^Ig are known. They correspond to hexachlorethane, 
 CqClg. From all the.se data we ob.serve the analogy, based on formulas alone, between 
 .silicon and carbon. 
 
 Silicon Fluoride, SiFl„ 
 silica : 
 
 SiO., ^ 
 
 is formed when hydrofluoric acid acts upon 
 
 4IIFI := SiFl, -P 211 , 0 . 
 
 To y)reparc it, a mixture of fluorite and ])owdered glass, or sand (SiO.,), 
 is warmed with suliduiric acid ; by the action of sulphuric acid upon the 
 
SILICON. 
 
 163 
 
 fluorite h3'drogen fluoride (p. 64) is disengaged, and this reacts upon the 
 silicon dioxide in tlie manner indicated in the above equation. I'he liber- 
 ated gas is collected over mercury. 
 
 It is colorless, has a penetrating 
 odor, and fumes strongly in the 
 air. Its vapor density corresponds 
 to the molecular formula SiFl^. 
 
 Its deportment with water is very 
 characteristic ; it is decomposed 
 thereby into silicic acid (H^SiOj 
 and hydrogen silico-fluoride : 
 sSiFl, + 4 H ,0 = H.SiO, + 
 
 2H2SiFlg. For the execution of 
 this method, conduct the tetra- 
 fluoride through a glass tube 
 into a vessel containing water 
 (Fig. 63). Gelatinous silicic acid 
 separates out, and the gaseous 
 hydrogen silico-fluoride (H^SiFlg) 
 remains dissolved in the water. 
 
 As the separating silicic acid 
 may easily obstruct the opening 
 of the glass tube, the latter is 
 allowed to project a slight distance into mercury. The solid silicic acid 
 is separated from the aqueous solution by filtration. 
 
 Hydrogen Silico-fluoride, H2SiFlg, is only known in aqueous solu- 
 tion. Upon evaporation at a low heat it decomposes into the tetrafluor- 
 ide and hydrofluoric acid(SiFl^ and 2HFI). In its chemical deportment 
 it is an acid similar to the hydrogen-halogen acids. Its aqueous solution 
 reddens blue litmus-paper, dissolves many metals, and saturates bases, form- 
 ing salts with them, in which two hydrogen atoms are replaced by metal. 
 
 The potassium and barium salts dissolve with difficulty in water. 
 
 Fig. 63. 
 
 Silicon Carbide, Carborundum, SiC. This remarkable substance was discovered by 
 Acheson when, at the suggestion of Edison, he endeavored to dissolve carbon in fused and 
 highly heated aluminium silicate, hoping that when the mass cooled diamonds would appear. 
 Miihlhauser determined its composition. Carborundum is produced by heating a mixture 
 of sand and coke, with salt as a flux, in the electric furnace to about 3500° : SiO.^ -f- 3C 
 = SiC -f- 2CO. When pure it consists of brilliant green crystals of specific gravity 3.22. 
 It is insoluble in the ordinary solvents and burns with difficulty even in oxygen. It is 
 attacked by molten caustic alkalies and alkaline carbonates. In hardness it ranks next 
 to boron carbide and the diamond ; hence it is beaten into powder and used as an abra- 
 sive, .superior to anything known at present in this line. See Metallic Carbides ; also Z. 
 f. anorg. Ch. 5 (1894I, 104; O. N. Witt, Die chemische Industrie auf der Kolum- 
 bischen Weltaiusstellung zu Chikago (1894), 124; F. Ahrens, Die Metallkarbide und 
 ihre Verwendung, Stuttgart, 1896. 
 
 The heats of for 7 ualio 7 i of the halogen derivatives of silicon are as follows 
 
 (Si,Cb):- 157.6; (Si,I 5 r,) = 120.4; (Si,I,) - 58.0. 
 
164 
 
 INORGANIC CHF.MISIRY. 
 
 Their ready decoinposahility l)y water may he aceomited for if llie lieat f)f formation 
 of silieic acid he eonsidered : Si, O.^ ^ 219.0 Cal. 'I'lieir depoilmeiil toward oxy{;en is 
 
 also very interesting. The heat disengagement, in their transposition, following the 
 equation 
 
 Six, + 0 , = SiO., 4- 4X, 
 
 equals 61.4 Cal., with the ehloride, 98.6 Cal., with the hromide and with the iodide 
 161.0 [= 219-58) Cal. Therefore, oxygen must liberate the halogen atoms present in 
 them. And this occurs with the chloride and hromide when raised to a red heat, while 
 the iodide inflames in the air, yielding silicon dioxide and iodine vapor. 
 
 Tin, Sn = 118.5, recently discovered (1886) Germanium, 
 
 Ge == 72, attach themselves in chemical character to the groiij) contain- 
 ing carbon and silicon. They bear the same relation to the latter 
 elements which arsenic and antimony bear to the members of the nitrogen 
 group. This is evident in their atomic weights : 
 
 N = 14.04 P = 3i.o As =: 75 Sb = 120 
 
 C — 12.00 Si=28.4 Ge=72 80=118.5. 
 
 Germanium and tin, like carbon and silicon, form volatile compounds 
 with four atoms of the halogens, e. g., GeCl, and SnCl,. In a similar 
 manner they unite with oxygen to yield dioxides, e. g., GeOg and SnO^, 
 which resemble silicon dioxide in having an acid character. On the 
 other hand, they do not form volatile hydrogen derivatives, and are thus 
 markedly distinguished from the other non-metals. Therefore germa- 
 nium and tin, with their higher analogue. Lead, will be treated with the 
 metals. 
 
 In the preceding pages we have considered four groups of elements, 
 comprising all the so-called metalloids (with the exception of boron). 
 In each group the last members, possessing the highest atomic weights, 
 exhibit very distinct metallic properties. This is clearly the case with ger- 
 manium and tin, antimony and arsenic in the free condition. Tellurium 
 and selenium also (in the crystalline modification) possess marked metallic 
 appearance; finally, iodine has a metallic luster. The affinity for hydro- 
 gen, on the other hand, diminishes with increase in metallic character ; 
 the hydrides of iodine, tellurium, antimony, and arsenic are very un- 
 stable, and decompose readily into their constituents ; finally, ger- 
 manium, tin, and bismuth do not combine with hydrogen. 
 
 The remarkable relations between the atomic weights of the elements 
 of the four groups are noticeable in the following table: 
 
 C = 12.00 
 
 N = 14.04 
 
 0 = 16.00 
 
 FI = 19 
 
 Si = 28.4 
 
 P = 31.00 
 
 S = 32.06 
 
 Cl = 35-45 
 
 (}e= 72 
 
 As= 75 
 
 Se = 79.1 
 
 Br = 79.96 
 
 Sn = 1 18.5 
 
 Sh = 120 
 
 Te = 127 
 
 I = 126.85. 
 
 A fuller consideration of these relations will be given in the presenta- 
 tion of the Periodic System of the Elements. 
 
THE VALENCE OF THE ELEMENTS. 
 
 165 
 
 THE VALENCE OF THE ELEMENTS. CHEMICAL 
 STRUCTURE OF THE MOLECULES. 
 
 Each of the elements of the group of halogens forms but one compound 
 with hydrogen : the halogen hydride. In the case of the elements of the 
 other three groups just considered this simplicity occurs exceptionally. 
 Thus, oxygen and sulphur each form two hydrogen derivatives : 
 
 H2O and H2O2 ; H^S and H^Sg. 
 
 The hydrides are even more numerous with phosphorus and nitrogen, and 
 with carbon they are almost too numerous to mention. In this confusing 
 multiplicity simple relations can be reached if attention is directed solely 
 to that derivative of the metalloid which contains the most hydrogen — 
 to the compound in which the power of the metalloid to unite with hydro- 
 gen attains its maximum. We would then have the following hydrides, 
 which in turn are arranged in groups : 
 
 CH4 
 
 NH3 
 
 OH2 
 
 FIH 
 
 SiH^ 
 
 PH3 
 
 SH2 
 
 CIH 
 
 
 ASH3 
 
 SeH2 
 
 BrH 
 
 
 SbH3 
 
 TeH2 
 
 IH. 
 
 Not one of the many hundred combinations of carbon with hydrogen 
 contains more of the latter element than methane; ethane, C 2 Hg, follows 
 methane in this respect, but it contains only three atoms of hydrogen to 
 one of carbon, and in all the other hydrides of carbon less even than three 
 atoms of hydrogen are present for each atom of carbon. It must, there- 
 fore, be concluded that the carbon atom under the conditions accessible 
 and known to us is not capable of combining with more than four atoms 
 of hydrogen. The same is true of the other elements which have been 
 mentioned. Facts oblige us, therefore, to ascribe a peculiar function of 
 affinity to each element, in its relation to hydrogen, which it does not 
 exceed. This is called its valence or ato7nicity. The elements of the 
 fluorine group are univalent; the elements of the oxygen group bivakfit ; 
 nitrogen and its analogues, U'ivalent ; carbon and silicon, finally, are 
 quadrivalent elements. The lary;est number of hydro ge7i atoins with which 
 one atoni of an element can combine determines the valence of the latter. It 
 has been customary to attribute to the elements, in accordance with their 
 valency, affinity units or valence units, i. e., to say that carbon has 
 four affinity units, oxygen two, etc. This is intended, however, only 
 to indicate that one carbon atom can combine with but four hydro- 
 gen atoms, one oxygen atom with but two hydrogen atoms, etc. (see 
 
 p. .66). 
 
 The valence of the elements is frequently designated bylines or Roman 
 numerals placed above the symbols : 
 
 I II HI IV 
 
 Cl O N C. 
 
i66 
 
 INORGANIC CHEMISTRY. 
 
 The mutual union of affinity units is indicated by one line: 
 
 II— Cl 
 
 Hydroeen 
 
 chloride. 
 
 Water. 
 
 II 
 
 I 
 
 N 
 
 / \ 
 
 II II 
 Ammonia. 
 
 II 
 
 I 
 
 II— C— II. 
 
 I 
 
 H 
 
 Methane. 
 
 In these formulas the atoms of oxygen, of nitrogen, of carbon — indeed, 
 of all the multivalent elements — represent, as it were, the nuclei to which 
 the hydrogen atoms attach themselves. 
 
 'I'he hydrogen atoms in these molecules can be replaced or substituted 
 by other elements (p. 159). According to this, an atom of each of the 
 univalent halogens can replace one hydrogen atom: 
 
 II u 
 «<H 
 
 
 II II 
 
 o<ci 
 
 II pi III /II 
 
 o<ci 
 
 III /Cl 
 
 Sb^Cl 
 
 \ci. 
 
 Water. 
 
 
 Hypochlorous 
 
 Chlorine Nitrogen Antimony 
 
 
 
 acid. 
 
 oxide. iodide. 
 
 trichloride. 
 
 It is more convenient to employ brackets instead of the lines: 
 
 
 1 
 
 fll 
 
 1 
 
 fCl 
 
 fll 
 
 
 'c 
 
 Cl 
 
 'c 
 
 Cl 
 
 IV Cl 
 
 
 
 1 
 
 
 1 
 
 Cl 
 
 
 1 
 
 Cl 
 
 1 
 
 Cl 
 
 [Cl. 
 
 
 Chloroform. 
 
 Carbon 
 
 Silicon 
 
 
 
 
 tetrachloride. 
 
 chloroform. 
 
 By the replacement of hydrogen in the preceding hydrogen compounds 
 by the univalent metal potassium, we obtain 
 
 
 II H 
 
 0<K 
 
 Ill fH 
 
 IV 1 
 
 I H 
 
 1 H 
 
 K— Cl 
 
 Potassium chloride. 
 
 K 
 
 Ik 
 
 
 
 Potassium hydroxide. 
 
 Dipotassium 
 
 imide. 
 
 [K. 
 
 Potassium methide. 
 
 The bivalent elements, like oxygen and sulphur, replace two hydrogen 
 atoms in the compounds of the latter element : 
 
 III /Cl 
 Sb=0 
 
 Antimony 
 oxychloride 
 
 Finally, tri valent nitrogen can replace three atoms of hydrogen, or 
 three of the halogens : 
 
 III 
 
 IV f N I IV III 
 
 C\II or II-C=N. 
 
 Ilydroj^en cyanide. 
 
 Ill 
 
 IV f N I IV III 
 
 C\C1 or C1-C“N. 
 
 Cyanogen chloride. 
 
 IV /H 
 C^H 
 
 Methylene 
 
 oxide. 
 
 IV^O 
 
 Carbon 
 
 dioxide. 
 
 iv^S 
 
 Carbon 
 
 bisulphide. 
 
 IV^O 
 
 s<o. 
 
 Silicon 
 
 dioxide. 
 
THE VALENCE OF THE ELEMENTS. 
 
 167 
 
 In these substitutions, in which the elements replace one another atom for atom, we 
 possess a means of determining the molecular weights, previously found from the densities, 
 in a purely chemical manner. Thus, methane, the molecule of which must from its 
 density be represented by CH^, can have its hydrogen atoms replaced one after the other 
 by chlorine and four compounds, CH3CI, CH2CI2, CHCI3 and CCl^ result. Analysis has 
 determined their composition. If the molecule of methane were represented by a larger 
 formula, say CjHg, then upon chlorination there would result not four but eight different 
 chlorinated methanes : CgH^Cl, CjHgCl.^, C2H5CI3, etc., of which four (C2H.JCI, C2HgCl3, 
 C2H3CI5, C2HCIY) would, so far as composition is concerned, differ very decidedly from 
 the preceding. Since, however, never more than four chlorinated methanes have been 
 made, and never one which contained more hydrogen than CH3CI or one having less 
 hydrogen than CIICI3, we conclude in a purely chemical way that methane must have the 
 formula CH^. This method of determining molecular weights can also be applied where 
 the vapor density is not applicable. 
 
 These substitutions (replacements) afford a clue to the magnitude of the atomic 
 weights of many elements, in so far as, as has already been repeatedly emphasized, the 
 atouiic tveight is the smallest quantity of an element which is present in the moleczile of 
 any of its compounds. If it is found that the smallest quantity by weight of chlorine 
 which enters the molecule of any of its derivatives is to the replaced hydrogen as 
 35.45 : 1. 01 (O = 16) then it must be assumed that these numbers represent the atomic 
 weights of the two elements. 
 
 Like valence is designated by the word equivalence. One atom of chlorine is equiva- 
 lent to one atom of hydrogen ; 35.45 parts, by weight, of chlorine are, then, equivalent 
 to 1. 01 parts, by weight, of hydrogen. One atom of oxygen is equivalent to two atoms 
 of hydrogen ; consequently, 2.02 parts, by weight, of hydrogen are equivalent to 16.00 
 parts, by weight, of oxygen, or i.oi parts of hydrogen to 8 parts of oxygen. Further, 
 one atom of nitrogen, or 14.04 parts, is equivalent to three atoms or 3.03 parts of 
 
 hydrogen ; i.oi parts of hydrogen is, therefore, equivalent to = 4.68 parts of nitro- 
 gen, etc. These quantities, equal to 8 parts, by weight, of oxygen, are termed equiva- 
 lent weights y and were formerly employed instead of the atomic weights (p. 70). As 
 may be observed from the preceding, the equivalent weights of multivalent elements 
 are parts of the atomic weights which correspond to the valence of the elements. 
 
 If, consequently, the valence of the elements, in relation to hydrogen 
 (also to other elements), has a definite value, the question naturally 
 arises — what will result if an atom of hydrogen be withdrawn from the 
 saturated molecules, e. g., water, H2O, ammonia, NH3, or methane, CH^? 
 The resulting groups or residues : 
 
 II III IV 
 
 -O-H -N^H2 
 
 Hydroxyl. Amide. Methyl. 
 
 can plainly not exist in a free condition. When set free, these groups, 
 therefore (like the elementary atoms), unite and form saturated molecules. 
 Thus, for example, we obtain the bodies: 
 
 II II III III 
 
 HO-OH H2N-NII2 
 
 Hydrogen Diamide, 
 
 peroxide. 
 
 Ill III 
 
 H2P-PH2 
 
 Liquid 
 
 hydrogen 
 
 phosphide. 
 
 IV IV 
 H3C-CH3. 
 Dimethyl 
 or 
 
 Ethane. 
 
 Carbon is particularly inclined to such combination. Upon remov- 
 ing an atom of hydrogen from dimethyl or ethane (C2Hg) the so-called 
 ethyl group remains : 
 
 IV IV 
 
 C,U,- or CH3-CII2-. 
 
68 
 
 INORGANIC CHEMISTRY. 
 
 in wliicli one carlion affinity is nnsatnrated ; this can again unite with llie 
 methyl gronj), CII,. 'hhc resnlting com})onnd would be: 
 
 IV IV IV 
 
 cyig or HaC-CII^-CIIa or CjHg.CIIj. 
 
 By the continuation of this process of a chain-like union, as it were, 
 of the carbon atoms, we obtain a whole series of hydrocarbons (CJ I, 
 etc.) with the general formula (compare ])age 153). 
 
 Not only similar residues or groups, but dissimilar also, comlnne in 
 this way : 
 
 III II IV II IV HI 
 
 H^N-OII IIjC-OII II3C-NII2. 
 
 Hydroxylamine. Methyl hydroxide. Methylamiiie. 
 
 Such combinations are generally effected by reactions of double decomposition. 'Fluis 
 methyl hydroxide (wood-spirit) results from the action of methyl iodide (CII3I) upon 
 silver hydroxide (AgOII) : 
 
 CII3I -f AgOII = CII3 . Oil b- Agl ; 
 
 methylamine by the action of methyl iodide upon ammonia : 
 
 CH3I -f NII3 CII3 . NII2 + III. 
 
 Dimethyl is produced when sodium acts upon methyl iodide : 
 
 2CH3T + Na^ = CJIg + 2NaI. 
 
 By the withdrawal of iodine by the sodium the methyl groups are liberated, and then 
 combine with each other. 
 
 Further, the atoms with several valences can unite with two and three 
 affinities {liouble and triple union ; there being no intention to assert with 
 what firmness this union exists) : 
 
 III III IV IV IV IV 
 
 N^N H2C=CH2 HC=CH. 
 
 Ethylene. Acetylene. 
 
 o 
 
 Hyponitrous 
 
 oxide. 
 
 The molecules of elements result by the complete mutual union of 
 their atoms : 
 
 H-H 0-^0 
 
 Hydrogen. Oxygen. 
 
 N=N 
 
 Nitrogen. 
 
 P-P 
 
 P^P. 
 
 Phosphorus. 
 
 This manner of linking or combination of the atoms in the molecule 
 with single affinities is designated cheinical co?istitution, or chemical 
 structure oi compounds; the formulas representing them are called con- 
 stitution or structural formulas. Of course, the actual ])osition of the 
 atoms in space (of which we have no knowledge) is not indicated by the 
 chemical structure. The fundamental ])rinci])le of cliemical structure 
 consists in this, that the affinity unit of one atom unites with the affinity 
 unit of another atom. 
 
 'I'he following circumstances, however, complicate these simple rela- 
 
THE VALENCE OF THE ELEMENTS. 
 
 69 
 
 tions: Among the elements of the nitrogen group we saw that phos- 
 phorus and antimony combine with 3 and 5 atoms of chlorine and the 
 other halogens; that sulphur, selenium, and tellurium take up 2 and 4 
 atoms of chlorine and bromine; that iodine unites with i and 3 atoms 
 of chlorine and 5 of fluorine. Only the quadrivalent elements, carbon 
 and silicon, are capable of combining with 4 hydrogen atoms, and with 
 not more than 4 of the halogens : 
 
 IV 
 
 Ill 
 
 II 
 
 I 
 
 CCl, 
 
 PC13 
 
 SC12 
 
 ICl 
 
 
 V 
 
 IV 
 
 III 
 
 • • 
 
 PCI, 
 
 SCfi 
 
 IC13 
 
 V 
 
 . . 
 
 . . 
 
 . . 
 
 IF15. 
 
 Hence it appears that the elements (excepting carbon and silicon) do not 
 express such a constant valence in their relation to chlorine (and to the 
 halogens) as they do to hydrogen. Phosphorus and its analogues appear 
 to be tri- and quinquivalent ; the elements of the sulphur group, bi- and 
 quadrivalent; iodine finally appears to be uni-, tri-, and quinquivalent. 
 
 The varying valence of the metalloids shows itself more distinctly and 
 frequently in the more stable oxygen derivatives. We are acquainted 
 with the following oxygen compounds of the elements of the four groups 
 already mentioned ; the members of each group afford perfectly analo- 
 gous compounds : 
 
 II 
 
 I 
 
 III 
 
 I 
 
 CO 
 
 N20 
 
 S-^Og 
 
 C120 
 
 IV 
 
 II 
 
 IV 
 
 IV 
 
 C02 
 
 NO 
 
 S02 
 
 C102 
 
 
 III 
 
 VI 
 
 V 
 
 
 N2O3 
 
 SO3 
 
 I205. 
 
 
 IV 
 
 VII 
 
 
 . . 
 
 N02 
 
 S20, 
 
 . . 
 
 
 V 
 
 
 
 . . 
 
 N ,05 
 
 
 . . 
 
 The valence of iodine and of nitrogen reaches five, and that of sulphur 
 seven affinity units, assuming that oxygen is constantly bivalent and 
 wholly united to the other element. The elements of the nitrogen group 
 are not more than quinquivalent, both in their relation to the halogens 
 and oxygen : 
 
 PCI,, PF1„ SbCl,. 
 
 Carbon, finally, does not show more than four affinities for hydrogen, 
 chlorine, or oxygen. As we will observe later, these relations are more 
 apparent in the hydroxyl derivatives of the oxides — the acids. 
 
 Hence we conclude that valence is not an absolute property belonging 
 per se to and for the elements solely, but that it is dependent upon the 
 nature of the elements combining with one another and upon the exter- 
 nal conditions under which they unite. We can, in general, distinguish 
 two valences: the hydrogen valence and the halogen or oxygen valence. 
 The hydrogen valence is constant ior all elements ; for Cl = i, for 0 = 2, 
 for N = 3, for C = 4. The valence of most of the elements appears 
 *5 
 
INORGANIC CHEMISTRY. 
 
 I 70 
 
 to be variable in respect to oxygen and chlorine. It varies, indeed, as 
 seen from the formulas given above for the chlorine group from i to 4 to 
 5 ; for the elements of the sulphur group from 4 to 6, and from 3 to 7 ; 
 for those of the nitrogen group from i to 3 to 5, and from 2 to 4. 
 
 It is perfectly immaterial whether we ascribe a varialile valence to the 
 elements or accept the maximum as the true measure, and regard the 
 lower compounds as unsafurated, because we possess no concei)tion of 
 the nature of valence. Later, we will discover that this alteration of 
 valence finds full ex})ression and generalization in the periodic system of 
 the elements, which is based upon the grouping of the elements according 
 to their atomic weights. 
 
 From these facts, we have evolved the idea of a variable valence. Until recently 
 another opinion prevailed. This, denying the alterability, regards valence as an abso- 
 lute constant property of the elementary atoms. According to this idea, the true valence, 
 or atomicity, is only derivable from the hydrogen comjjounds ; the halogens are absolutely 
 univalent, the elements of the oxygen group bivalent, and those of the nitrogen group 
 trivalent, etc. Different propositions have been made in order to carry out the constant 
 atomicity for all compounds. A chain-like union is assumed in order to explain the 
 OKygen compounds. This is similar to the carbon atoms in the carbon compounds 
 (p. 168). It is represented in the following examples: 
 
 I II II II II II I I II II II II I 
 
 I_ 0 — O— O— O— O— I Cl— O— O— O— ( )— H 
 
 Iodic anhydride. Perchloric acid. 
 
 Sulphur dioxide. 
 
 II /O 
 
 s( )0 
 
 Sulphur trioxide. 
 
 II o— O— O— H 
 
 ^O— H 
 
 Sulphuric acid. 
 
 II III II II II III II 
 
 o N— O— O— O— N =0 
 
 Nitrogen pentoxide. 
 
 II III II II I 
 O N— O— O— H. 
 
 Nitric acid. 
 
 Hence it would appear that the oxygen atoms can unite without end, in a chain-like 
 manner, similar to the carbon atoms in the carbon compounds. However, a maximum 
 combining affinity of different groups for oxygen does actually occur ; this varies regularly 
 from group to group (for iodine = 5, for sulphur ^ 7, for nitrogen = 5, for carbon =4). 
 Consider it as we may, on the supposition of constant affinity, the reason for the different 
 tmmber of linking oxygen atoms must be based on the nature of the other element. The 
 fact, then, that the higher oxides or their hydrates (HCIO4, H2SO4, HNO3) are more 
 stable than the lower ( p. 182), argues decidedly against the chain-like union of the 
 oxygen atoms, which occurs in the unstable peroxides. 
 
 To explain the other compounds according to the constant atomicity theory, a differ- 
 ence is conceived between atomic and molecular compounds. The former are such as can 
 be explained by constant atomicity. All others are regarded as molecular compounds, 
 resulting from the union of two or more molecules, upon the basis of newly acquired 
 molecular affinities. Thus the compounds PCI5, SCI4, ICI3 are viewed as addition prod- 
 ucts of atomic compounds with chlorine molecules : 
 
 VC\,Cl, SCl2,Cl2 ICfiCb- 
 
 Phosphorus Sulphur Iodine 
 
 peutachloridc. lelrachloridc. trichloride. 
 
THE VALENCE OF THE ELEMENTS. 
 
 171 
 
 It was thought that a proof of these bodies being differently constituted from the real 
 atomic compounds was the fact that when vaporized they decomposed into simpler 
 derivatives ; the molecular compounds were not regarded as capable of existing as 
 such in the gaseous state. We saw, however, that the decomposition of the molecules 
 PCI5 and SCl^ is only gradual, increasing with the temperature, and that they do exist at 
 lower temperatures, undecomposed, in a vapor form (compare p. 141). True atomic 
 compounds, like sulphuric acid and nitric acid, frequently separate into simpler molecules, 
 in their conversion into vapor (compare Sulphuric Acid). 
 
 Phosphorus pentafluoride, PFI5, a gas at ordinary temperatures, is a strong argument 
 in favor of the quinquivalence of phosphorus ; iodine, too, forms a volatile pentafluoride, 
 IFI5. It is noteworthy that, as a general thing, the metalloids yield more stable and 
 higher compounds with the lower halogens (fluorine and chlorine) than with bromine 
 and iodine, which possess higher atomic weights (p, 142). The idea that molecular com- 
 pounds can also exist in vapor form, would render their distinction from the atomic com- 
 pounds purely arbitrary — not present in the nature of the bodies themselves. But other 
 gaseous compounds exist, which can, in no light, be regarded as molecular. Thus, the 
 usually quadrivalent or sexivalent tungsten (WClg, WBrg, WOCI4) forms a gaseous penta- 
 chloride, WCI5, and molybdenum, perfectly analogous to tungsten, yields a pentachloride, 
 M0CI5. Further, the quinquivalent vanadium (VdOClg) yields a gaseous tetrachloride, 
 VdCl^. 
 
 The salts of ammonia, according to the theory of constant atomicity, are not regarded 
 as ammonium compounds : 
 
 V V V 
 
 NH,C 1 , NH^NOg, (NHJgSO, 
 
 (p. 129), but as addition products of ammonia with acids : 
 
 NH 3 ,HC 1 , NH3,HNOg, (NH3)3,H3SO„ 
 
 which would make the analogy of the same with metallic salts appear rather doubtful. 
 Further, the properties of many compounds like 
 
 V V IV VI V 
 
 POCI3, (C3Hg)3PO, (CH3)3S.0H, (CH3)3S0, (CH3),N.0H, 
 
 and others, cannot be interpreted by the constant valence theory. The existence of 
 potassium permanganate, KMnO^, is not compatible with a constant bi- or quadrivalence 
 of manganese (see p. 178). 
 
 Yet, down to the present time, the acceptance of molecular additions could not be en- 
 tirely dispensed with, especially for the so called water of crystallization compounds; the 
 proposed effort, however, continues to reduce all such compounds on the basis of higher 
 valence of the elements. While, consequently, the constant valence theory comprises 
 only the so-called atomic compounds, the extended valence idea draws all others into the 
 circle of generalization. 
 
 We must, first of all, remember that the nature of chemical union and the cause of the 
 valence of the atoms are entirely unknown to us, and that, therefore, neither the idea of 
 a variable, nor yet of a constant valence constitutes a final explanation. Our task must 
 be restricted rather to representing the varying combination relations of atoms in accord- 
 ance with the facts and free them as far as possible from hypothetical additions. The 
 supposition of a constant valence, formerly given preference as the simpler, has shown 
 itself to be insufficient. Because it maintains real differences, where such are not percep- 
 tible, it departs from the ground of induction. The idea of a variable valence, which 
 takes all facts into consideration, is not prejudicial. It is their simplest expression. 
 Indeed, it finds a pregnant analogy in the deportment of the hydrocarbon radicals. By 
 the elimination of hydrogen from the saturated molecules (C2H6, for example), we obtain 
 radicals or groups of increasing valence (CjHg, C2.H3, C2H2'), The group 
 
 C2H2, however, is bi- and quadrivalent; the group Cgllg, uni- and trivalent. It is very 
 probable that the elementary atoms are, like these radicals, of a compound nature. 
 
172 
 
 INORGANIC CHEMISTRY. 
 
 The principles of chemical structure, just presented, appear most 
 clearly, and with the greatest regularity, with carbon. The consti- 
 tution of the innumerable varieties of carbon compounds is exjilained 
 by the quadrivalent nature of the carbon atoms, and their ability to com- 
 bine with each other by single affinities. In the other, the so-called 
 inorganic compounds, the valence and structure relations are more 
 complicated, and are far less investigated, but even in them so many 
 regularities appear that the actual material is greatly simplified therebv, 
 and is made more comprehensible. The doctrine of valence and struc- 
 ture is the first attempt to refer the facts underlying the law of multiple 
 proportions to the properties of the elementary atoms. As this theory 
 merely comprises actual relations, it cannot be negatived, but only further 
 developed. 
 
 The idea of varying valence of the atoms lay undeveloped as the basis of Gerhardt’s 
 type theory, according to which chemical compounds were referred to a few fundamental 
 forms or types : 
 
 H 
 
 H 
 
 11 
 
 Cl 
 
 H 
 
 IIN. 
 
 H 
 
 This mode of comparison will be repeatedly adopted in the following pages (p. 198). 
 After Edward Frankland (1852) had established the doctrine of the saturating capacity 
 of the elementary atoms, August Kekule added the type CIb to those just given and 
 solved in a clear and convincing manner the question how carbon atoms were able to 
 combine with one another and to satisfy their affinities (1857-1859). In this way he 
 developed one of the fundamentals of the chemistry of to-day — doctrine of the linking 
 of the atoms. See the historical sketch in Richter’s Organic Chemistry. 
 
 OXYGEN COMPOUNDS OF THE METALLOIDS. 
 
 Almost all the oxygen derivatives of the metalloids yield acids with 
 
 water : 
 
 H2O = 2HIO3; 
 
 Iodic 
 
 SO3 -f H2O 
 
 = H2SO, ; 
 
 Iodine 
 
 Sulphur 
 
 Sulphuric 
 
 pentoxide. 
 
 acid. 
 
 trioxide. 
 
 acid. 
 
 
 P 2 O 5 + 3^2^^ 
 
 Phosphorus 
 
 pentoxide. 
 
 = 2H3PO4. 
 
 Phosphoric 
 acid. 
 
 
 Conversely, these oxides may be obtained by the removal of water from 
 the acids; therefore, they are ordinarily termed anhydrides of the corre- 
 sponding acids: laO^, anhydride of iodic acid; SO3, sulphuric anhydride; 
 PjOj, phosphoric anhydride, etc. 
 
 Salts are formed by the replacement of the hydrogen of the acids by 
 metals. The acids are distinguished as monobasic, dibasic, tribasic, 
 and polybasic, de])ending upon the number of their hydrogen atoms 
 reidaceable by metals. 
 
 Like the metalloids, the metals possess different valences ; the univalent 
 metals (sodium, potassium, silver) each replace one hydrogen atom; 
 
OXYGEN COMPOUNDS OF THE HALOGENS. 
 
 173 
 
 those of higher valence replace more. Therefore the metals of higher 
 valence can unite several acid residues. This explains the import, e. g., 
 of the following chemical formulas : 
 
 (1)2 
 
 K2SO, 
 
 Potassium sulphate. 
 I 
 
 KNO3 
 
 Potassium nitrate. 
 
 I I II 
 
 KNaSO^ CaSO^ 
 
 Potassium-sodium sulphate. Calcium sulphate. 
 
 II III 
 
 Cu(N 03)2 Bi(N 03 ) 3 . 
 
 Copper nitrate. Bismuth nitrate. 
 
 The salts of sulphuric acid are called sulphates, those of nitric acid, 
 nitrates, those of phosphoric acid, phosphates, etc. The symbols of the 
 metals are sometimes written at the beginning and sometimes at the end 
 of the chemical formulas of the salts ; in the second case we mean to indi- 
 cate that the metal atoms are in union with the oxygen. 
 
 1. OXYGEN COMPOUNDS OF THE HALOGENS. 
 
 Oxygen derivatives of fluorine are not known. Iodine is the only one 
 of the remaining halogens which unites directly with oxygen ; how'- 
 ever, the oxides of chlorine may be prepared by indirect methods. The 
 only known acids of bromine are those corresponding to the oxides. 
 
 Chlorine forms the following anhydrides and acids : 
 
 Anhydrides. Acids. 
 
 CljO HCIO, Hypochlorous acid. 
 
 . . (HCIO2, Chlorous acid.) 
 
 CIO2 or CI2O, 
 
 . . HCIO3, Chloric acid. 
 
 . . HCIO4, Perchloric acid. 
 
 The compound Cl20^ exists, and must be regarded as a mixed anhy- 
 dride of chlorous and chloric acids. Chlorous acid itself is not known 
 free, but is stable in the form of salts. Further, the anhydrides of chlorous, 
 chloric and perchloric acids, having the formulas CI2O3, CljO^and Cl20^, 
 are not known. 
 
 The following formulas express the chemical structure of these com- 
 pounds : 
 
 I II I 
 Cl-O-Cl 
 
 Hypochlorous 
 
 anhydride. 
 
 I II I 
 Cl-O-H. 
 Hypochlorous 
 acid. 
 
 Ill V 
 
 OCI-O-CIO2 
 
 Chlorous-chloric 
 
 anhydride. 
 
 V 
 
 CIO2-OH. 
 
 Chloric 
 
 acid. 
 
 VII 
 
 CIO3-OH. 
 Perchleric acid. 
 
174 
 
 INORGANIC CHEMISTRY. 
 
 In tlic acids vve assume tlie presence of the univalent group OH 
 (hydroxyl or water residue) the hydrogen of which can be replaced by 
 metals, through the action of metals or bases. The grouj) (CIO,^ or CIO.,) 
 combined with hydroxyl is called an acid 7 'esidue, or radical. In the 
 anhydrides of the monobasic acids two acid radicals are united by an 
 oxygen atom; water converts them into 2 molecules of acid : 
 
 g>0 + 0<J| 
 
 Cl-OII 
 
 Cl-OII. 
 
 The salts of perchloric acid are called perchlorates ; those of chloric 
 acid, chlorates ; those of chlorous acid, chlorites ; and those of hypo- 
 chlorous acid, hypochlorites. 
 
 Hypochlorous Oxide, Cl^O (chlorine monoxide, hypochlorous an- 
 hydride), is i)roduced by conducting dry chlorine gas over ]jrecipitated 
 mercuric oxide, cooled by ice or cold water; the brown oxychloride of 
 mercury is formed at the same time : 
 
 2HgO + 2CI, = IlgpCl, + Cip. 
 
 The precipitated mercuric oxide should be heated to 300° before it is used, otherwise 
 it will be acted upon too energetically by the chlorine, and oxygen instead of hypochlo- 
 rous oxide will result. Red mercuric oxide is affected too slowly by chlorine. 
 
 The disengaged gas is condensed in a bent glass tube, cooled by a 
 freezing mixture of ice and calcium chloride. 
 
 It is a reddish-brown liquid, boiling at -j-5°. In the gaseous condition it 
 has a yellow-brown color, an extremely disagreeable odor, and it attacks the 
 respiratory organs very strongly. The vapor density corresponds to the 
 molecular formula, CI2O = 86 9. The gaseous or liquid oxide is very 
 unstable and explosive. If heated, or if brought in contact with a flame, 
 or exposed to the action of the electric spark, it decomposes with deto- 
 nation into oxygen and chlorine : 
 
 2CI.2O = 02 + 2CI2. 
 
 2 vols. I vol. 2 vols. 
 
 It is similarly exploded when it is brought in contact with sulphur, 
 phosphorus, and organic substances, such as caoutchouc. It has strong 
 oxidizing and bleaching properties. It dissolves in water to hypo- 
 chlorous acid : 
 
 CI2O + H2O = 2IICIO. 
 
 It yields calcium hypochlorite and chlorine when it is conducted over 
 porous calcium chloride : 
 
 CaCl2 -}- 2CI2O = Ca(C 10)2 + 2CI2. 
 
 Hypochlorous Acid, HCK), is only known in aqueous solution. 
 Its presence in chlorine water has been mentioned (p. 52). It is 
 ol)tained liy conducting clilorine into water in which there is suspended 
 
HYPOCHLOROUS ACID. I 75 
 
 freshly precipitated mercuric oxide; the liquid is agitated constantly, 
 while at the same time it is cooled and shut olf from the light : 
 
 HgO - 1 - 2CI2 -h H2O = HgCh + 2HCIO. 
 
 The mercuric chloride produced combines with the excess of oxide to 
 form oxychloride (see p. 174). After filtration the solution may be dis- 
 tilled ; a strong chlorine odor is usually noticeable at the beginning. The 
 first distillates are especially rich in hypochlorous acid. The free chlorine 
 is expelled by blowing air through the solution (Balard). 
 
 Hypochlorous acid may be similarly prepared by conducting chlorine 
 into water in which chalk is suspended : 
 
 CaCOg -j- 2CI2 T HjO = CaCq -f- CO2 T 2HCIO (Williamson) ; 
 
 also by the distillation of bleaching lime solutions with as much dilute 
 acid as is required for the decomposition of the hypochlorite, e. g. : 
 
 NaClO -f NaCl + HNO3 = NaNOg -f NaCl -f HCIO (Gay-Lussac). 
 
 Boric acid is well adapted for this purpose as it does not decompose 
 the chloride. The concentrated aqueous solutions (5 per cent.) are 
 yellow in color. Upon distillation there ensues a partial decomposition 
 into chloric acid and chlorine: 
 
 {a) 3HCIO = 2HCI + HClOg ; 
 
 \b) 2HCI + 2HCIO = 2CI2 + 2H2O. 
 
 The decomposition is very slight with dilute solutions. These are, 
 unlike the concentrated liquids, quite stable in the light. Their odor is 
 peculiar and they oxidize and bleach energetically. The bleaching action, 
 due to the separation of oxygen in statu nascendi, is twice as great as that 
 of free chlorine, as is evident from the following equations: 
 
 CI2 + H2O = 2HCI -f O, 
 
 2HCIO = 2HCI + O2. 
 
 Hydrochloric acid decomposes hypochlorous acid into chlorine and 
 water, when just twice as much chlorine is evolved as there is present in 
 the hypochlorous acid. 
 
 The acid itself is very feeble, more feeble than carbonic acid, hence it 
 is produced free if chlorine be conducted into aqueous solutions of car- 
 bonates : 
 
 K2COg -f CI2 -h H2O = KCl -f- KHCOg -f HCIO. 
 
 Its salts (bleaching powder, see Chloride of Lime) are formed together 
 with chlorides by the action of chlorine, in the cold, upon strong bases: 
 
 2 NaOH + CI2 NaCl -f NaClO + H2O. 
 
 Within certain limits this transposition takes place independently of 
 the temperature and strength of solution if chlorine is not present in 
 excess, otherwise the formation of chlorates will occur even at low tem- 
 peratures. When the solutions are heated two decompositions of the 
 
INORGANIC CHEMISTRY. 
 
 I 76 
 
 hy})ochlorite ])roceed simultaneously; in one instance chloride and 
 chlorate result : 
 
 3 NaC 10 ^ 2NaCl + NaClOj, 
 and in the other case chloride and oxygen : 
 
 2 NaC 10 = 2NaCl + O,. 
 
 [Compare Foerster and Jorre, Jr. prakt. Ch. 59 (1899), 53 I.] 
 
 A hypochlorite is also produced when chlorine is conducted over cal- 
 cium hydrate powder (see Bleaching Lime). 
 
 Oh shaking the aqueous solution of hypochlorous acid with mercury, there is produced 
 a yellow-brown precipitate of ligO. llgCh, soluble in hydrochloric acid (salts of hypo- 
 chlorous acid form HgO). This behavior serves to distinguish hypochlorous acid from 
 chlorine, which under like circumstances forms mercurous chloride, Ilg^Ck, insoluble in 
 hydrochloric acid (reaction of Wolters). 
 
 Chlorine Trioxide, CI2O3 = (C10)20, the anhydride of chlorous 
 acid (p. 131), and the free Chlorous Acid, HClOj = CIO. OH, are 
 unknown. It was formerly supposed that the tri oxide resulted by the 
 reduction of chloric acid, if potassium chlorate, KCIO3, was decomposed 
 by nitric acid in the presence of reducing substances (arsenic trioxide, 
 sugar and tartaric acid). Later research has, however, shown that in 
 such decompositions it is not the trioxide, but the dioxide, CIO2, mixed 
 with chlorine and oxygen, which results. 
 
 The salts of chlorous acid, the chlorites, e. g . , potassium chlorite, are formed by the 
 transposition of the aqueous solution of chlorine dioxide with the alkalies (see below). 
 The silver salt, KgQXO^, and the lead salt, Pb(C102).^, are precipitated from the aqueous 
 solutions of the alkali salts by silver nitrate and lead acetate. They dissolve with diffi- 
 culty in cold water, and crystallize from hot water in yellowish-green scales. They are 
 decomposed by carbonic acid ; the liberated chlorous acid decomposes immediately. 
 
 Chlorine Dioxide, CIO2, or Chlorine Tetroxide, CI2O4, formerly 
 called hypochloric acid, is formed when concentrated sulphuric acid acts 
 in the cold upon potassium chlorate. The chloric acid, which first 
 appears, at once decomposes into chlorine dioxide, water, and the stable 
 perchloric acid : 
 
 3HCIO3 =z CI2O, -h HCIO, + HjO. 
 
 It is best obtained (mixed with carbon dioxide) by adding dilute sul- 
 jdmric acid (with 2 parts of water) to a mixture of potassium chlorate (i 
 ])art) and oxalic acid (4.5 parts). Upon the application of a gentle heat 
 a yellow-green gas escapes. It can be condensed by a freezing mixture 
 to a reddish-brown liquid, boiling at 9.9° (under a pressure of 740 mm.), 
 and solidifying at — 79°. This liquid may be distilled without decom- 
 position from glass vessels, free from organic matter, if the temperature 
 does not exceed 30°. Violent explosions occur at higher temperatures, 
 so that it is advisable, if ])roper precautions are not taken, not to pour 
 sulphuric acid U])on ])otassium chlorate. Liquid as well as gaseous chlo- 
 rine-dioxide is very explosive, especially in contact with organic matter 
 and when heated, but sunlight is without effect. 
 
 The formation of the oxide can be effected in a perfectly harmless way. 
 
CHLORIC ACID. 
 
 177 
 
 and its powerful oxidizing action be illustrated, by throwing some potas- 
 sium chlorate and a few pieces of yellow phosphorus into water, con- 
 tained in a measuring glass, then allowing sulphuric acid to touch the 
 bottom of the tube, drop by drop, by means of a pipette. By the action 
 of the disengaged dioxide the phosphorus will burn under water with a 
 brilliant light. 
 
 When concentrated sulphuric acid is added to a mixture of potassium 
 chlorate and sugar, a violent combustion occurs. 
 
 Chlorine dioxide dissolves in water with a yellow color. Alkalies de- 
 colorize the solution, forming salts of chloric and chlorous acids ; 
 
 no + 2KOH = CIO.OK + CIO2.OK -f H^O. 
 
 Potassium Potassium 
 chlorite. chlorate. 
 
 Sunlight decomposes the aqueous solution of chlorine dioxide into 
 chloric acid, oxygen, and chlorine : 
 
 3CI2O4 + 2H2O = 4CIO2.OH + O2 4 - CI2. 
 
 The gas density of the dioxide corresponds to the molecular formula 
 CIO2. It is, however, very probable that at lower temperatures, in the 
 liquid condition or in aqueous solution, the molecules possess the doubled 
 formula, Cl20^, and that then the compound represents the mixed anhy- 
 dride of chloric and chlorous acids : 
 
 Cl O — '>0 
 
 - ci02->^- 
 
 The decomposition of its aqueous solution by alkalies is an argument 
 in favor of this view, as is also its analogy with nitrogen dioxide, NOj, 
 or nitrogen tetroxide, N2O4 (see this), the existence of both of which 
 molecules has been proved. 
 
 Chloric Acid, HCIO3 or CIO2.OH, is obtained by decomposing 
 an aqueous solution of barium chlorate with sulphuric acid : 
 
 Ba(C103)2 + H2SO, = BaSO, + 2HCIO3. 
 
 Barium chlorate. Barium 
 
 sulphate. 
 
 The barium sulphate separates as a white, insoluble powder, and can 
 then be filtered off from the aqueous solution of the acid. This is con- 
 centrated without decomposition, under diminished pressure, until the spe- 
 cific gravity becomes 1.28, and it then contains about 40 per cent, of 
 chloric acid ; it is oily and, when heated to 40°, decomposes into chlorine, 
 oxygen, and perchloric acid, HCIO^. The concentrated aqueous solu- 
 tion oxidizes strongly; sulphur, phosphorus, alcohol, and paper are 
 inflamed by it. Hydrochloric acid eliminates chlorine from the acid and 
 its salts: 
 
 HCIO3 + SHCI = 3II2O + 3CI2. 
 
 The chlorates are slowly formed when aqueous hypochlorite solutions 
 are heated ; the latter partly break down into oxygen and chlorides. To 
 
178 
 
 INORGANIC CHEMISTRY. 
 
 obtain the chlorate the liypochlorite solution should be slightly super- 
 saturated with chlorine; then the chlorate will form even in the cold but 
 much more rapidly if the solution be heated to 130°. This may be thus 
 explained. The excessive chlorine liberates hy^jochlorous acid from the 
 hypochlorite : 
 
 2KCIO + 2CI2 f 211,0 = 2KCI + 4IICIO, 
 
 which in turn is rapidly and completely trans])osed by the hypochlorite 
 into chlorate when chlorine or hydrochloric acid is again liberated : 
 
 or 
 
 2KCIO T 2IICIO = KCIO3 + KCl + 11,0 + Cl,, 
 KCIO + 2IICIO = KCIO3 + 2IICI ; 
 
 and hypochlorous acid is evolved by the liberated chlorine or hydro- 
 chloric acid, etc. The statement previously made that chlorine acting 
 upon cold lyes produces hypochlorites, and chlorates when the solutions 
 are hot, does not harmonize with the facts. The chlorate formation is 
 rapid and complete only when chlorine is in excess (compare Foerster 
 and Jorre). The electrolytic production of chlorates from chlorides will 
 be described in connection with potassium chlorate. 
 
 Perchloric Acid, HCIO^ or CIO3 . OH. This is the most stable 
 of all the oxygen derivatives of chlorine. Its sodium salt is present in 
 Chile saltpeter. As previously stated, it is produced by the decomposi- 
 tion of chloric acid, but is more easily obtained from its salts. Upon 
 heating potassium chlorate to fusion, oxygen escapes and potassium per- 
 chlorate results : 
 
 2KCIO3 = KCIO4 -f KCl -f O,. 
 
 When this decomposition approaches completion, the fused mass becomes 
 a thick liquid and finally a solid. As potassium perchlorate dissolves with 
 difficulty in water it can be readily separated from the chloride formed 
 simultaneously. 
 
 A solution of perchloric acid containing a little sodium chloride, but applicable for all 
 analytical purposes, may be prepared as follows : Commercial sodium chlorate is con- 
 verted by heat into perchlorate and chloride. Concentrated hydrochloric acid is poured 
 upon the powdered mixture and dissolves almost nothing but perchloric acid : 
 
 NaClO, + HCl = NaCl + HCIO^. 
 
 The solution is filtered and evaporated until dense white fumes of perchloric acid com- 
 mence to escape ; by distillation under greatly reduced pressure the dihydrate may be 
 obtained [Z. f. anorg. Ch. 9 (1895), 342; 13 (1897), 166]. 
 
 Roscoe obtained a monohydrate of perchloric acid by distilling potas- 
 sium perchlorate with four times its quantity of concentrated sulphuric 
 acid until the drops, passing over, no longer solidified. The mass was 
 then heated to 110° until crystals a])peared in the neck of the small 
 retort. From this point on anhydrous perchloric acid distilled over. 
 
 The pure acid is a mobile, colorless liquid, fuming strongly in the air; 
 its specific gravity is 1.78 at 15°. It cannot be preserved even in the 
 dark, since after a few days it decomposes with violent explosion. Heat 
 jiroduces the same result. It also explodes in contact with phosphorus, 
 
BROMIC ACID. 
 
 179 
 
 paper, carbon, and other organic substances. It produces painful 
 wounds when brought in contact with the skin. It absorbs water with 
 avidity, and with one molecule of the same forms the crystalline hydrate 
 HClO^-f H2O, melting at 50° and passing at 110° into the anhydrous 
 acid and the dihydrate, HCIO^ + 2H2O : 
 
 2 C 10 ,H . H2O = HCIO, + HCIO^ 2H2O. 
 
 The dihydrate may also be obtained by evaporating the aqueous solutions 
 of perchloric and chloric acids. It is a colorless, oily liquid of com- 
 paratively great stability. It boils unchanged at 203°. Its specific gravity 
 equals 1.82. 
 
 Perchloric acid is also formed at the positive pole by the electrolysis 
 of aqueous solutions of inorganic chlorine compounds; this is due to the 
 liberation of oxygen at the kathode. 
 
 Bromine yields the following compounds with oxygen and hydrogen : 
 
 HBrO, Hypobromous acid. 
 
 HBrOg, Bromic acid. 
 
 HBrO^, Perbromic acid. 
 
 The corresponding anhydrides are not known. The acids are perfectly 
 analogous to the corresponding chlorine compounds. 
 
 Hypobromous Acid, HBrO, is formed when bromine water acts 
 upon mercuric oxide ; the aqueous solution can be distilled in vacuo, and 
 possesses properties very similar to those of hypochlorous acid. 
 
 Bromic Acid, HBrOg. Bromates are like the chlorates. An 
 aqueous solution of the acid can be obtained from the barium salt by 
 decomposing the latter with sulphuric acid. A more practical method 
 of getting the free acid is to let bromine act upon silver bromate : 
 
 SAgBrOg -f 6Br -f 3H2O = SAgBr -f fiHBrOs, 
 or to oxidize bromine with hypochlorous acid : 
 
 5CI2O + Br2 -f H2O 2HBr03 + loCl. 
 
 The aqueous solution may be concentrated in vacuo until its content 
 reaches 50.6 per cent. HBrOg, and then closely corresponds to the 
 formula HBrO, -f- yH^O. When heated it decomposes into bromine, 
 oxygen and water. Heat breaks down the alkali bromates into bromide 
 and oxygen without the production of a perbromate. 
 
 Perbromic Acid, HBrO^, is said to be formed in the action of 
 bromine vapor upon perchloric acid : 
 
 IICIO, + Br = HBrO, -f Cl, 
 and is perfectly similar to the latter. 
 
i8o 
 
 INORGANIC CHEMISTRY. 
 
 Iodine forms llie following anhydrides and acids; 
 
 I-A IIIO3 11104.211/) 
 
 loclitie pentoxide, Iodic acid. Periodic acid (dihydrate). 
 
 Iodic anhydride. 
 
 Iodic Acid, HIO3. The potassium and sodium salts of this acid 
 occur in Chile saltpeter and are at present the best sources for iodine (see 
 p. 55). The iodates are formed in the same manner as the chlorates and 
 bromates, by dissolving iodine in a hot solution of j)otassium or sodium 
 hydroxide : 
 
 6KOII-|-3l, = 5KI 4- KI()3 + 311.0. 
 
 Upon adding barium chloride to this solution sparingly soluble barium 
 iodate separates; it may be transposed by sulphuric acid into insoluble 
 barium sul})hate and free iodic acid. 
 
 When iodine acts upon cold .sodium or potassium hydrate hypoiodites are formed. 
 Schonbein suspected this, but Lonnes proved it : 
 
 2NaOII + 2l = NalO + Nal + Iip. 
 
 These salts, however, quickly pass into iodates : 
 
 3XaIO = NalOa + ^Nal. 
 
 Therefore iodine acts upon hydroxides just like chlorine and bromine, but with iodine 
 the second step of the reaction, the formation of iodate, proceeds very rapidly at the 
 ordinary temperature. The experiments of Binz and also those of Lonnes, however, 
 show that iodine can remain uncombined for quite a time when dissolved in alkaline 
 liquors if potassium iodide be present. 
 
 Willgerodt and V. Meyer have prepared organic derivatives of the yet unknown 
 III 
 
 H — I O. Indeed, Meyer has made organic, strongly basic compounds which may be 
 
 III 
 
 derived from iodoniiim hydroxide, H2 = I — OH, also unknown. The chemical struc- 
 ture of these derivatives forbids the assumption that iodine is invariably univalent (see 
 Ber. 27 (1894), 426, 1592, and also Richter’s Organic Chemistry). 
 
 The free acid can be obtained by the oxidation of iodine with strong 
 nitric acid, or by means of chlorine : 
 
 3I2 + loHNOj = 6HIO3 + loNO + 2H2O. 
 
 Iodates are also produced by the action of iodine upon chlorates or bro- 
 mates in aqueous solution, whereby the iodine directly eliminates the 
 chlorine and bromine: 
 
 KCIO3 + I = KIO3 + Cl. 
 
 Upon evaporating the aqueous solution the free iodic acid crystallizes in 
 colorless rhombic jirisins of sj)ecific gravity 4.63. When iodic acid is 
 heated to 170° it decomposes into water and iodic anhydride : 
 
 211103 = 1203 + 1130. 
 
 It is decomposed, like chloric acid, by hydrochloric acid : 
 
 2IIIO3 + loIICl = I2 + 5CI2 + 6II2O. 
 
PERIODIC ACID. 
 
 l8l 
 
 Reagents, like hydrogen sulphide, H2S, sulphur dioxide, SOj, and hydri- 
 odic acid, HI, reduce it to iodine : 
 
 HIO3 + 5HI - 3H3O + 3lr 
 
 Iodic Anhydride, IjO^, is a white erystalline powder, which dis- 
 solves in water to form iodic acid. It decomposes at 300° into iodine 
 and oxygen. It can be obtained directly from ozone and iodine. 
 
 Periodic Acid, HIO^. Normal periodic acid is not known. The 
 hydrate, HIO^ . 2H2O, is produced by the action of iodine upon per- 
 chloric acid : 
 
 2HC10^ -f I2 + 4H2O = 2 [HIO^. 2H2O] + Cb- 
 
 Upon the evaporation of the aqueous solution, the acid crystallizes in 
 colorless forms, which deliquesce, melt at 130°, and at about 140° decom- 
 pose into water, oxygen and periodic anhydride : 
 
 2 (HI 0 ,. 2H2O) = 1 , 0 , + O2 + 5H2O. 
 
 Periodates result on conducting chlorine into hot alkaline solutions of 
 iodides or iodates. A sodium periodate has been found in Chile salt- 
 peter. 
 
 The existence of the hydrates of periodic and perchloric acids as well as of many 
 others (see Sulphuric and Nitric Acids), which were once regarded as molecular compounds, 
 is interpreted at present by the acceptance of hydroxyl groups, directly combined with 
 the element of higher valence : 
 
 VII 
 
 CIO^H -|- H2O = €102(011)3, Trihydrate or trihydric acid. 
 
 VII 
 
 CIO4H 2H2O = CIO (OH)^, Pentahydrate or pentahydric acid. 
 
 VII 
 
 CIO^H 3H2O = C1(0H)7, Heptahydrate or heptahydric acid. 
 
 The maximum hydrates, ClfOH).^ and I(OH).j, in which all seven affinities of the 
 halogen atom are attached to hydroxyl groups, are not known, but probably exist in 
 aqueous solution. As they give up water, and one atom of oxygen becomes simultaneously 
 united with two bonds to the halogen, they yield the lower hydrates — even to the mono- 
 hydrate CIO3 . OH. Perchloric acid continues monobasic in the polyhydrates, since but on^ 
 hydrogen atom is replaced by metals : 
 
 C\ 0 ,U, -b KOH =: CIO4K + 3H2O. 
 
 On the other hand, periodic acid (IO3 . OH) is not only monobasic, but as a pentahydrate, 
 I0(0H)5 = HIO4 . 2H2O, can, like the polybasic acids, furnish also polymetallic 
 salts, as : 
 
 VII / (011)3 VII f ( 0 H )3 VII VII 
 
 10 \ (0Na)2 10 \ (0Ag)2 I0(0.Na)5 I0(0Ag)3. 
 
 Salts also exist which are derived from condensed polyiodic acids, as : 
 
 ^O , Diperiodic acid, etc. 
 
 io<(on)/ 
 
 (Compare Uisulphuric Acid, Dichromic Acid, Pyrophosphoric Acid, etc.) 
 
i 82 
 
 INORGANIC CHEMISTRY. 
 
 The existence of such salts plainly indicates that the hydrates of acids must be looked 
 upon as hydroxyl compounds, and that iodine and the halogens are, in fact, heptads in 
 their highest combinations. 
 
 The oxygen compounds of the halogens in some respects display a charac- 
 ter exactly opposite to that of the hydrogen derivatives. While the affinity 
 of the halogens for hydrogen diminishes with increasingatomic weight from 
 fluorine to iodine (see p. 65), the affinity for oxygen is the exact reverse. 
 Fluorine is not capable of combining with oxygen; the chlorine and 
 bromine compounds are very unstable, and are generally not known in 
 free condition ; the iodine derivatives, on the contrary, are the most stable. 
 In accord with this is the fact that in the higher oxygen compounds 
 chlorine and bromine are set free by iodine, while in the hydrogen and 
 metallic compounds of the halogens the direct reverse is the case, viz., 
 that iodine and bromine are replaced by chlorine. 
 
 Further, the oxygen compounds exhibit the remarkable peculiarity that 
 their stability increases with the addition of oxygen. The lowest acids, 
 HCIO, HBrO, are very unstable, even in their salts; they possess a very 
 slight acid character, and are separated from their salts by carbon 
 dioxide. The most energetic and most stable are the highest acids, 
 HCIO4, HBrOg, HIO^, in which the higher valence of the halogens 
 appears. In the sulphur and nitrogen groups those oxides, in which the 
 elements manifest their maximum valence, are the most stable (compare 
 p. 169). 
 
 2. OXYGEN COMPOUNDS OF THE ELEMENTS OF 
 THE SULPHUR GROUP. 
 
 The elements sulphur, selenium, and tellurium combine with oxygen in 
 several proportions. Sulphur and oxygen form the following derivatives : 
 
 S,03 
 
 Sulphur 
 
 sesquioxide. 
 
 SO3 
 
 Sulphur trioxide, 
 Sulphuric anhydride. 
 
 SO2 
 
 Sulphur dioxide, 
 Sulphurous anhydride. 
 
 S2O,. 
 
 Sulphur heptoxide, 
 Persulphuric anhydride. 
 
 The oxygen derivatives of selenium and tellurium correspond to sulphur 
 di- and trioxides : 
 
 Se02 wSeOj 
 
 Selenium dioxide, Selenic 
 
 Selenious anhydride. anhydride. 
 
 TeOj 
 
 Tellurous 
 
 anhydride. 
 
 TeOg. 
 
 Telluric 
 
 anhydride. 
 
 Each of these oxides, sulphur sesquioxide excepted, combines with one 
 
OXYGEN COMPOUNDS OF SULPHUR. 1 83 
 
 molecule of water to form a dibasic acid (p. 172), of which the respective 
 oxide is the anhydride. The acids of sulphur are : 
 
 H2SO3 HjSO^ H2S20g. 
 
 Sulphurous Sulphuric Persulphuric 
 
 acid. acid. acid. 
 
 By the replacement of one atom of hydrogen by one atom of metal the 
 so-called acid or primary salts result, while the neutral or secondary salts 
 are obtained by the replacement of both hydrogen atoms : 
 
 KHSO^ K2SO,. 
 
 Acid potassium sulphate, Neutral potassium sulphate, 
 
 Monopotassium sulphate. Dipotassium sulphate. 
 
 1. OXYGEN COMPOUNDS OF SULPHUR. 
 
 Sulphur eombines in four proportions with oxygen. There are in addi- 
 tion nine compounds which contain hydrogen besides these two elements ; 
 they are dibasic acids. The only one which can be readily prepared is 
 sulphuric acid. The others are stable only in aqueous solution or in the 
 form of salts. In the table which follows the anhydrides are arranged 
 opposite to the acids into which they pass by the addition of water. 
 
 Oxides— Anhydrides. 
 
 Sulphur sesquioxide, SjOg 
 Sulphur dioxide, SOj 
 Sulphur trioxide, SO, 
 
 Sulphur heptoxide, SgO^ 
 
 Polythionic 
 
 acids 
 
 Both sulphur dioxide and trioxide 
 di-acid ox pyro-acid : 
 
 2502 + H2O = H2S2O5, Disulphurous acid. 
 
 2503 -j- HjO = H2S2O7, Disulphuric acid. 
 
 The latter alone is known in a free state. 
 
 Sulphur Dioxide, SO2, or sulphurous anhydride, is formed by burn- 
 ing sulphur or sulphides in the air : 
 
 S -I O2 = SO2. 
 
 I vol. I vol. 
 
 A little trioxide is always produced. 
 
 It may also be obtained by heating sulphur with the oxides of copper, 
 manganese and lead : 
 
 2CuO -f 2S = CujS + SOj. 
 
 Cuprous 
 
 sulphide. 
 
 Acids. 
 
 Thiosulphuric acid, HjSjOj 
 
 Hyposulphurous acid, HjSOg 
 
 Sulphurous acid, HjSO, 
 
 Sulphuric acid, HjSO^ 
 
 Persulphuric acid, H2S2O8 
 
 Dithionic acid, HgSjOg 
 
 (Hyposulphuric acid) 
 Trithionic acid, HjSgOg 
 
 Tetrathionic acid, H2S4O5 
 
 Pentathionic acid, H,ScOc. 
 
 yield a so-called anhydro-acid^ 
 
i84 
 
 INORGANIC CHEMISTRY. 
 
 It is most conveniently prepared for laboratories by heating concentrated 
 sulphuric acid (i part) with coi)per jiart) : 
 
 2H2SO4 + Cu = CuSO^ -h SO2 4 - 2H2O. 
 
 Copper 
 
 sulphate. 
 
 Usually a little cuprous sulphide separates : 
 
 5Cu + 411280, = CU2S + 3CUSO, -f 41130. 
 
 This is due to a far-reaching reduction. 
 
 Sulphuric acid is similarly decomposed (reduced) by heating it with 
 carbon : 
 
 2lbSO, -f C = 2SO2 4 - CO2 + 2II2O. 
 
 By this method we get a mixture of carbon and sulphur dioxides, which 
 are separated with difficulty. A more convenient method for prepar- 
 ing sulphur dioxide consists in allowing ordinary sulphuric acid to 
 act upon calcium sulphite, CaSOg. The latter is mixed with burnt 
 gypsum {yz part) and water, then moulded into cubes, which can be 
 introduced into a Kipp generator, as in the preparation of oxygen 
 (p. 81). Owing to its solubility in water, sulphur dioxide must be col- 
 lected over mercury. 
 
 Sulphur dioxide is a colorless gas, with a suffocating odor. One liter 
 of it weighs 2.8615 grams under normal conditions. Its specific gravity 
 equals 2.21 (air= i) or 64.06 (O2 = 32), corresponding to the molecular 
 formula SOj. It condenses at — 15°, or at ordinary temperatures under 
 a pressure of three atmospheres, to a colorless liquid, of specific gravity 
 1.43 at 0°, which crystallizes at — 76° and boils at — 8°. Its critical 
 temperature is 157°; its critical pressure 79 atmospheres. Upon evapo- 
 ration the liquid sulphur dioxide absorbs much heat ] being easily 
 accessible it is used in ice machines. If some of the liquid is poured 
 upon mercury in a clay crucible, and the evaporation accelerated by 
 blowing air upon it, the metal will solidify. Water dissolves 50 volumes 
 of sulphur dioxide gas with liberation of heat. The gas is again set free 
 upon application of heat. The solution shows all the chemical properties 
 of the free gas. 
 
 Sulphur dioxide has great affinity for oxygen. The gases combine 
 when dry, if their mixture be conducted over feebly heated platinum 
 sponge ; * sulphur trioxide results : 
 
 2SO2 + O., = 2SO3. 
 
 2VOls. IVOl. 
 
 Winkler’s method for producing sulphuric acid technically is based on 
 this reaction (compare p. 193). 
 
 In a(pieous solution the dioxide slowly absorbs oxygen from the air, and 
 becomes sulphuric acid : 
 
 S()2 + H20-f 0=11.280,. 
 
 * Instead of platinum s])onge, platinized asbestos may be applied ; this is obtained by 
 immersin}:^ asbestos in a platinic chloride solution, then in ammonium chloride, and after- 
 ward drying and igniting. 
 
SULPHUROUS ACID. 185 
 
 Aqueous sulphur dioxide is converted more rapidly into sulphuric acid 
 by the action of the halogens chlorine, bromine and iodine: 
 
 H,S03 + H^O -f Cl, = H,SO, + 2HCI. 
 
 Here the decomposition of a molecule of water is effected in conse- 
 quence of the affinity of the halogen for hydrogen and of sul})hurous acid 
 for oxygen. On adding sulphurous acid to a dark-colored iodine solution 
 the latter is decolorized. 
 
 Similarly, sulphurous anhydride and its solution withdraw oxygen from 
 many compounds rich in that element ; hence it deoxidizes strongly and 
 passes over into sulphuric acid. Thus chromic acid is reduced to oxide, 
 and the red solution of permanganic acid is decolorized with formation 
 of manganous salts. Many organic coloring substances, like those of 
 flowers, are decolorized by it.* This property is what leads to its appli- 
 cation in the bleaching of wools and silks, which are strongly attacked 
 by the ordinary chlorine bleaching agents (p. 53). 
 
 Again, the dioxide may be deoxidized by stronger reducing agents (it 
 acts with them as an oxidant); thus in the presence of water, sulphur is 
 separated from it by hydrogen sulphide : 
 
 SO2 + 2H2S = 2H2O + 3S. 
 
 If, however, both gases are perfectly dry or strongly diluted by other 
 neutral gases, the action is very slow. (See Hydrosulphurous and Penta- 
 thionic Acids.) 
 
 A mixture of equal volumes of sulphur dioxide and chlorine unites in direct sunlight 
 to sulphuryl chloride SO^Ch (p. 195). When sulphur dioxide acts upon warmed 
 phosphoric chloride, the products are phosphorus oxychloride, and the compound SOC’b : 
 
 SO2 + PCI5 = POCI3 -f SOCI2. 
 
 Chlorthionyl, SOC^, may be viewed as the chloride of sulphurous acid or as sul- 
 phur dioxide in which one atom of oxygen is replaced by two atoms of chlorine (p. 186). 
 Thionyl chloride may be made by the interaction of sulphur dichloride and sulphur tri- 
 oxide : 
 
 SO3 + SCI2 = SOCI2 + SO2 
 
 [see Michaelis, Ann. Chem. 274 (1894), 184]. It is a colorless liquid with a sharp 
 odor, and boils at 78°. Its specific gravity equals 1.67. Water decomposes it into 
 hydrogen chloride and sulphurous acid : 
 
 SOCI2 + H2O = SO2 + 2PICI. 
 
 Sulphurous Acid, H2SO3, is not known in free condition, but is 
 probably present in the aqueous solution of sulj^hur dioxide. On cool- 
 ing the concentrated solution to 0°, colorless cubical crystals separate, 
 containing probably six molecules of water (Geuther). If the aqueous 
 solution is allowed to stand for some time, especially in sunlight, sulphur 
 separates with the formation of sulphuric acid : 
 
 3SO2 -p 2ii20 = ri22SO^ -j- s. 
 
 *The acid forms colorless compounds with dyestuffs of the flowers. Dilute sulphuric 
 acid or heat breaks down these derivatives, i. e., the original colors reappear. 
 
i86 
 
 INORGANIC CHEMISTRY. 
 
 Sulphurous acid is dibasic and forms two series of salts; the jirimary 
 (KHSO3) and secondary (KjSOg). 
 
 The sulphites, with the exception of those of the alkalies, are insolu- 
 ble or dissolve with difficulty in water [see Seubert and Elten, Z, f. 
 anorg. Ch. 4 (1893), 44]. When sulphurous acid is separated out 
 from its salts by stronger acids it decomposes into its anhydride and water : 
 
 Na^SOg -f 2IICI = 2NaCl -f SO, + II^O. 
 
 The following is all that is known regarding the chemical structure of sulphurous acid 
 and its derivatives. Its anhydride and chloride have the formulas : 
 
 IV IV 
 
 O = S = O and O = 
 
 By water absorption or by decomposition with water the hydrate 
 
 IV 
 
 O = S< 
 
 OH 
 
 OH 
 
 results. This formula indicates that the hydrogen atoms or the metals which may replace 
 them are not directly combined with the sulphur, but are linked through oxygen. Or- 
 ganic compounds, esters of sulphurous acid, are known which undoubtedly are derived 
 from this symmetrical formula. The inorganic salts of the acid, however, very probably 
 contain one metal atom in direct union with sulphur, hence their basal acid must have 
 the formula 
 
 IV 
 
 O = S< 
 
 O-OH 
 
 H 
 
 O^vi OH, 
 
 accordingly as sulphur is regarded as quadrivalent or sexivalent. Organic derivatives are 
 also known of this unsymmetrical acid. From this it would appear that the anhydride 
 readily yields compounds which may be derived either from a symmetrical or an un- 
 symmetrical hydrate, SOgHj. See Sodium Sulphite and also Richter’s Organic Chemistry. 
 
 The 7 ’ino-che 7 nical Depo 7 -t 77 ie 7 it . — Sulphur dioxide is a very powerful exothermic com- 
 pound. 71. 1 Cal. are set free in its formation from rhombic sulphur and oxygen. When 
 it dissolves in much water there is an additional disengagement of 7.7 Cal., so that the 
 heat of formation of the hypothetical sulphurous acid in dilute aqueous solution (from 
 sulphur, oxygen and water) equals 78.8 Cal. : 
 
 (S,02) gas = 7 I-I ; (S02,Aq) = 7.7 ; (S,02,Aq) = 78.8. 
 
 In consequence of this great loss of energy the dioxide is a very stable compound ; it 
 is only at high temperatures that it sustains a partial separation into sulphur and oxygen. 
 For its behavior toward oxygen, see p. 187. 
 
 Hydrosulphurous Acid, H.2SO2 or H2S2O4. On adding zinc, iron, and some other 
 metals to the aqueous solution of sulphurous acid they dissolve without liberation of 
 hydrogen to yellow-colored liquids. Schonbein (1852) observed that such solutions pos- 
 sessed, when ajjplied to indigo, strong bleaching properties, and assumed that there was 
 in them a jjeculiar acid, the composition of which was first determined by Schiitzenberger 
 in 1869. The hydrosulphurous acid is formed there by the action of the hydrogen set 
 free by the zinc upon a second molecule of sulphurous acid : 
 
 H2SO3 + Zn = ZnSOg + H2 
 
 and 
 
 H2SO3 + H2 = H2SO2 + FI2O. 
 
 The aqueous solution of the acid has an orange-yellow color, reduces powerfully, bleaches 
 and soon decomposes with separation of sulphur and the formation of sulphurous acid. 
 
 'I'lie salts are more stable than the acid. The sodium salt is obtained by the action of 
 zinc filings upon a concentrated solution of primary sodium sulphite. It is used in dye- 
 
SULPHUR TRIOXIDE. 
 
 187 
 
 ing and cotton printing to bleach indigo. Its composition is not established with cer- 
 tainty ; it corresponds to either the formula NaHSOg or NujSjO^. The salt solutions 
 absorb oxygen very rapidly from the air and change to sulphites. 
 
 Sulphur Sesquioxide, SjOg, is obtained by the solution of flowers of sulphur in 
 liquid sulphuric anhydride ; it separates out in blue drops, which solidify to a mass resem- 
 bling malachite. It decomposes gradually, more rapidly on warming, into sulphur dioxide 
 and sulphur. It is very violently broken down by water, with formation of sulphur, 
 sulphurous, sulphuric, and polythionic acids. It dissolves with a blue color in fuming 
 sulphuric acid. 
 
 Sulphur Trioxide, SO3, or sulphuric anhydride, is produced, as pre- 
 viously described, by the union of sulphur dioxide and oxygen, aided by 
 heated platinum sponge. Platinized balls of white clay are used in tech- 
 nical operations. It is also formed when sulphur dioxide and air are con- 
 ducted over ignited oxide of iron, chromic oxide, or manganese oxide. 
 These oxides act like the platinum sponge, platinized asbestos or clay. 
 They are merely contact substances. It can also be made by heating 
 sodium or potassium pyrosulphate (p. 194) and anhydrous sulphates, e. g., 
 ferric sulphate : 
 
 FcjlSOJa = Fe^Oa + 3SO3 ; 
 
 and is most conveniently obtained by heating fuming (Nordhausen) sul- 
 phuric acid (p. 193); the escaping white fumes are condensed in a chilled 
 receiver. It may be obtained pure by repeated distillation, by fusion at 
 moderate temperature (20-30°), and then pouring it off from the re- 
 maining solid portions. It crystallizes in long, broad, transparent 
 needles, which melt at 14.8° to a very mobile liquid. At 16° its specific 
 gravity is 1.940. It distils at 46°. The vapor density agrees with the 
 formula SO3. The perfectly pure anhydride does not change on preser- 
 vation. If, however, by absorption of water it contains traces of sulphuric 
 acid, it soon becomes an asbestos-like mass, which does not melt until at 
 about 50°. This was formerly regarded as a peculiar form of sulphuric 
 anhydride. The pure anhydride can be readily obtained from it by 
 distillation [R. Weber; see also Rebs, Ann. Chem. 246 (1888), 379]. 
 
 Sulphuric oxide fumes strongly in the air, and attracts moisture with 
 avidity. When thrown on water it dissolves with hissing to form sul- 
 phuric acid (SO3 -f- H2O = H2SOJ. 
 
 When the vapors are led through heated tubes they are decomposed 
 into sulphur dioxide and oxygen. 
 
 Thermo- chemical Deportment . — When sulphur dioxide and oxygen combine to form 
 liquid sulphur trioxide 32. 1 Cal. are disengaged, so that its heat of formation from the 
 elements is 103.2 Cal. : 
 
 (S 0 ^, 0 ) liquid = 32.1, (8,03) liquid = 103.2, 
 
 inasmuch as the heat of formation of the dioxide 71. i Cal. (p. 186). This is another 
 contradiction of Berthelot’s princi{)le of the greatest evolution of heat. According to it 
 when sulphur burns in the air or in oxygen it should form not the dioxide but the trioxide, 
 because in the latter case there would occur the greater heat evolution. The fact is that 
 
i88 
 
 INORGANIC CHEMISTRY. 
 
 from 91. 1-95.5 percent, of the sulphur is converted into dioxide, and only 2. 9-2. 5 per 
 cent, into trioxide. In the presence of porous bodies (ferric oxide, etc.) the <)uantity 
 of trioxide reaches as much as 13 per cent. 
 
 On dissolving .sulphur trioxide in much water to form afjueous sulphuric acid, 39.2 
 Cal. are disengaged. The j)roduction, therefore, of the aqueous acid from sulphur, 
 oxygen, and water equals (including the heat of formation of SO3) 142.4 Cal. : 
 
 (S03,Aq) = 39.2 ; (S,03,Aq) = 142.4. 
 
 If we add to this the heat of formation of water (liquid), 68.3 Cal., the heat of f«^r 
 mation of sulphuric acid (II2SO4 — SO3 -(- HjO) from the elements in dilute aqueous 
 solution will be : 
 
 (Il2,S,04,Aq) = 210.7. 
 
 The heat of .solution of anhydrous sulphuric acid, in much water, equals 17.8 Cal. ; 
 hence the heat of formation of anhydrous sulphuric acid from its elements is 210.7 — 
 17.8 = 192.9 : 
 
 (Il2S04,Aq) = 17.8 ; (Il2,S,04) = 192.9. 
 
 Sulphur Heptoxide, was obtained by Berthelot on conducting a silent elec- 
 
 tric discharge of considerable intensity through a mixture of equal volumes of dry .sulphur 
 dioxide and oxygen. It separates in oily drops which .solidify at 0° to a crystalline mass. 
 This compound must be regarded as the anhydride of persulphuric acid. It decomposes 
 upon standing, immediately when heated, into oxygen and sulphur trioxide : 
 
 S2O7 = 2SO3 + O. 
 
 It fumes strongly in the air and with water, as with heat, decomposes into sulphuric 
 acid and oxygen : 
 
 ^2^7 "f" 2H2O 2H2SO4 -J- O. 
 
 Persulphuric Acid, H2S20g, corresponding to sulphuric heptoxide, has not been 
 obtained in a pure condition. A solution of it may be formed at the anode when sul- 
 phuric acid (cooled, 40 per cent.) is electrolyzed. Richarz thinks that it is produced in 
 this instance by the sulphuric acid breaking down into the ions H and HSO4 (p. 92), 
 the latter then combining to 1^28203. Inactive oxygen and ozone are formed simulta- 
 neously. A solution of the acid can be obtained by dissolving sulphur heptoxide in dilute 
 sulphuric acid, also by the addition of aqueous hydrogen peroxide to cooled, concentrated 
 sulphuric acid, when ozone is also produced. 
 
 The solution of persulphuric acid in sulphuric acid exhibits oxidation reactions similar 
 to those of hydrogen peroxide. It oxidizes ferrous to ferric sulphate, gradually liberates 
 iodiive from potassium iodide and decolorizes the blue solution of indigo-sulphuric acid ; 
 however, it is not capable of decolorizing a permanganate solution, neither does it oxi- 
 dize chromic to perchromic acid or affect titanic acid solutions (p. 103). 
 
 Hydrogen peroxide is produced if the electrolyzed sulphuric acid contains more than 
 60 per cent, of sulphuric acid ; this is caused by the breaking down of the persulphuric 
 acid which has been formed (p. loi). 
 
 d'he persulphates result in the action of strong bases upon the heptoxide, or, as shown 
 by Hugh Marshall and Berthelot, by the electrolysis of sulphate solutions [Elbs, Jr. 
 j)rakt. Ch. 48 (1893), 185, and Chem. Zeit. 1895, 1120]. They will be described under 
 the respective metals. Their dilute solutions show the following properties, character- 
 istic for persulphuric acid : when heated they give out ozone, and in the presence of 
 hydrochloric acid chlorine ; they precipitate the hydrated dioxide of manganese upon 
 the addition of a manganous salt ; in the presence of sulphuric acid they oxidize aniline 
 to aniline black. See p. 198 for the probable structure of persulphuric acid. 
 
SULPHURIC ACID. 
 
 189 
 
 SULPHURIC ACID. 
 
 H2SO4. 
 
 it is certain that this acid was known in the fifteenth century, probably 
 long before that time. It is the most important of all the acids; nearly 
 all of them can be prepared directly or indirectly by means of it. In 
 technical chemistry and in the arts it meets with an unusually exten- 
 sive and varied application, but it is especially valuable, in the Le Blanc 
 soda process, in the refining of petroleum and tar-oils, for the production 
 of aluminium sulphate, blasting material, dyes, artificial manures, etc., 
 etc. 
 
 Besides the reactions already mentioned, it is formed in the oxidation of 
 sulphur by nitric acid. It was obtained formerly by heating ferrous sul- 
 phate (FeSOJ, which breaks down into ferric oxide, sulphurous acid and 
 sulphuric anhydride : 
 
 2FeSO^ = FejOj SOj + SO3. 
 
 At present, however, it is almost exclusively manufactured in large quan- 
 tities, after the so-called English lead-chamber process. This method is 
 based upon the conversion of sulphur dioxide into sulphuric acid by means 
 of nitric acid. Sulphur, pyrite (FeSj), or other blendes are roasted in 
 ovens, and the disengaged sulphur dioxide immediately conducted, to- 
 gether with air, into a series of large leaden chambers in which it is fre- 
 quently brought in contact with nitric acid and steam. 
 
 The nitric acid gives up a portion of its oxygen to the sulphur dioxide 
 and thereby oxidizes it in the presence of water to sulphuric acid, which 
 collects on the floors of the lead chambers. The nitrogen oxides, NO, 
 NOj, and N2O3, arising simultaneously from the nitric acid, are capable 
 of transferring, in various ways, the oxygen of the air, which enters the 
 chambers, in the presence of water, to the sulphurous acid so that it 
 passes over into sulphuric acid. As the nitrogen oxides are being con- 
 stantly regenerated, a given quantity of nitric acid should, according to 
 theory, be capable of converting unlimited amounts of sulphur dioxide 
 into sulphuric acid if the water and oxygen are present in sufficient 
 quantity. Facts show that this is really not true, because a portion of 
 the nitric acid (1-2 parts by weight for 100 parts of sulphurous acid) is 
 reduced to nitrous oxide, nitrogen, and probably ammonia, — compounds 
 which do not participate in the oxidation of the sulphurous acid. Hence, 
 in the lead chambers the nitrogen oxides, particularly nitric oxide, play 
 the role of oxygen carriers. 
 
 The chemical changes occurring side by side in the lead chambers are influenced by 
 the quantities of the reacting substances and by the temperature ; consequently, they are 
 not the same for all chambers, indeed not the same for different parts of the same cham- 
 ber. The chemical and physical conditions are being constantly altered by the violent 
 mixing together of the gases, and it seems almost a fruitless task to attempt to establish a 
 definite theory for the sul])huric acid manufacture. The most important of these changes 
 appear to be the following : 
 
 In the presence of water, the nitric acid oxidizes the sulphur dioxide to sulphuric acid, 
 and the former is reduced to nitric oxide (NO) or nitrogen dioxide (NOj) : 
 
 3SO2 -f 2IINO3 -f 2II2O = 3 H.,S 0 , 1 2NO. 
 
190 
 
 INORGANIC CHEMISTRY. 
 
 The nitric oxide unites with the oxygen of the air (which entered the chambers simul- 
 taneously with the sulphur dioxide) to form nitrogen dioxide, which, in the presence of 
 water, converts a fresh portion of sulphur dioxide into sulphuric acid : 
 
 SO2 + 11,0 + NO, = HjSO* + NO. 
 
 Tlie regenerated nitric oxide is again subjected to the same transformations. Tn Tunge’s 
 ojunion the lead-chamber process is not an alternating reduction and oxidation of the oxides 
 of nitrogen, but rather a condensation of nitric oxide and nitrogen dioxide with sulphur 
 
 dioxide, oxygen, and water to nitrosylsulphuric acid, (P- 207) : 
 
 2SO, + 2NO O3 -t- 11,0 = 2S0,<0-,N0 
 
 The excess of water in the lead chamber immediately converts this product into sulphuric 
 acid and nitrogen trioxide or a mixture of nitric oxide and nitrogen dioxide : 
 
 2 SO,< 2 i- ■+ N, 0 .. 
 
 N,03 = NO + NO, ; N,03 II ,0 ^ 2IINO,. 
 
 In the anterior portion of the chamber system the nitrosylsulphuric acid breaks down 
 with the assistance of the sulphurous acid and with the formation of nitric oxide : 
 
 2S0,<0jj^° + so, + 211,0 = 3S0,(0H), + 2NO. 
 
 The regenerated nitric oxide and nitrogen dioxide act together with air and steam upon 
 new quantities of sulphur dioxide in the manner already indicated. According to Lunge, 
 nitrosylsulphuric acid is formed upon the very first action of the nitric acid introduced into 
 the lead chamber : 
 
 SO, NO, . OH = 
 
 and it then reacts in the manner indicated. [See Lunge, Z. f. anorg. Ch. 7 (1894), 
 212 ; R. Hasenclever, Ber. 29 (1896), iii, 2861 ; H. Ost, Lehrbuch des technischen 
 Chemie, 3 Aufl. (1898), and especially Jurisch, in Dammer’s Handbuch der chemischen 
 Technologic (1895) I, 163.] 
 
 In the lead-chamber process the active nitrogen oxides (NO and NO,) are carried along 
 and withdrawn from the action by means of the escaping nitrogen and excess of air. To 
 avoid any further loss of nitric acid by this means, the escaping brown gases are conducted 
 through the so-called Gay-Lussac tower. This is constructed of lead sheets, and filled 
 with pieces of coke, or, as these are apt to become coated with a mud of lead sulphate 
 which stops them up, with fire-brick or cylinders over which concentrated sulphuric acid 
 constantly trickles. The acid completely absorbs the nitrogen oxides K^O,, NO, and 
 NO, with formation of nitrosylsulphuric acid (see p. 208). The nitrogen oxides can be 
 regained from the acid — the so-called nitroso-acids — collected at the bottom of the tower, 
 and made useful in the production of sulphuric acid in the chambers. This is effected at 
 present in the so-called Glove)' to^veVy which is constructed of lead plates and fire-proof 
 bricks, and inserted between the sulphur ovens and lead chambers. In this the nitroso- 
 acid (diluted with the previously obtained chamber acid) is allowed to run over fire-brick, 
 while the hot gases of combustion from the sulphur ovens stream against it. This cools 
 the hot gases to the required temperature (70-80°), water evaporates from the chamber 
 acid, and, at the same time, the nitrogen oxides are set free (see above), and carried 
 into the lead chambers. Hence, the (Hover tower serves not only for complete utiliza- 
 tion of the nitrogen oxides, but also for the concentration of the chamber acid (to 82 
 j)er cent. I I,SOJ. 
 
 'I'he chamber j)rocess may be illustrated by the following experiment : A large glass 
 flask A (Fig. 64) replaces the lead chamber; in its neck are introduced, by means of a 
 cork, several glass tubes, which serve to introduce the various gases. In rr, sulphur di- 
 
SULPHURIC ACID. 
 
 I9I 
 
 oxide is generated by heating a mixture of sulplmric acid and mercury or copper turn- 
 ings. The flask b contains some dilute nitric acid and copper turnings, from wliich nitric 
 oxide (NO) is evolved. Water is boiled in c to get steam. Air enters through d while 
 the excess of gases escapes through e. By the meeting of nitric oxide (NO) with the 
 air, red fumes of nitrogen dioxide (NO2') arise, and these in presence of water change 
 the sulphur dioxide to sulphuric acid. The regenerated nitric oxide yields nitrogen di- 
 oxide with the oxygen of the air, and converts another portion of sulphur dioxide into 
 sulphuric acid. In time aqueous sulphuric acid collects upon the bottom of the vessel. 
 
 If, at first, only sulphur dioxide, nitric oxide and air enter without the steam, we get 
 
 (by aid of the moisture of the air) the compound SO2 
 
 { 
 
 O.NO 
 
 OH 
 
 (the so-called ni tro.su Iphuric 
 
 acid) which covers the walls of the vessel with a white crystalline sublimate. These 
 crystals, known as lead-chamber crystahy are also formed in the technical manufacture 
 
 Fig. 64. 
 
 of sulphuric acid, when an insufficient quantity of steam is conducted into the chambers. 
 Water decomposes them into sulphuric acid and nitrogen oxides. 
 
 The acid collecting in the chambers (chamber acid) possesses, when 
 the operation has been properly conducted, the specific gravity of 1.5 (50^ 
 according to Beaiime) ; it contains about 60 per cent, of sulphuric acid 
 and 40 ])er cent, of water.* For concentration the chamber acid is first 
 heated in lead jians until the specific gravity reaches 1.71 (60° Beaume ; 
 
 *The Beaum6 araeometer {apuLoq, thin) — hydrometer — used so much in technical work 
 has an empirical scale. For liquids heavier than water an instrument is used the zero 
 j)oint of which is obtained by placing it in pure water. This will give the highest point 
 of the scale. A second fixed point is obtained by dipping the instrument into a solution 
 of 15 parts of salt in 85 parts of water. The distance between the two points is divided 
 into fifteen equal parts or degrees. This division is also continued further downward. 
 
192 
 
 INORGANIC CHEMISTRY. 
 
 78 per cent. H^SOJ. The lead pans are strongly attacked by further 
 evaporation. However, in vessels of cast iron, an acid of specific gravity 
 1.8 (64.2° Beaume ; 86.9 per cent. HjSOj may be olitained. d'hese 
 are the crude sulphuric acids (^Acidum sulphuriciwi crudutn) of commerce. 
 They still contain arsenic, .selenium, iron and lead, d'he jiurified acid 
 obtained in various ways and diluted is again reduced in lead vessels to 60° 
 Beaume; further concentration to 65.5-66° Beaume (92-94 per cent. 
 HjSO^; specific gravity 1.83-1.837) is conducted in vessels of glass, 
 porcelain or platinum. Tlie most concentrated acid (98-98.5 percent. 
 
 always made in vessels of platinum, which in recent years have 
 been lined within with gold, because they are then less attacked. 
 
 By the distillation of the crude bhiglish acid an acpieous solution at first 
 distils over (one-third of the distillate), but at 338° we obtain almost pure 
 sulphuric acid {Acidum sulphuricufu purum or destillatuvi). This has the 
 specific gravity 1.854 at 0° or 1.842 at 12°, and contains about 1.5 per 
 cent, of water. On cooling this to — 35° white crystals separate, which 
 after repeated recrystallization melt at 10.5° ; this is the anhydrous acid, 
 HjSO^; also called suljihuric acid nionohydrate. The crystalline acid 
 is more readily obtained by cooling the 96-98 i)er cent, sulphuric acid 
 to 0° or — 10°, and then adding already formed crystals. This is the 
 manner in which the anhydrous acid is produced technically ; the crys- 
 tals are separated from the liquid hydrous acid by a centrifugal machine. 
 
 Pure anhydrous sulphuric acid, H2SO^ (called monohydrate), has the 
 specific gravity 1.8384 at 15° (water of 4° = i), and is, therefore, lighter 
 than slightly hydrous acid. When the anhydrous acid is heated, white 
 fumes of sulphur trioxide escape at 40° ; the liquid begins to boil at 200°, 
 and at 338° the acid, with 1.5 per cent, of water, again distils over. 
 
 From these data it is obvious that sulphuric acid, even at a gentle heat, sustains a par- 
 tial decomposition (dissociation) into sulphur trioxide and -water. The vapor density of 
 sulphuric acid has been found to be 72.0 (O.^ = 32) at 332° (near its boiling point). It 
 diminishes at higher temperatures, and is 49 at 416°, where it is constant. This 
 behavior is explained by the dissociation of the acid molecules, according to the equation : 
 
 H2SO, = SO, + H2O. 
 
 I vol. I vol. I vol. 
 
 Hence the dissociation of the acid is complete at 416°, while it is only about 34 per 
 cent, at 332° (p. 94). 
 
 Concentrated sulphuric acid is a thick, oily liquid. On cooling a sul- 
 phuric acid, containing about 15 per cent, of water, to 0°, large six-sided 
 prisms of the hydrate H2SO^-|-H20 (called sulphuric acid dihydrate) 
 separate; these melt at H-8°, and at 205° break down into the anhydrous 
 acid and water. The second hydrate, H2SO4 2H2O, has the specific 
 gravity 1.63, and yields water at 195°. The concentrated acid possesses 
 an extremely great affinity for water, and absorbs aqueous vajior energet- 
 ically, hence is applied in the drying of gases and in desiccators. It mixes 
 with water with the evolution of considerable heat, and, for this reason, it 
 is especially directed, in diluting the acid, to pour the latter in a thin 
 stream into the water, and not the reverse, as explosive phenomena occur. 
 In mixing suljihuric acid with water, a contraction of the mixture takes 
 
SULPHURIC ACID. 
 
 193 
 
 place; its maximum corresponds to the hydrate H2SO^-j-2H20 (/. e., 
 98 parts : 36 parts). 
 
 The existence of the hydrates of sulphuric acid is explained, as in the case of periodic 
 acid, by the assumption of hydroxyl groups. Later investigations confirm this idea : 
 
 H2S0^ + = S(OH)g, Hexahydroxyl sulphuric acid. 
 
 11280^ -(- 11.2^ — SO(OH)^, Tetrahydroxyl sulphuric acid. 
 
 11280^ . . = S02(0H)2, Normal sulphuric acid. 
 
 The tetrahydroxyl and the hexahydroxyl sulphuric acids yield only salts of the normal 
 dibasic acid, when they are acted upon by bases. 8alts corresponding to the hydrates are 
 not known, as was the case with periodic acid. 
 
 The affinity of sulphuric acid for water is so great that the former with- 
 draws the hydrogen and oxygen from many substances, with the produc- 
 tion of water. In addition to carbon, many organic compounds contain 
 hydrogen and oxygen in the proportion in which these elements yield 
 water. The withdrawal of hydrogen and oxygen from such substances 
 leaves the carbon. This explains the charring action of the acid upon 
 wood, sugar, and paper. When sulphuric acid acts upon alcohol (C2HgO), 
 ethylene, C2H^, results (p. 153). 
 
 By conducting sulphuric acid vapors over red-hot porous bodies, it is 
 decomposed into sulphur dioxide, water, and oxgen (p. 194) : 
 
 H28O, SOj + lIjO + O. 
 
 When heated with sulphur, phosphorus, carbon, and some metals (mer- 
 cury, copper), the acid is reduced to sulphur dioxide (p. 183). Nearly 
 all the metals are dissolved by it, forming salts; only lead, platinum, and 
 a few others are scarcely attacked at all. It is a very strong acid, and, 
 when heated, expels most other acids from their salts; upon this depends 
 its application in the manufacture of hydrochloric, nitric, and many other 
 acids, especially the organic acids. The barium salt (BaSOJ is character- 
 ized by its insolubility in water, acids, and alkalies; therefore, sulphuric 
 acid added to solutions of barium compounds produces a white pulveru- 
 lent precipitate, which serves to detect small quantities of the acid. 
 
 The structure of sulphuric acid and its anhydride can, assuming sul- 
 phur to be sexivalent, be expressed by the following formulas: 
 
 
 VI 
 
 8 O. 
 
 Pyrosulphuric or Disulphuric Acid, H2S20^. On withdrawing 
 one molecule of water from two molecules of the acid there results the 
 
 compound 1^2820^, whose formation and structure may be represented by 
 the following formula : 
 
 SO 
 
 SO 
 
 - H20 
 
 S02< 
 
 S02< 
 
 OH 
 
 O 
 
 OH. 
 
 As this acid contains two hydroxyl groui)s it is dibasic; yet its manner 
 of formation shows that it j)ossesses an anhydride character. Later, we 
 
 17 
 
94 
 
 INORGANIC CHKMISTRY. 
 
 will observe that almost all polybasic acids, like i)hosphoric acid, 
 PC)(0H)3, silicic acid, Si()(OII)2, and chromic acid, Cr()2(()n)2, arc 
 capable, by the condensation of several molecules and the elimination of 
 water, of forming like derivatives, which bear the name Poly- or Pyro- 
 acids. The pyro-acid, corresponding to sulphurous acid, is also known 
 in salts. 
 
 The disulphuric acid is contained in the so-called fuming:; or Nordhau- 
 sen sulphuric acid (Acidum sulphuricum fumans, oil of vitriol), which was 
 formerly obtained by heating dehydrated ferrous sulphate — green vitriol 
 (FeSO^). It is a thick, oily, strongly fuming li([uid, of specific gravity 
 1.86-1.9. When it is cooled, large colorless crystals of disulphuric acid 
 (H2S2O2) separate ; these melt at 35°. Heat breaks it down into sulphuric 
 acid and suli)hur trioxide, which volatilizes: 
 
 II2S2O7 = up(\ -f SO3. 
 
 Conversely, disulphuric acid may be obtained by dissolving sulphur 
 trioxide in sul})huric acid : 
 
 II2SO, + SO3 = II2S2O7. 
 
 The production of fuming sulphuric acid also depends on this, as it may 
 be regarded as a solution of sulphur trioxide or pyrosuli)huric acid in an 
 excess of sulphuric acid. 
 
 Technically, fuming sulphuric acid is obtained from pyrites (FeS2) — (at present only 
 in Bohemia). The decomposition of the pyrites in the air yields ferrous sulphate and 
 ferric sulphate, which can be dissolved out with water. The solution is evaporated, and 
 the residue roasted in a reverberatory furnace, whereby the ferrous salt is changed to 
 ferric salt. The latter is then distilled from earthen retorts, when sulphuric acid and sul- 
 phur trioxide pass over and are collected in the receivers : 
 
 Fe,(SO.)3 = Fe^O, + 3SO3. 
 
 The residue, consisting of red ferric oxide, finds application as colcothar {caput 7 nor- 
 tuum) in polishing and as a paint. 
 
 Solid, crystalline pyrosulphuric acid has been recently introduced into 
 the market as a substitute for the fuming liquid sulphuric acid. It is made 
 by conducting the theoretical amount of sulphur trioxide into concen- 
 trated sulphuric-acid (see above). Sulphur trioxide is prepared by two 
 distinct methods at present. 
 
 In Winkler’s method sulphuric acid of 66° Baume is first allowed to run into retorts 
 raised to a red heat. The mixed gases, sulphur dioxide, oxygen and water (p. 193 b 
 resulting from this action, are freed from steam by passing through a coke tower through 
 which trickles concentrated sulphuric acid. The dry mixture is then conducted over 
 heated jdatinized balls of white clay and the resulting sulphur trioxide is collected in 
 concentrated suljfiniric acid. 
 
 'I'lie more recent method of Wolters consists in ]>roducing sodium pyrosulphate by heat- 
 ing sodium sulphate with concentrated sulphuric acid : 
 
 Na2SC), -j- ILSOi -f II./X 
 
 Sodium Sodiiiin 
 
 sulphate. pyrosulphate. 
 
SULPHURIC ACID. 
 
 95 
 
 An intermediate product in this reaction is primary sodium sulphate — NallSO^ — 
 which upon the application of heat gives up water and passes into the pyrosulphate. By 
 the action of concentrated sulphuric acid upon the latter the anhydride is liberated and 
 distilled off in a vacuum : 
 
 Na^S.O, + H^SO^ = 2NaHS04 + SO3. 
 
 The residual sodium sulphate can again be converted into pyrosulphate. Sodium pyro- 
 sulphate decomposes at about 600° into sodium sulphate and sulphur trioxide, thus : 
 
 Na2S207 = Na2S04 -j- SO3. 
 
 Sulphuric Acid Chloranhydrides. — Under the name of halogen anhydrides we 
 understand the derivatives resulting from the replacement of hydroxyl in acids by 
 chlorine. Conversely, the chloranhydrides, by the action of water, pass into the corre- 
 sponding acids ; 
 
 so,<^} + 2H,0 = + 2HCI. 
 
 The ordinary method for the preparation of the chloranhydrides consists in permitting 
 phosphorus pentachloride to act on the acids. Sulphuric acid has two hydroxyl groups ; 
 therefore it can furnish two chloranhydrides. 
 
 Cl 
 
 The first, S02<^j^q, Sulphuryl Hydroxy-chloride, or Chlorstdphonic Acid, results 
 when one molecule of phosphorus pentachloride acts upon one molecule of sulphuric acid : 
 
 + PCI5 = SO,<g'j^ + FOCI, + HCl. 
 
 The resulting oxychloride of phosphorus acts upon two additional molecules of sul- 
 phuric acid, with formation of metaphosphoric, hydrochloric, and chlorsulphonic acids : 
 
 2S02(0H)2-f POCI3 = 2S02 <qj^ + HPO3 + HCl. 
 
 It is formed, too, by the direct union of sulphuric anhydride with hydrochloric acid : 
 SO3 -f HCl = SO3CIH. 
 
 The mo.st practical method for its formation consists in conducting chlorine gas through 
 sulphuric acid (15 parts), and gradually adding phosphorus trichloride (7 parts). Or, 
 hydrochloric acid gas is led into solid fuming sulphuric acid (H2S2O.J), as long as absorp- 
 tion occurs, and then the acid is distilled (Otto). 
 
 Chlorsulphonic acid is a colorless, strongly fuming liquid, of specific gravity 1.716 at 
 
 18°, and boils without decomposition at 152°. The salt S02 -<qj^ results from the 
 
 union of sulphur trioxide with potassium chloride. 
 
 The second chloranhydride, SO2CI2, or Sulphury! Chloride,* forms when phosphorus 
 pentachloride acts upon sulphuric anhydride ; and also by the direct union of sulphur di- 
 oxide with chlorine in sunlight : 
 
 SO2 -(- Clj = SO2CI2. 
 
 I vol. I vol. I vol. 
 
 * 'Hie group SO2 combined with 20H-groups in sulphuric acid, is known as Sul- 
 phuryl, 
 
96 
 
 INORCJANIC CHEMISTRY. 
 
 'I’he most convenient method for its formation is to pass equal volumes of sulphur dioxide 
 and chlorine gas into a cajxacious flask, containing some camphor, which causes the 
 union of the gases to form sulphuryl chloride. A colorless, suffocating, strongly fuming 
 liquid, of specific gravity 1.667 at 20° (referred to water at 4°), results. It boils at 
 69. 1°. Water decomposes it slowly at the ordinary temperature into sulphuric and hydro- 
 chloric acids. A little water will first change it to chlorsulphonic acid : 
 
 SOj<°[ + 11,0 = SO,<°'jj + IICI. 
 
 Its vapor density is normal at 130° ; at higher temperatures it breaks down into sulphur 
 dioxide and chlorine : 
 
 SO2CI2 = SO2 + CI3. 
 
 Pyrosulphuryl Chloride, S.^O^Clj, is the chloranhydride of pyrosulphuric acid. It 
 is obtained by several reactions, chiefly, however, by the action of phosjdiorus pentachlo- 
 ride or phosphorus pentoxide upon chlorsulphonic acid : 
 
 It is a thick liquid, fuming in the air; has a specific gravity of 1.819 at 18°, and 
 boils at 142°. At 210° it shows a normal density ; it is dissociated at higher tempera- 
 tures, corresponding to the equation : 
 
 S2O5CI2 SO3 + SO2 + CI2. 
 
 It dissolves gradually in water, without hissing, and decompo.ses into sulphuric acid and 
 hydrochloric acid ; at first, with a little water it yields chlorsulphonic acid. 
 
 AMIDO-DERIVATIVES OF SULPHURIC ACID. 
 
 The amide of an acid is produced when its hydroxyl is replaced by the univalent 
 group NH2. In the case of polybasic acids there are several amides as there are sev- 
 eral chlorides, depending upon whether the hydroxyl groups are partly or completely re- 
 placed by amido-groups. Sulphuric acid yields : 
 
 S02< 
 
 NH 
 
 OH 
 
 3 
 
 Amidosulphonic acid, 
 Sulphanilic acid 
 or Amidosulphuric acid. 
 
 SO 
 
 ^^2^NH2 
 
 Sulphamide, 
 Sulphuric acid amide. 
 
 They may be obtained by treating the corresponding chlorides with dry ammonia. 
 
 Siilphamide, formed by Traube on conducting ammonia into a solution of sulphuric 
 acid chloride : 
 
 SO2CI2 + 4NH3 = S 02 (NH 2)2 + 2 NH,C 1 , 
 
 consists of large, colorle.ss crystals melting at 8l°. It dissolves readily in water ; its solu- 
 tion is neutral in reaction and has a .slightly bitter taste. Alkalies and acids partly 
 .saponify it in acjueous solution, forming at first sulphamic acid, then sulphuric acid and 
 a\nmonia — a reaction corresponding to the decomposition of the chlorides of sulphuric 
 acid by water. Suli)hamide, like sulphuric acid, contains two hydrogen atoms replaceable 
 by metals. It yields a silver salt of the composition S02(NIIAg)2 [see W. Traube, Ber. 
 26 (1893), I, 607]. 
 
 In the action of ammonia upon sulphuryl chloride, sulphimide, (SOg) = Nil, is pro- 
 d iced together with suljdiamide. It is also formed when the latter is heated to 200-210°. 
 It acts like a monobasic acid ; its salts, e. g.^ 
 
 SO 2 N.K, S02N.Na, (S02N)2Ba -f 2 H 2 O, 
 
POLYTHIONIC ACIDS. 
 
 197 
 
 are readily soluble in water and crystallize well. Crystallized sulphiinide is not yet 
 known ; by the absorption of water it passes readily into sulphuric acid and ammonia 
 (\V. Traube). 
 
 In addition to the preceding there are other compounds derived either from ammonia 
 or from hydroxylamine in such manner that the hydrogen in them in direct union with 
 nitrogen is replaced by the univalent sulpho-group — SO 3 H. Two series result : 
 
 NH^. SO3H 
 
 Amidosulphuric acid (see p. 196). 
 
 NHISO^H)^ 
 
 Imidosulphonic acid, 
 Disulphamic acid. 
 
 N(S 03 H )3 
 
 Nitrolosulphonic acid, 
 Trisulphamic acid. 
 
 HO. NH. SO 3 H 
 
 Hydroxylamine sulphonic acid. 
 
 HO. N(S03H)2. 
 
 Hydroxylamine disulphonic acid. 
 
 The acids of the first series are obtained from ammonia and sulphur trioxide or chlorsul- 
 phonic acid ; also, just like the second, from sulphurous acid and nitrous acid or hydroxyl- 
 amine in aqueous solution [Raschig, Ann. Chem. 241 (1887), 161 ; also Chem. Central- 
 blatt, 1897, I, 10 ; compare p. 131]. 
 
 OH 
 
 Thiosulphuric Acid, H2S2O3 = ? generally known as hypo- 
 
 sulphurous acid, can be considered as sulphuric acid in which the oxygen 
 of an hydroxyl group is replaced by sulphur. It is not known in a free 
 condition, since as soon as it is liberated from its salts by stronger acids 
 it at once decomposes into sulphur dioxide, sulphur and water (com- 
 pare Ber. 22 (1889), 1686, 2703): 
 
 Na2S203 -f 2HCI = 2NaCl + SO2 -f S -f H2O. 
 
 Its salts, called hyposulphites, are of practical importance (compare Sodium 
 Hyposulphite). They are formed by the direct addition of sulphur to 
 sulphites : 
 
 Na 2 S 03 S = Na 2 S 203 ; 
 
 similar to the formation of sulphates by the addition of oxygen to the sul- 
 phites. 
 
 The formation of its sodium salt by the action of iodine upon a mixture of sodium sul- 
 phite and sodium sulphide, 
 
 NaSNa -f NaS03 . Na + I2 = NaS . SO.p^a -f 2NaI, 
 
 Sodium hyposulphite. 
 
 is taken as evidence of the chemical structure of the acid (Spring). However, this reac- 
 tion can take place in two phases : 
 
 I . Na2S + B 2NaI + S ; 2. S + Na2S03 = Na2S203. 
 
 Sodium amalgam resolves sodium hyposulphite into sodium sulphite and sodium sulphide. 
 
 POLYTHIONIC ACIDS. 
 
 By this name (from tzoXuc;, many, and Osiov, sulphur) are understood the 
 acids of sulphur, containing two or more atoms of the latter with six 
 
198 
 
 INORGANIC CIIEMISIRY. 
 
 oxygen atoms and two Iiydrogen atoms. Salts of the following are 
 known : 
 
 l^ithionic acid. 
 ll-^SgOg, 'rrithionic acid. 
 llgS^Og, d'etrathionic acid. 
 ll.^SjOg, rcMtathionic acid. 
 
 Following Blomstrand and Mendelejeff it maybe assumed that there are two univalent 
 \’I 
 
 groups — SOgll or SOg.OH — in the.se acids which in dilhionic acid are directly com- 
 bined, but in the other acid.s are joined through S, S.^, S3, with one another. 
 
 SO3II SO3II SO3II SO.II 
 
 SO3II 'SO3II 
 
 VI 
 
 The group SO3II is known as the sulpho-growY> •, it is also present in organic sulpho- 
 acids, and corresponds to the acid-forming carbon group, COOII, called carboxyl. From 
 this group are derived the acids just mentioned : 
 
 H.SOj.OH HO.SOg.OII ^<80 011 IIS.SOg.OH. 
 
 Sulphurous acid. Sulphuric acid. Pyrosulphuric acid. Thiosulphuric acid. 
 
 The formulas of the polythionic acids may also be derived from the formulas of the 
 hydrogen sulphides, HjS, IlgSg, and HgSg, if the hydrogen atoms in the.se be replaced 
 by the sulpho-group. If in the case of the ordinary hydrogen sulphide this replacement 
 extends merely to one hydrogen atom, thiosulphuric acid results ; the acids corresponding 
 to it are not known in the other two series : 
 
 c/H 
 
 e .SO3 . OH 
 
 CJ .SOg. OH 
 ^^SOg. OH 
 
 s<r^ 
 
 • • 
 
 C. ^SO., . OH 
 . OH 
 
 
 
 Q ^SOg . OH 
 ^ 3 <-S 0 ., .OIL 
 
 Similarly, sulphuric acid and pyrosulphuric acid are derived from water ; whereas 
 persulphuric acid must, according to Michaelis and Richarz, be referred to hydrogen 
 peroxide : 
 
 H 
 
 H 
 
 Q^SOgH 
 
 O.SO3H 
 
 ^^SOgH 
 
 
 Sulphuric acid. 
 
 Pyrosulphuric acid. 
 
 -H 
 
 H 
 
 0 <-S 03 H 
 
 Q ^SOgH 
 
 
 Unknown. 
 
 Persulphuric acid. 
 
 This is an example of the type idea mentioned on p. 1 72. Whether the formulas 
 deduced above actually repre.sent the true chemical structure of the polythionic acids is 
 at present very questionable [see Debus, Ann. Chem. 244 (1888), 181]. 
 
 Dithionic Acid, Il^S./Ig, is only known in aqueous solution. When concentrated 
 in vacuo or when heated it decompo.ses into sulphuric acid and sulphur dioxide. Its 
 manganc.se .salt (together with manganous sulphate) results from the action of sulphur 
 dioxide upon manganese dioxide (MnOj) suspended in water: 
 
 MnO., + 2SO., = MnS-Pg 
 
 and 
 
 MnOg }- SOg = MnSO^. 
 
OXYGEN DERIVATIVES OF SELENIUM AND TELLURIUM. I99 
 
 Barium hydroxide converts this into the barium salt, from which the free dithionic acid in 
 aqueous solution is obtained by means of sulphuric acid. 
 
 Trithionic Acid, HgSgOg, is not known in a free condition. Its salts are produced 
 when an aqueous solution of primary potassium sulphite is digested with flowers of 
 sulphur : 
 
 6KHSO3 -f 2S = 2K2S3O6 + K2S2O3 + 3 H, 0 . 
 
 Separated from its salts by other acids it decomposes rapidly into sulphuric acid, sulphur 
 dioxide, and sulphur. Its production by the action of iodine upon a mixture of sodium 
 sulphite and hyposulphite is especially interesting : 
 
 Na^SjOg -j- Na^SOj -|- I^ = NagSjOg -f 2NaI. 
 
 It is also formed from sulphur dichloride and sodium sulphite : 
 
 SCI3 + 2Na2S03 — Na^SgOg -}- 2NaCl (Spring). 
 
 The tri- and not a tetrathionate is produced when sulphur dichloride is used. 
 
 Tetrathionic Acid, H^S^Og. Its salts are produced when iodine acts upon solutions 
 of the hyposulphites : 
 
 2KS . SO3K 4- I2 = K,S,Og + 2KI. 
 
 Potassium tetrathionate. 
 
 This reaction is applied in volumetric analysis for the quantitative determination of 
 iodine, and such substances as separate iodine from potassium iodide (see Sodium Thio- 
 sulphate). 
 
 'fhe tetrathionic acid, separated from its salts by stronger acids, is very unstable, and 
 when its aqueous solution is concentrated it decomposes into sulphuric acid, sulphur 
 dioxide, and sulphur. An aqueous solution of the acid can be easily prepared by con- 
 ducting hydrogen sulphide into aqueous sulphurous acid : 
 
 4SO2 -|- 3^2^ — 2^2^ “f" 3 ^- 
 
 If the solution be saturated with bases, neutral and acid tetrathionates result, e, g.y 
 BaS^Og -f 2H2O and Zn(HS^Og)2. 
 
 Pentathionic Acid, H2S50g, is supposed to be produced together with tetrathionic 
 acid when hydrogen sulphide acts upon aqueous sulphurous acid : 
 
 5H2S + 5SO2 - H2SgOg + 4H2O + 5 S 
 
 (Wackenroder’s liquid ; the dry gases do not react). It was from this solution that Debus 
 first prepared pure, crystallized pentathionates. 
 
 All the polythionic acids are distinguished from sulphuric acid by the solubility of their 
 barium salts. The alkali salts of pentathionic acid decompose in aqueous solution in 
 such a manner that after some time the salts of the other polythionic acids are also present 
 in the solution. 
 
 2. OXYGEN DERIVATIVES OF SELENIUM AND TELLURIUM. 
 
 Se02 H 2 Se 03 
 
 Selenium dioxide. Selenious acid. 
 
 . . . IlgSeO,. 
 
 Selenic acid. 
 
 Selenium Dioxide, Se02, or selenious anhydride, is produced when 
 selenium burns in the air or in oxygen. It consists of long white needles, 
 which sublime at about 320° without fusing. It dissolves readily in 
 
200 
 
 INORGANIC CHEMISTRY. 
 
 water, forming selenious acid, H^SeO,. Tlie latter is also obtained by 
 dissolving the metal in concentrated nitric acid. When the solution is 
 evaporated it crystallizes in large, colorless jirisms, which decompose, on 
 heating, into the anhydride and water. Sulphurous oxide reduces seleni- 
 ous acid, with separation of free selenium : 
 
 HjSeOj -f 2SO2 + 11,0 = 2lI,SO, -f Se. 
 
 It is similarly reduced by potassium iodide in hydrochloric acid solu- 
 tion with the simultaneous sejiaration of iodine: 
 
 SeO, + 4KI + 4IICI = 4KCI + 2 ll ,0 4- Se -f 2I,. 
 
 Selenic Acid, H2SeO^, is obtained by conducting chlorine gas into 
 an acpieous solution of selenious acid : 
 
 H2Se03 -f II2O + CI2 = IljSeO^ + 2IICI. 
 
 The solution may be concentrated until it attains a specific gravity of 2.6 when it 
 becomes an oily liquid, similar to suljdiuric acid, and contains 95 per cent, of selenic acid, 
 n2Se04. If the solution be heated above 260°, the acid breaks down into selenium 
 dioxide, oxygen and water. It was from this solution that (.larneron and Macallan 
 obtained the pure acid as a crystalline mass, melting at 58°. It forms a hydrate with 
 water, H2SeO^ . H2O, melting at 25°. Its anhydride is not known. 
 
 The salts of selenic acid are known as selenates, those of selenious acid as selenites 
 [see Jahr. Chem, 1889, 388]. 
 
 The derivatives of tellurium are very similar to those of selenium. 
 
 Tellurium Dioxide, TeO,, results when tellurium is burned, and forms a white 
 crystalline mass, fusing at a red heat and subliming. It is almost insoluble in water. 
 
 Tellurous Acid, H2Te03, is produced when the metal is dissolved in concentrated 
 nitric acid. 'Water will precipitate it from such a solution in the form of a white amor- 
 phous powder. On warming, it readily decomposes into the dioxide and water. 
 
 Telluric Acid, H.2TeO^. Potassium tellurate is produced when tellurium or its di- 
 oxide is fused with saltpeter. Tellurium is added to an excess of dilute nitric acid in 
 which it dissolves with violent reaction. The resulting dioxide is oxidized by means of 
 a little more than the calculated quantity of chromic acid : 
 
 STeO, + 2Cr03 +3H2O = 3H2Te04 -f Cr203. 
 
 The greater portion of the telluric acid separates on evaporation. It is purified by pre- 
 cipitation with concentrated nitric acid from its aqueous solution [Z. f. anorg. Ch. X 
 (1895), 190], The acid crystallizes from its aqueous solution with two molecules of 
 water (Il2Te04 -f- 211,0) in isometric or triclinic forms, which at 100° lose all their 
 water and break down into a white powder of telluric acid, H2Te04. The latter is not 
 very soluble in water, and manifests a slight acid reaction. When carefully heated, 
 telluric acid breaks down into water, and the trioxide, TeOg, which is a yellow mass in- 
 soluble in water, and by further application of heat decomposes into the dioxide and 
 oxygen. 
 
 3. OXYGEN DERIVATIVES OF THE ELEMENTS OF THE 
 NITROGEN GROUP. 
 
 'Phc halogens combine with one atom of hydrogen to form halogen 
 hydrides; their oxygen acids also containing one atom of hydrogen. 
 
OXYGEN DERIVATIVES OF NITROGEN. 
 
 201 
 
 The elements of the sulphur group contain two atoms of hydrogen in 
 their hydrogen derivatives and oxygen acids. Accordingly, we find that 
 the elements of the nitrogen group combine with three atoms of hydro- 
 gen, and form acids which also contain three atoms of the same element: 
 
 HCl H^S PHg 
 
 HCIO3 H2SO3 P0,H3 
 
 HCIO, H^SO^ P 0 ,H 3 . 
 
 The acids of the nitrogen group containing three atoms of hydrogen, 
 designated normal or ortho-acids {opdoq, correct, real, true), as HgPO^, 
 H3 AsO^, H3ASO3, can yield monobasic acids by the removal of one mole- 
 cule of water. Such derivatives, having one atom of hydrogen, are called 
 jneta-acids (^era, a word which here expresses a change in condition) : 
 
 H3PO, 
 
 Orthophosphoric acid, 
 Na 3 As 03 
 
 Sodium orthoarseniate. 
 
 HPO3 
 
 Metaphosphoric acid. 
 
 NaAsO.^. 
 
 Sodium meta-arseniate. 
 
 These meta-acids of phosphorus pass into the ortho-acids by the absorp- 
 tion of water. The ortho-acids of nitrogen, on the other hand, are less 
 stable and only exist in some salts. The ordinary acids and salts of 
 nitrogen belong to the meta-series and contain one atom of hydrogen 
 (or metal) ; 
 
 (H3NOJ HNO3 
 
 Orthonitric acid. Ord. Nitric acid. 
 
 (H3NO3) 
 Orthonitrous acid. 
 
 HNO2. 
 
 Ord. Nitrous acid. 
 
 The further exit of water produces the true anhydrides (p. 183). 
 
 1 . 
 
 OXYGEN DERIVATIVES OF NITROGEN. 
 
 N205 
 
 
 
 HNO3 
 
 Nitrogen pentoxide. 
 
 N.2O4 and NOj — 
 Nitrogen Dioxide, 
 
 tetroxide. 
 
 
 Nitric acid. 
 
 
 XT- 
 
 JNitrogen tnoxide. 
 
 NO 
 
 Nitric oxide. 
 
 > HNO2 
 Nitrous acid. 
 
 N,0 ^ 
 
 Nitrous oxide. 
 
 > (HNO).,. 
 
 Hyponitrous acid. 
 
 The following formulas express the structure of these compounds: 
 
 III III 
 N=N 
 V 
 o 
 
 Nitrous oxide. 
 
 Ill III 
 
 0 -=N N^O 
 
 V 
 
 o 
 
 Nitrogen 
 
 trioxide. 
 
 Ill III III 
 
 TIO-N=N-OTI OrrN-OII 
 
 Hyponitrous acid. Nitrous acid. 
 
 Ill V 
 0 =N 
 
 V 
 
 o 
 
 Nitric-nitrous 
 
 anhydride 
 
 (tetroxide). 
 
 V V 
 
 V 
 
 o 
 
 Nitrogen 
 
 pentoxide. 
 
 O-N-OU. 
 Nitric acid. 
 
202 
 
 INORGANIC CHEMISTRY. 
 
 Tlie salts of nitric acid arc called ^titrates ; those of nitrous acid, 
 nitrites. 
 
 NITRIC ACID. 
 
 IINO3. 
 
 This acid occurs in nature only in the form of salts, — potassium, 
 sodium, and calcium nitrates (comj)are these), — which have resulted from 
 the decay of nitrogenous organic substances in the j^resence of strong 
 bases (the alkalies) and Bacillus nitrificans. It is sometimes present in 
 the air as ammonium salt. The free acid is formed in very slight quan- 
 tity by conducting electric sparks through moist air; most readily with 
 a mixture of 5 volumes of air and 6 volumes of oxygen, which may 
 obtain technical value if the electricity and oxygen can be prepared 
 cheaply enough. At present Chile saltpeter is the j)rincii)al source of 
 nitric acid. 
 
 To prepare nitric acid heat potassium or sodium nitrate in a retort with 
 sulphuric acid, when the nitric acid will distil over and acid sodium or 
 potassium sulphate remain : 
 
 NaNOa -f HjSO^ = IINaSO^ + IINO3. 
 
 If one molecule of sulphuric acid be used with two molecules of the 
 nitrate, the resulting acid sulphate will at higher temperatures act upon 
 the second molecule of nitrate in the sense of the equation : 
 
 NallSO, + NaNOa = HNO3 -f Na^SO,. 
 
 The temperature requisite to complete the reaction is, however, so high 
 that a portion of the nitric acid will be decomposed. 
 
 Perfectly anhydrous nitric acid has not yet been obtained. The most 
 concentrated acid (99.8 per cent. HNO3) is a colorless liquid of specific 
 gravity 1.56 at 0° ; it fumes in the air, and at — 47° solidifies to a crys- 
 talline mass. At ordinary temperatures it undergoes a partial decomposi- 
 tion (similar to sulphuric acid) into water, oxygen, and nitrogen dioxide, 
 NO2, which dissolves in the acid, with a yellowish-brown color; the color- 
 less acid therefore becomes colored upon standing, and in sunlight soon 
 turns yellow. At 86° the acid commences boiling and sustains a partial 
 decomposition ; the first portions are colored yellow by the dissolved 
 nitrogen dioxide, but subsequently, aqueous acid distils over. Nitric 
 acid is completely decomposed into nitrogen dioxide, oxygen, and water, 
 when its vapors are heated to about 
 
 2IINO3 = 2NO2 + H3O -f O. 
 
 The acid mixes in all proportions with water. Upon distilling the 
 dilute aqueous solution, only pure water passes over at first ; the boiling 
 temperature gradually rises, and between 120° and 121° a solution goes 
 over, which contains 68 ])er cent, of HNO3, and has a s])ecific gravity of 
 1. 414 at 15°. Tliis is the ordinary concentrated nitric acid of trade. 
 When this is distilled with 5 parts of sul})huric acid, an almost anhydrous 
 acid is obtained, which may be freed from nitrogen dioxide, contained in 
 it, by conducting a stream of air through it. 
 
NITRIC ACID. . 203 
 
 The liquid boiling at 121°, however, can be regarded as a mixture of the trihydrate, 
 N0(0H)3, and pentahydrate, N(OH)5 (pp. 193, 201). 
 
 Nitric acid is a very powerful acid, oxidizing or dissolving almost all 
 metals (gold and platinum excepted). Nearly all the metalloids, like 
 iodine sulphur, phosphorus, and carbon, are converted by it into their 
 corresponding acids. It acts as a very strong oxidizing agent, destroy- 
 ing organic coloring substances, and decolorizes a solution of indigo very 
 readily. In so doing the nitric acid itself is deoxidized to the lower 
 oxidation products of nitrogen (NO and NOj). Some substances even 
 reduce the acid to ammonia. Thus, for example, if zinc be brought into 
 dilute nitric acid (5-6 per cent.) the metal will be dissolved without the 
 liberation of hydrogen. The latter, in stain nascendiy acts at once upon 
 the excess of acid and reduces it to ammonia, which forms an ammonium 
 salt with the acid ; hence, in solution, we have ammonium nitrate in 
 addition to the zinc nitrate : 
 
 2HNO3 + Zn = Zn(N03)2 + H, 
 
 and 
 
 2HNO3 + 4H2= NH^N03 + 3H2O. 
 
 If the aqueous nitric acid be less dilute (containing more than 10 per 
 cent, of HNO3) it will be reduced by zinc and other metals, not to am- 
 monia, but to the nitrogen oxides, N2O, NjOg, and N20^. The more 
 concentrated the acid, the higher will the oxides be. 
 
 The reduction of nitric acid to ammonia by nascent hydrogen occurs 
 more easily in alkaline solution. If a solution of nitrates, made alkaline 
 with sodium or potassium hydroxide, be heated with zinc or aluminium 
 filings or iron powder, then all the nitrogen of the nitric acid will be 
 converted into ammonia : 
 
 HNO3 + 4H2 = NH3 + 3H2O. 
 
 Hydroxylamine (p. 130) and ammonia are produced when nitric acid 
 acts on tin. 
 
 Nitric acid usually forms salts of the form MeNOg, with one equivalent 
 of the metals; these are called nitrates, and are all readily soluble in water. 
 
 Red Fuming Nitric Acid {Acidum nitricuni fumans') is the name given 
 a nitric acid containing much nitrogen dioxide in solution. It is 
 obtained by the distillation of two molecules of saltpeter with one mole- 
 cule of sulphuric acid (p. 202), or better, by the distillation of com- 
 mercial concentrated nitric acid with concentrated sulphuric acid. It 
 generally has the specific gravity of i. 5-1. 54, and possesses greater 
 oxidizing power than the colorless nitric acid. 
 
 A mixture of i volume of nitric acid and 3 volumes of concentrated 
 hydrochloric acid is known as Aqua regia, as it is able to dissolve gold 
 and platinum, which neither of the acids alone is capable of doing. The 
 powerful oxidizing action of ihe mixture is due to the presence of free 
 chlorine and the chlorine derivative NOCl, which may be considered 
 the chloranhydride of nitrous acid : 
 
 3IICI + IINO3 = 2II2O + NOCl -f- CI3. 
 
204 
 
 INORGANIC CHEMISTRY. 
 
 Nitrogen Pentoxide, N.^( >5, nitric anliydi ide, is produced hy carefully heating phos- 
 phoric anhydride with nitric acid : 
 
 211NO3 -f I’A - N/), -I 2 iiro 3 ; 
 
 further, together with oxygen, on conducting chlorine over silver nitrate : 
 
 2Ago. NO, + 2CI = + o. 
 
 It forms colorless, rhombic prisms, melting at 30° and boiling with partial decomposi- 
 tion at 47°. It is very unstable, decomposing readily into nitrogen tetroxide and oxygen, 
 and sometimes exploding spontaneously. It yields nitric acid with water and evolves 
 much heat by the union : 
 
 j;5g>>o + ii,o = 2NO,.on. 
 
 Nitroxyl Chloride, NO.2CI, the chloranhydride of nitric acid, results from the union 
 of nitrogen dioxide with chlorine, and according to the ordinary method of forming chlor- 
 anhydrides (see p. 195), by the action of phosphorus pentachloride or oxychloride upon 
 nitric acid, or better, its silver salt : 
 
 3NO, . OAg + POCI3 = 3NO3CI + P0(0Ag)3. 
 
 Silver nitrate. Silver phosphate. 
 
 It is said to be a heavy, yellow liquid, boiling at -1-5°. Water immediately decomposes 
 it into hydrochloric and nitric acids. However, Geuther’s experiments make the exist- 
 ence of this acid rather doubtful [Ann. Chem. 245 (1888), 98]. 
 
 Nitrosyl Chloride, NOCl, is produced by the union of nitric oxide (2 volumes) with 
 chlorine (l volume): 2NO CI2 = 2NOCI ; when phosphoric chloride, PCI5, is 
 allowed to act upon liquid nitrogen tetroxide, N2O4, and by heating lead-chamber crystals 
 (p. 208) wdth sodium chloride to 80-90° : 
 
 + NaCl = + NOCl. 
 
 The reddish -yellow vapors that escape, if cooled to — 20°, condense to a red liquid of 
 specific gravity 1. 416 at — 12°, and boils at -|-2°. It forms nitrous and hydrochloric 
 acids with water : 
 
 NOCl -f H2O = HNO2 + HCl. 
 
 It may, therefore be regarded as the chloranhydride of nitrous acid — NO . OH. 
 
 O. O. 
 
 Nitranaide, NO2 . NH2 = '^N - NHj or ^N = NH, the amide of nitric acid, 
 
 O^ HO^ 
 
 was di.scovered by Thiele and Lachmannon decomposing pota.ssium nitrocarbaminate (see 
 Organic Cliemistry) with sulphuric acid. It is also produced upon adding nitric acid to a 
 .solution of potassium imidosulphonate (p. 196) in concentrated sulphuric acid : 
 
 NH(S03H)2 + NO2. OH NO2. NH2 + H2S2O,. 
 
 Nitramide crystallizes from ligroine, in which it dissolves with difficulty, in brilliant white 
 plates. 1 1 dissolves readily in ether, alcohol and water. It melts at about 73° 'vith imme- 
 diate decomposition into nitrous oxide and water : 
 
 NO2. NH2 = N2O 4 IP/). 
 
 It volatilizes partly at the ordinary temperature. It is especially sensitive to alkalies; in 
 
NITROGEN TRIOXIDE. 
 
 205 
 
 contact with sodium hydroxide it explodes with fire phenomena. Its aqueous solution 
 reacts strongly acid. It also yields a readily decomposable mercury salt : 
 
 NO^NHg. 
 
 When nitramide is reduced hydrazine results : 
 
 NO2NH2 + 3H2 = NHj . NHj + 2 H, 0 . 
 
 Nitramide is isomeric with hyponitrous acid : HO-N=N-OH. [See Ann. Chem. 288 
 (1895), 267; 296 (1897), 95. J 
 
 Nitrogen Trioxide, NjOg, nitrous anhydride, is formed by the 
 direct union of nitrogic oxide (4 volumes) with oxygen (i volume) at 
 —18°: 
 
 4NO + 02 = 2N2O3 ; 
 
 4 vols. I vol. 2 vols. 
 
 by mixing liquid nitrogen tetroxide, N20^, with a little cold water : 
 
 2 JJgqo +hp = noj.o + 2N0,.0H; 
 
 by the introduction of nitric oxide into liquid nitrogen tetroxide 
 below — 21° : 
 
 N2O4 + 2NO = 2N2O3 ; 
 
 and by conducting nitric oxide into cold anhydrous nitric acid : 
 
 2HNO3 + 4NO = 3N2O3 + HgO. 
 
 It is most easily obtained by the action of nitric acid upon arsenious 
 oxide, AS2O3. Nitrogen tetroxide is simultaneously produced, from which 
 it is readily separated by fractional distillation and condensation. 
 
 Nitrogen trioxide condenses at — 21° to a dark-blue liquid of specific 
 gravity 1.449 begins to boil at 3.5°. It decomposes when 
 
 distilled ; its vapors consist of a mixture of the tetroxide and nitric 
 oxide (N2O4 2NO) and at more elevated temperatures of nitrogen di- 
 oxide and nitric oxide (2NO2 -f- 2NO). Upon cooling they reunite to 
 the liquid nitrous anhydride. The latter is therefore only known in the 
 liquid condition. Its dissociation begins at — 21° [Geuther, Ann. 
 Chem. 245 (1888), 96; see also Lunge and Porschnew, Z. f. anorg. 
 Chem. 7 (1894), 209]. 
 
 The trioxide mixed with a little cold water probably forms nitrous 
 acid (N2O3 -f- H2O = 2HNO2) ; more water, aided by heat, decomposes 
 it into nitric acid and nitric oxide gas : 
 
 3HNO2 = HNO3 + 2NO + H2O. 
 
 Nitrous Acid, HNO2, is not known in a free state. Its salts (the 
 7 iitrites) are obtained by igniting the nitrates: 
 
 KNO^ = KNOj + O. 
 
2o6 
 
 INORGANIC CHEMISTRY. 
 
 The withdrawal of oxygen is rendered easier if oxidizalile metals, 
 lead, be added to the fusion (see Potassium and Sodium Nitrite). Traces 
 of nitrites are found in the air and in many waters. 
 
 On adding sulphuric acid to the nitrites, brown vapors are disengaged ; 
 these consist of nitrogen dioxide and nitric oxide. It may be that the 
 nitrous acid, at first liberated, is broken down into water and the trioxide, 
 which, as we have seen above, gradually decomposes into nitrogen dioxide, 
 NO2, and nitric oxide, NO. Similar reddish-brown vapors are obtained 
 if nitric acid be permitted to act upon starch or arsenious oxide (As.^O.,). 
 
 According to lAinge, if we einjdoy nitric acid of specific gravity 1.30- 1.35, nitrogen 
 trioxide is produced almost exclusively, whereas in using the concentrated acid (1.4-1. 5) 
 we get a mixture rich in dioxide, and if the acid be dilute the chief product is nitric 
 oxide, NO, with a little nitrogen dioxide, NOj. 
 
 The nitrous acid which has separated out in the solution and its decom- 
 position products — N()2 and NO — are strong oxidizers, setting iodine 
 free from the soluble iodides. In other cases, however, they exhibit a 
 reducing action ; thus, e. g., the acidified red solution of jiotassium jier- 
 manganate is decolorized by the addition of nitrites, and nitrous becomes 
 nitric acid. In very dilute aqueous solution, the action proceeds ac- 
 cording to the following equation : 
 
 5HNO2 + 2KMnO^ + 311280^ == 5HNO3 4- K2S0^ + 2MnS(\ 3H2O. 
 
 This reaction serves for the quantitative determination of free nitrous 
 acid, as well as for its salts (p. 208). 
 
 Nitrogen Tetroxide, N20^, or nitrogen dioxide^ NO 2 (formerly called 
 hyponitric acid'), only exists at low temperatures; when heated, it suffers 
 a gradual decomposition into the simpler molecules NO2 which recom- 
 bine to N20^ upon cooling. We here meet the interesting case of disso- 
 ciation, occurring even at the ordinary temperature. The tetroxide, 
 N20^, is colorless, while the dioxide, NO2, is colored red-brown ; it appears, 
 therefore, that the color gradually becomes darker as the temperature 
 rises, and that it corresponds to the increasing dissociation of the com- 
 plex molecules N20^. The same was observed with sulphur tetrachloride 
 (p. no) and nitrogen trioxide (p. 205). 
 
 At ordinary temperatures, nitrogen tetroxide is a liquid of specific 
 gravity 1.49. When cooled to — 20° it solidifies to a colorless crystal- 
 line mass, melting at — 12°. In consequence of the dissociation which 
 begins at 0°, the liquid, at first colorless, becomes yellow, and the in- 
 tensity in color grows with rising tem])erature. The liquid begins to 
 boil at about 22°, and is converted into a yellowish-brown vapor which 
 becomes dark as the temiierature is increased. 
 
 'flic theoretical vapor density of nitrogen tetroxide, N20^ ecjuals 92.08, while that of 
 the dioxide, N()2, ecpials 46.04. 'i'lie exj)erimental vaj)or density has been found to equal 
 76 at the point at which the licpiid compound boils (26°) ; it may l)e calculated from this 
 that, at this t(nn])cniture, 34.4 per cent, of the tetroxide molecules are decomposed into 
 dioxide molecules. IJeiice we conclude that the dissociation of the compound com- 
 
NITROGEN TETROXIDE. 
 
 207 
 
 mences even in the liquid state ; this is confirmed by the yellow coloration appearing at 0°. 
 With rising temperature the density of the vapor steadily diminishes, becomes constant 
 finally at 150° and equals 46. Then all the molecules (N2O4) are decomposed into the 
 simpler molecules NO2 ; and the dark coloration of the vapors attains its maximum. 
 
 Nitrogen tetroxide is produced by conducting electric sparks through 
 a mixture of dry nitrogen and oxygen ; it is also formed by the union of 
 two volumes of nitric oxide with one volume of oxygen : 
 
 2NO + O2 = N2O,. 
 
 2 vols. I vol. I vol. 
 
 We can get it more conveniently by heating dry lead nitrate, which 
 decomposes according to the following equation : 
 
 Pb(N 03)2 = PbO + O + 2NO2. 
 
 The escaping dioxide condenses in the cooled receiver to liquid 
 tetroxide, N20^. 
 
 Nitrogen tetroxide varies in its behavior with water, according to the 
 temperature of the latter. We saw that by the action of a little cold 
 water, it was decomposed into nitrogen trioxide and nitric acid (p. 205). 
 With an excess of cold water, and also with an aqueous solution of alkalies, 
 it yields nitric and nitrous acids, that is, their salts: 
 
 + ^2^ = NO2 . OH -L NO . OH. 
 
 Both reactions plainly indicate that the liquid tetroxide represents the 
 mixed anhydride of nitric and nitrous acids (p. 201) ; similarly, the com- 
 pound Cl20^ constitutes the mixed anhydride of chloric and chlorous 
 acids (p. 1 1 7). AVarm water converts the tetroxide into nitric acid and 
 nitric oxide — because under these conditions the nitrous acid decom- 
 poses into nitric acid and nitric oxide (see p. 205). The dioxide behaves 
 similarly : 
 
 3NO2 -f H2O = 2HNO3 + NO. 
 
 The tetroxide and dioxide possess strong oxidizing properties; many 
 substances burn in their vapors; iodine is set free from the soluble 
 metallic iodides by them. 
 
 Many metals, just after the reduction of their oxides in hydrogen, absorb large quanti- 
 ties of nitrogen dioxide at lower temperatures, e. g., copper takes up 1000 times its 
 volume. “ Nitro-metals ” are produced in this way. “ Nitro-copper,” CiqNOg, is an 
 amorphous substance, which is decomposed by water into nitric oxide, copper, copper 
 nitrite and nitrate [Ber. 26 (1893), iv, 361]. 
 
 Nitrosylsulphuric Acid, SO^NH — S02<Qf|^^^, termed nitrosul- 
 phonic acid, is an intermediate product in the manufacture of commercial 
 sulphuric acid (see p. 190), and is quite important for the determina- 
 tion of the nitrogen oxides. It is produced by conducting nitrogen 
 tetroxide into concentrated sulphuric acid : 
 
 b<^h<CoH ~ ' -f- NO2 . OH. 
 
2o8 
 
 INORGANIC CHEMISTRY. 
 
 Nitrogen monoxide, NO, is not absorl)cd by pure suli)huric acid, but 
 will be if the same contains nitric acid : 
 
 3 H,S 0 . + HNO 3 + 2 NO = 3 S 03 <J]-yo + 211 , 0 . 
 
 Therefore it results upon conducting nitrous acid vapors into concentrated 
 sulphuric acid : 
 
 2S0,<0[{ + NO, + NO = 2S0,<^„^” + 11,0. 
 
 Further, the nitrosylsulphuric acid results from the combined action of 
 sulphur dioxide, nitrogen tetroxide, oxygen, and a little water: 
 
 2 SO, + N,0. + 0 + 11,0 = 2 S 0 ,<^',^*^. 
 
 It is obtained most readily by conducting suliihur dioxide into strongly 
 cooled anhydrous nitric acid : 
 
 so, + UNO, = 
 
 there results a thick magma, which may be dried upon porous earthen 
 plates under the desiccator. 
 
 Nitrosylsulphuric acid forms a leafy or granular crystalline, white 
 mass (lead-chamber crystals, p. 191), which melts at about 73° and decom- 
 poses into its anhydride, sulphuric acid, and nitrogen oxides. It deli- 
 quesces in moist air, and yields sulphuric and nitrous acids with water : 
 
 + ^=0 = S0=<0H + NO. OH ; 
 
 the nitrous acid decomposes further into nitric acid and nitric oxide 
 (p. 205). 
 
 Nitrosylsulphuric acid dissolves in concentrated sulphuric acid without any change ; 
 the solution, called nitroso-add, is also produced in the sulphuric acid manufacture in 
 the Gay-Lussac tower (p. 190), is very stable and may be distilled without decomposi- 
 tion. When diluted with water it remains unaltered at first, but when the specific 
 gravity of the solution reaches 1.55-1.50 (51-48° Be. ), then all the nitrogen oxides escape, 
 especially on warming. When the nitroso-acid is poured into a large quantity of water, 
 the nitrososulphuric acid breaks down into sulphuric and nitrous acids, and the latter in 
 part into nitric acid and nitric oxide. 
 
 All the nitrogen oxides and acids are separated as nitric oxide (NO) on shaking the 
 nitro.so-acid with mercury — a procedure serving equally well for estimating the amount 
 of nitroso-acid by means of the nitrometer. Sulphur dioxide acts similarly and upon 
 this depends the denitrating action of the Glover tower (see ]:)p. 189, 190). 
 
 The anhydride of nitrosylsulphuric acid, S^N.^Oy = O | C) N( )’ P^'^^biced upon 
 
 heating the acid (together with sulphuric acid, nitric oxide, and nitrogen dioxide) beyond 
 its point of fusion. It is obtained pure by .saturating sulphur trioxide with nitric oxide : 
 
 3S()., -I 2N() = ()(S(),,.O.NO),, -t- SO.,. 
 
 1 1 is a crystalline;, colorless mass, molting at 217°, and boiling without decomposition at 
 about 360°. Much water decomposes it, the same as nitrosylsulphuric acid. 
 
NITRIC OXIDE. 
 
 209 
 
 The chloranhydride of nitrosylsulphuric acid, SNO^Cl = S02 <Cq‘ is formed by the 
 
 union of sulphur trioxide with nitrosyl chloride : 
 
 SO3 + NOCl = SO^C^O . NO). 
 
 It forms white leaflets, is decomposed by heat into its components, and with water breaks 
 down into sulphuric, hydrochloric, and nitrous acids. 
 
 Nitric Oxide, NO. When different metals are dissolved in some- 
 what dilute nitric acid this oxide is formed. It is conveniently ob- 
 tained by pouring dilute nitric acid (specific gravity 1.2) upon copper 
 filings : 
 
 3Cu -f 8HNO3 = 3 Cu(N 03)2 + 4H2O -f 2NO. 
 
 A better procedure consists in pouring a saturated solution of salt- 
 peter upon thin copper turnings and then gradually introducing concen- 
 trated sulphuric acid, or a concentrated sodium nitrite solution is 
 allowed to flow gradually into a hydrochloric acid solution of ferrous 
 chloride or sulphate : 
 
 FeCb + NaNOj + 2HCI = FeClg -f NaCl + HjO + NO. 
 
 The evolution begins in the cold [see Thiele, Ann. Chem. (1889) 253, 
 
 246]- 
 
 A constant stream of very pure nitric oxide may be obtained by pour- 
 ing concentrated sulphuric acid, containing about 2 per cent, of sodium 
 nitrite [Emich, Monatsh. f. Chem. (1892), xiii, 73], upon mercury. 
 Also by mixing a nitrite solution with one of potassium ferrocyanide and 
 acetic acid [Deventer, Ber. 26 (1893), i, 589]. 
 
 A colorless gas escapes, which, however, immediately forms brown 
 vapors when it comes in contact with the air, as it unites with the oxygen 
 of the latter to form nitrogen dioxide. Therefore, all the air must be 
 expelled from the generating vessel by nitric oxide, and the gas collected 
 over water after the interior of the apparatus has become colorless. 
 
 Nitric oxide is a colorless gas. Under normal conditions one liter of it 
 weighs T.3426 grams; its gas density also remains unchanged at — 100° 
 (in comparison with air of the same temperature, V. Meyer). It is con- 
 densed with difficulty. Its critical temperature is — 93°, and its critical 
 pressure 71 atmospheres (p. 48). Liquid nitric oxide is colorless. It 
 boils at — 153-6° and solidifies at — 167° to a white mass. It is slightly 
 soluble in water, but dissolves very readily in an aqueous solution of fer- 
 rous salts, imparting a dark-brown color to the liquid ; heat expels it 
 from the same. 
 
 Nitric oxide is readily soluble in nitric acid. As its solution becomes more concen- 
 trated, it assumes a brown, yellow, green or blue color, as nitrogen trioxide is formed 
 finally with anhydrous nitric acid : 
 
 2IINO3 + 4NO = 3N2O3 + H2O. 
 
 Potassium permanganate oxidizes it, like nitrous acid (p. 206), to nitric acid : 
 loNO 4- 6KMnO^ -f 9H2SO* = 10IINO3 p 3K2SO, -f fiMnSO^ + 4H2O. 
 
210 
 
 INORGANIC CHEMISTRY. 
 
 As nitric acid contains 53.3 |)cr cent, of oxygen, it is capable of sustain- 
 ing the combustion of some substances, but only those in whose com- 
 bustion enough heat is liberated to effect the breaking down of nitric 
 oxide into nitrogen and oxygen. Hence, iihosphorus continues to burn 
 in this gas, while a sulphur hame, developing only a slight heat, is ex- 
 tinguished ; ignited charcoal does the same, while a sihinter, burning 
 energetically, will continue to do so, when introduced into the gas. 
 On shaking a cylinder filled with nitric oxide with a few drops of readily 
 volatile carbon bisulphide, and bringing a flame to the mouth of the 
 vessel, the carbon bisulphide va])ors will (piietly burn in the gas, giving 
 a bright, luminous flame, emitting strong actinic rays ; in this combustion, 
 the carbon and suliihur of the carbon bisulphide unite with the oxygen 
 of the nitric oxide and form carbon dioxide and sulphur dioxide. 
 
 Nitric oxide is a strongly endothermic compound (see p. 212), and is consequently 
 exjdosive (p. 30). 
 
 On determining the (juantity of heat disengaged in the combustion of jdiosphorus, car- 
 bon or other sul)stances in this gas, it will be discovered that the same is greater (about 
 21.5 Cal.) than that which is developed by the combustion of these bodies in oxygen. 
 This can only be explained upon the theory that less heat is neces.sary for the separation 
 of NO into N and O than for the separation of the molecules of combined oxygen atoms 
 — an additional proof that the molecules of free oxygen (as of other elements) consist of 
 atoms. 
 
 With oxygen or air, nitric oxide at once forms brown vapors of nitro- 
 gen dioxide : 
 
 2NO -b O2 = 2NO.,. 
 
 2 vols. I vol. 2 vols. 
 
 With less oxygen, nitrogen trioxide is produced (p. 205). Nitric oxide 
 also combines with chlorine to nitro.syl chloride, NOCl (p. 204), and with 
 bromine it yields nitrosyl bromide, NOBr. At a red heat nitric oxide 
 becomes nitrogen dioxide and nitrogen. With hydrogen and moderate 
 heat it forms water and nitrogen : 
 
 NO + = N + H.p ; 
 
 a mixture of both gases burns with a greenish-white flame. On conduct- 
 ing a mixture of nitric oxide and hydrogen over platinum sponge, water 
 and ammonia are produced : 
 
 2NO + 5H2 = 2NH3 + 2H2O. 
 
 The volumetric analysis of nitric oxide gas may be easily executed as 
 follows: Fill a bent glass tube over mercury with the gas; introduce 
 into the same a piece of sodium and heat the latter with a lamp. The 
 sodium combines with the oxygen, and free nitrogen separates; the 
 volume of the latter always ecpials half the volume of the nitric oxide 
 gas emi)loyed ; this follows from the formula NO : 
 
 2NO = N.2 + O.,. 
 
 2 vols. I vol. 1 vol. 
 
 Tn nilrogeti dioxide nitrogen apjicars to 1)0 ((uadrivalent, in nitric oxide bivalent, while 
 usually it is Irivalent or (|uin(|uivalenl. Attention has already been called to the similarity 
 of nitrogen dioxide and chlorine dioxide, CIO.^ (p. 207). 
 
NITROUS OXIDE. 
 
 21 I 
 
 Nitrous Oxide, N.^ 0 , is formed when zinc or tin acts ni)on dilute 
 nitric acid of specific gravity i.i. It may be best obtained by heating 
 ammonium nitrate, which at about 170° breaks down directly into water 
 and nitrous oxide : 
 
 NH.NOs = N ,0 + 2 H, 0 . 
 
 It is very probable that nitraniide, NOg . NH2, is produced transitionally, but it immedi- 
 ately decomposes into water and nitrous oxide (p. 204). The decomposition is very 
 energetic at about 240° and is frequently accompanied with explosion. It is advisable 
 to use a well-dried mixture of equivalent quantities of sodium nitrate and ammonium sul- 
 phate. The reaction will then proceed quietly, beginning at 230°, and no explosion will 
 occur (Smith). 
 
 The formation of the oxide from hydroxylamine nitrite (the isomeride of ammonium 
 nitrate) is interesting. The decomposition takes place in aqueous solution and proceeds 
 according to the equation : 
 
 NO.2 . NH3OH = N.p + 2H2O. 
 
 Nitrous oxide is a colorless gas, of sweetish taste and slight odor. Its 
 density corresponds to the molecular formula N.^O = 44.08. One liter 
 weighs under normal conditions 1.9706 grams. In cold water it is toler- 
 ably soluble (i volume of water dissolves at 0° 1.305 volumes of nitrous 
 oxide) ; therefore it must be collected over warm water or mercury. By 
 pressure and cold it condenses to a colorless liquid of specific gravity 
 0.937 and boils under the ordinary pressure at — 89.8°. By rapid evapo- 
 ration, liquid nitrous oxide solidifies to a crystalline mass which remelts 
 at — 102.3°. If the aqueous nitrous oxide be evaporated over an air- 
 pump its temperature falls to — 140°. Its critical temperature is -{-38.8° ; 
 critical pressure, 775 atmospheres (p. 48). 
 
 Although this oxide contains less oxygen than nitric oxide, it supports 
 the combustion of many bodies more readily than the latter, because it is 
 more easily decomposed into oxygen and nitrogen. A glimmering chi}) 
 inflames in it, as in oxygen ; phosphorus burns in it with a bright, luminous 
 flame, while a sulphur flame liberating very little heat is extinguished. A 
 mixture of equal volumes of nitrous oxide and hydrogen explodes like 
 detonating gas, only less violently : 
 
 N2O -f H2 = N2 -f HjO. 
 
 I vol. I vol. I vol. I vol. 
 
 It resembles oxygen very much, but can be distinguished from it by 
 not producing brown vapors with nitric oxide, as does the former. It 
 is not capable of combining with oxygen. When it is conducted 
 through a red-hot tube it is decomposed into nitrogen and oxygen. It 
 has an exhilarating effect when inhaled in slight quantity, and is, there- 
 fore, termed laughing gas. 
 
 Its volume composition may be determined in the same manner as 
 with nitric oxide, viz., by heating a definite volume of the gas with 
 potassium. Then we learn that from one volume of nitrous oxide an 
 equal volume of nitrogen will be separated — corresponding to the molec- 
 ular formula : 
 
 N./! + K.2 = N.2 + K2O. 
 
 I vol. 1 vol. 
 
212 
 
 INORGANIC CMF.MISTRY. 
 
 Hyponitrous Acid, — IIO-N N-() 1 I. In tlie list of reduction prodm Is 
 
 of nitric acid extending as far down as ammonia : 
 
 lINOg, UNO,, (UNO),, Nil, Oil, NII3, 
 
 this acid occupies a position between nitrous acid and hydroxylamine. Accordingly, it 
 can be obtained — 
 
 (1 ) by the reduction of nitric and nitrous acids ; 
 
 (2) by the oxidation of hydroxylamine ; and 
 
 (3) by the union of nitrous acid and hydroxylamine with the exit of water. 
 
 The reduction of the salts of nitric or nitrous acids may be effected by sodium amal- 
 gam or by electrolysis ; the oxidation of hydroxylamine in acjueous solution by various 
 metallic oxides, but especially by mercuric oxide: 2NII2OII -(- 2llg() = IKJN 
 NOll (- 2 ll ,0 T 2llg. To obtain hyponitrous acid according to (3) allow sodium 
 nitrite and hydroxylamine sulphate to interact, when sodium sulphate and hydroxylamine 
 nitrite will result: 2NaN(), -j- SO^(NH^())2 r- Na,S()^ -f- 2N(J2. Nil/), which for the 
 most part breaks down into water and nitrous oxide (p. 212), but in small amount to 
 water and hyponitrous acid: IIONII, -t- ON. OH — HON NOH 11 , 0 . Divers 
 claims that good yields of potassium hyponitrite can be obtained by boiling ])otassium 
 hydroxylamine sulphonate (p. 197) with caustic potash : 2NH(OH)S03K -f- 4K(J)H = 
 K2N.,0, -f- 2lv2'^^\3 411,0. Upon neutralizing the hyponitrite solution, obtained by 
 
 one or the other method, with acetic acid and adding silver nitrate the silver salt Ag.^N,!), 
 separates as a bright yellow, amorphous powder, which slowly decomposes at 100°, but at 
 110° with violent explosion. Concentrated sulphuric acid liberates nitrous oxide from it. 
 
 The acid may be obtained in ethereal solution by adding gradually an excess of its 
 silver salt to a solution of dry hydrogen chloride in anhydrous ether. After evaporating 
 the ethereal solution in a vacuum, free hyponitrous acid remains in crystalline leaflets. It 
 explodes readily — with flame — oncoming in contact with solid caustic potash ; deliquesces 
 in the air, is very soluble in alcohol, and soluble in benzene and in chloroform. It is 
 more stable in aqueous solution than in the solid condition, but even then readily decom- 
 poses into nitrous oxide and water: HgN.^O, = N ,0 -)- 11 , 0 . It therefore bears the 
 same relation to nitrous oxide as nitric acid to nitrogen pentoxide, and nitrous acid to 
 nitrogen trioxide ; however, nitrous oxide does not take up water and it, therefore, does 
 not behave like the true anhydride of hyponitrous acid. Upon neutralizing its aqueous 
 solution with caustic soda the acid salt NaHN202 is produced, which quickly breaks 
 down into caustic soda and nitrous oxide : NaHN.,^^, = N,!^ + NaOH. The freshly 
 prepared solution of the acid liberates iodine from potassium iodide after it has stood for 
 some time ; this is because it decomposes very slowly and only to a slight degree into 
 nitrous acid and ammonia: 3H2N2O2 =: 2N2O3 -j- 2NH3. It decolorizes potassium 
 permanganate and becomes nitric acid. 
 
 The molecular magnitude of the acid was established by the determination of the vapor 
 density of its ethyl ester, N202(C2H5)2, by Zorn, and from the molecular weight of the 
 benzyl ester, N202(C7lI.j)2, by Hantszch and Kaufmann. 
 
 Hyponitrous acid is isomeric with nitramide (p. 204). 
 
 Compare the investigations of Hantzsch and his co-workers : Ann. Chem. 292 (1896), 
 317 ; 299 (1898), 94; Divers ibid. 295 (1897), 366; also Kirschner, Z. f. anorg. Chem. 
 16 (1898), 424. 
 
 d'he thermo-cheinical relations of the oxygen derivatives of nitrogen give some clue 
 to their chemical deportment. All nitrogen oxides are endothermic compounds, i. e., 
 they are produced from their elements with heat absorption corresponding to the symbols : 
 
 (N„ 0 )=-i 7.5 (N, 0 ) = -2i.6 (N, 0 ,)=-S. 
 
 I'rom this we observe that the oxides of nitrogen cannot be ]'>repared from the elements 
 without addition of energy. Proceeding from nitric oxide (Nt)), we see from the 
 above numbers that the formation of the liigher oxides from it occurs with heat disen- 
 gagetnent : 
 
 (2NO, i)) ■= 20.1 (NO, O) = 13.4 (2NO,, (),) = 40.5, 
 
OXYGEN COMPOUNDS OF PHOSPHORUS. 
 
 213 
 
 whereas lieat is absorbed in the conversion of nitrous into nitric oxide : (N2O, O) = 
 
 — 25.4. (See J. 'I'homsen’s Thermocheniische Untersuchungen, iv, 314.) 
 
 lieat disengagement, on the contrary, occurs in the production of nitric acid from its 
 elements : 
 
 (N,03,H-liquid) =41.6 (N,03,H,Aq) = 49.1. 
 
 This explains the relative stability of that acid. 
 
 Compounds of Nitrogen with Sulphur. — Two such bodies are known : nitrogen 
 sulphide, N^.S^, corresponding in composition to a quadrupled nitric oxide, and nitrogen 
 pentasulpJnde, corresponding to nitric anhydride, NgO^. 
 
 Nitrogen sulphide, N4S4, is produced when dry ammonia, in benzene solution, acts 
 upon sulphur dichloride : 
 
 4NH3 + 6SCI2 =: N4S4 + I2HCI + 83. 
 
 It consists of orange-red needles, melting at 178°. It detonates at higher temperatures 
 or when struck. Its molecular weight has been determined by methods which will be 
 described under Solutions. 
 
 Nitrogen pentasidphide, N4S5, is formed by a complicated reaction occurring on heat- 
 ing nitrogen sulphide with carbon bisulphide to 100° in a sealed tube. It is a deep-red 
 liquid, transparent in thin layers with a blood-red color. It is very mobile, not moisten- 
 ing glass and its specific gravity at 18° equals 1.90. It solidifies by cold to a mass resem- 
 bling iodine and melts at 10°. Its odor is like that of iodine and is at the same time 
 sw'eet. Nitrogen pentasulphide is not soluble in water, but in organic solvents. Its solu- 
 tions are more stable than the pure pentasulphide, w Inch easily decomposes into its ele- 
 ments. Boiling water, alkalies and hydrogen sulphide decompose the pentasulphide 
 very readily with the production of ammonia and the separation of sulphur. [See Z. f. 
 anorg. Chem. 13 (1897), 200.] 
 
 2. OXYGEN COMPOUNDS OF PHOSPHORUS. 
 
 P4O6 
 
 Phosphorus 
 
 trioxide. 
 
 P2O5 
 
 Phosphorus 
 
 pentoxide. 
 
 Hypophosp 
 
 acid. 
 
 horous 
 
 H3PO3 
 
 Phosphorous 
 
 acid. 
 
 Orthophosphoric 
 
 acid. 
 
 Oxides containing less oxygen, e. g., the tetroxide and the oxide 
 P./j, are not definitely known. 
 
 The following anhydride acids are derived from orthophosphoric acid : 
 
 HPO3, Metapho.sphoric acid. 
 
 H4P2O7, Pyrophosphoric acid. 
 
 Orthophosphoric acid also yields an anhydride acid with phosphorous 
 acid : hypophosphoric acid, H4P20g. 
 
 The structure of these compounds is expressed by the following 
 formulas : 
 
 V HI V OH V 
 
 II3PO-OII p ^OII iikxCh 
 
 Hypophosphorous Phosphorous Phosphoric 
 
 acid. acid. acid. 
 
214 
 
 INORGANIC CHEMISTRY. 
 
 In hypophos])horoiis acid two atoms of liydrogcn are in direct union 
 with quinquivalent phospliorus, while the third hydrogen atom forms an 
 hydroxyl group with oxygen. It is only this last atom of hydrogen which 
 is readily replaced by the action of bases ; hence hypophosphorous acid 
 is a 7 nonobasic acid. Phosphorous acid (like sulphurous and arsenious 
 acids) appears in two forms in its derivatives (pp. i86, 221). It is very 
 probable that in its salts it has one atom of hydrogen and two hydroxyl 
 groups joined to ])hosphorus; consecjuently it is a dibasic acid. [See 
 Michaelis and Pecker, Per. 30 ( 1897), 1003.] Finally, three series of salts 
 and three hydroxyls are assumed to be present in phosphoric acid. Py 
 the elimination of one molecule of water from ])hosi)horic acid, meta- 
 ])hosphoric acid results — an anhydride which, at the same time, is a 
 monobasic acid, as it contains one hydroxyl group : 
 
 V 
 
 POj-OII, Metaphosphoric acid. 
 
 On removing one molecule of water from two molecules of jihosphoric 
 acid, pyro- or diphosphoric acid is formed (see }). 193) : 
 
 V /OH 
 PO^OH 
 
 ^OH 
 
 V /OH 
 PO^OH 
 
 \OH 
 
 — H,0 
 
 2 molecules Phosphoric acid. 
 
 /OH 
 
 PO^OH 
 
 PO— OH 
 ^OH. 
 
 I molecule Pyrophosphoric acid. 
 
 Pyrophosphoric acid contains four hydroxyl groups, hence is tetra- 
 basic. Similarly, hypophosphoric acid results from one molecule each 
 of ])hosphoric acid and phosphorous acid, P(OH)3. 
 
 Finally, if from two molecules of phosphoric acid all the hydrogen atoms 
 be removed, in the form of water, an anhydride remains : 
 
 V V 
 O2P-O-PO2. 
 
 Phosphoric anhydride. 
 
 The salts of phosphoric acid are termed phosphates ; those of phos- 
 phorous acid, phosphites^ and of hypophosphorous acid, hypophosphites. 
 
 Hypophosphorous Acid, H3PO2. Hydrogen phosphide e.scapes 
 when a concentrated solution of sodium or barium hydroxide is warmed 
 with yellow phosphorus, leaving behind in solution a salt of hypophos- 
 phorous acid : 
 
 4? 4- 3NaOII -p 3II.P = 3TT.TO.ONa + PH, 
 
 8P -f 3l5a(OHh + 3II2O = 3(Il2l'0. 0)Ta+ 2Pli3. 
 
 d'he free acid may be separated from the barium salt by means of sul- 
 ])huric acid ; the insoluble barium suljihate being filtered off from the 
 acpieous solution of the acid, and the latter concentrated under the air- 
 
PHOSPHOROUS ACID. 
 
 215 
 
 pump. Hypophosphorous acid is a colorless, thick liquid, with a strong 
 acid reaction. Below 0° it sometimes solidifies to large, white leaflets, 
 which melt at Heat converts it, with much foaming, into 
 
 hydrogen phosphide and phosphoric acid : 
 
 2H3P0,^PH3 + II3PO,. 
 
 It absorbs oxygen readily, becoming phosphoric acid, hence acts as a 
 powerful reducing agent. It reduces sulphuric acid to sulphur dioxide, 
 and even to sulphur. It precipitates many of the metals from their solu- 
 tions ; from copper sulphate it separates the hydride CiqH^. 
 
 The acid is monobasic, H2PO.OH (see p. 214). Its salts dissolve 
 readily in water, and absorb oxygen from the air, thus becoming phos- 
 phates. When heated in a dry condition, they set free phosphine and 
 hydrogen and are converted into pyro- and metaphosphates; some also 
 yield metallic phosphides. Nascent hydrogen reduces hypophosphorous 
 acid to phosphine. 
 
 Phosphorous Acid, H3PO3, is formed at the same time with phos- 
 phoric acid and hypophosphoric acid in the slow oxidation of phosphorus 
 in moist air. The decomposition of phosphorus trichloride by water gives 
 it more conveniently : 
 
 PCI3 + 3H2O = H3PO3 + 3HCI. 
 
 By evaporating this solution under the air-pump the phosphorous acid 
 becomes crystalline. The crystals are readily soluble in water, and 
 deliquesce in the air. It melts at 70°, and decomposes on further heating 
 into phosphine and phosphoric acid : 
 
 4H3PO3 = PH3 + 3H3PO,. 
 
 In the air the acid absorbs oxygen, and changes to phosphoric acid. 
 Hence, it is a strong reducing agent, and precipitates the free metals 
 from many of their solutions. In the presence of water the halogens 
 oxidize it to phosphoric acid. 
 
 It is a dibasic acid, forming two series of salts, in which one or two 
 atoms* of hydrogen are replaced by metals. The phosphites do not oxi- 
 dize in the air, except under the influence of strong oxidizing agents 
 When heated, they generally decompose into hydrogen, pyrophosphates 
 and phosj)hide. Nascent hydrogen also reduces phosphorous acid to phos- 
 phine — a circumstance of importance in the detection of phosphorus- 
 poisoning. 
 
 Phosphorous Anhydride, or phosphorus trioxide according to the old formula P2O3, 
 has been shown by 4 'horpe and Tutten (1892) to be correctly expressed by 1 ^) 6 - B is 
 
 * Therefore, the structural formula, TIPO(OII)2 is assigned to this acid. There 
 appears to exist another phosphorous acid, at least in compounds, to which the formula 
 P(OII)3 belongs (pp. 213, 216). 
 
2i6 
 
 INORGANIC CUKMISTRV. 
 
 produced on conducting dry air over gently lieated jdiospliorus, or by carefully beating 
 phosphorous acid with i)hosphorus trichloride : 
 
 2II3PO3 + 2PCI3 =: P,Og + 6 IIC 1 . 
 
 It is a white, flocculent mass or it consists of colorless needles. It melts at 22.5°, sub- 
 limes readily and boils in an atmosphere of nitrogen at 173°. I ts vapor density corre- 
 sponds to the formula I^Og. At 400° it breaks down into phosphorus and phosphorus 
 tetroxide, PjO^, which crystallizes. It is decomposed by water in a very complicated 
 manner. 
 
 Phosphoric Acid, H3P0^, or Orthophosphoric acid, is produced when 
 the peiitoxide is dissolved in hot water, and by the decomposition of 
 phosphorus pentachloride or phosphorus oxychloride (POCl,) by water (see 
 p. 219). It may be obtained by decomposing bone ash, Ca3(PO^)2, with 
 sulphuric acid, or better, by oxidizing yellow phosithorus with nitric acid. 
 The aqueous solution is evaporated to dryness in a platinum dish. The 
 anhydrous acid consists of colorless, hard, prismatic crystals, which in 
 the air deliquesce to a thick, acid liquid. 
 
 Phosphoric acid is tribasic, forming three series of salts, q. 2\\^(\ primary 
 (KH2P0^), secondary (K^HPO^), and tertiary (K3POJ. I'hey may be 
 spoken of according to the number of hydrogen atoms rei)laced by metals, 
 as, e. g., monopotassium phosphate (KH2POJ, di])otassium phosphate 
 (K2HPO4), and tripotassiiim phosphate (K3POJ. The salts of the first 
 two series contain hydrogen, replaceable by metals, hence may be termed 
 acid, while the salts of the third series are neutral. Their behavior with 
 litmus does not harmonize with this view (see Sodium Phosphate). 
 
 The tertiary phosphates, excepting the salts of the alkalies, are insolu- 
 ble in water. With a silver nitrate (AgNOg) solution, soluble phosphates 
 give diyellow precipitate of silver phosphate, AggPO^. 
 
 Pyrophosphoric Acid, PI^P207 (structure, p. 214), is formed by 
 the continuous heating of orthophosphoric acid to 260°, until a portion 
 of it dissolved in ammonium hydroxide does not yield a yellow but a pure 
 white preci])itate with silver nitrate. The sodium salt is easily obtained 
 by heating disodium phosphate: 
 
 2Na2HPO^ = NaTjO; + H 2 O. 
 
 Other salts are similarly formed by heating the corresponding ortho- 
 phosphates. 
 
 'Fhe acid presents a white crystalline appearance, and is readily soluble 
 in water. When in solution, it slowly takes up water at ordinary tem- 
 ])eratures, more rapidly when heated, and, like all true anhydrides, passes 
 into the corresponding acid — orthophosi)horic acid. 
 
 Pyrophosphoric acid is tetrabasic. Its salts are very stable, and are 
 not altered by boiling with water ; warmed with dilute acids, they become 
 salts of the ortho-acid, 'bhe soluble salts yield a 7 vhife ])recipitate of silver 
 ])yro]jhosphate, AgJ^O^, with silver nitrate. 
 
 Hypophosphoric Acid, II,P20^^. While i^yrophosphoric acid is an anliydride acid 
 of |)hoH])lioric acid, the so-called hy|)ophosplu)ric may he viewed as a mixed anhydride 
 of pliosj)lioric and symmelrical phosj)horoiis acids. 
 
METAPHOSPHORIC ACID. 
 
 217 
 
 PO(OH), 
 
 It is produced, as demonstrated by Salzer in 1877, together with phosphorous and phos- 
 phoric acids, by the slow oxidation of moist phosphorus in the air. It is separated from 
 these acids by means of its difficultly soluble sodium salt, NajH.^P.^Og -(- 6H.^O ; by pre- 
 cipitating the solution of the latter with a soluble lead salt we get insoluble lead hypo- 
 phosphate, PbjP.^Og. Its silver salt is more easily obtained by oxidizing phosphorus in 
 the presence of silver nitrate. The free acid separated from the lead or silver salt by 
 hydrogen sulphide is rather stable in a dilute solution, and below 30° may be concen- 
 trated to a syrup. At higher temperatures, more readily in the presence of hydrochloric 
 or sulphuric acid, the acid decomposes into phosphoric and phosphorous acids. It does 
 not reduce metallic salts, but is oxidized by potassium permanganate to phosphoric acid. 
 [Salzer, Ann. Chem. 211 (1892), i. See Z. f. anorg. Chem. 6 (1894), 132, for a con- 
 venient method for its production.] 
 
 Metaphosphoric Acid, HPO3 or PO2 . OH, results upon heating the 
 ortho- or pyro-acid to 300°. It can be more conveniently obtained by 
 dissolving phosphorus pentoxide in cold water : 
 
 P2O5 -f HjO = 2HPO3. 
 
 It is a vitreous, transparent mass (^Acidum pJwsphoricum glaciale), which 
 generally contains less hydrogen than is required by the formula HPO3. 
 This probably is because some anhydride is present in it. It melts when 
 heated and volatilizes at higher temperatures, without suffering any 
 change. It deliquesces in the air, and dissolves with ease in water. 
 (Commercial glacial phosphoric acid contains sodium and magnesium 
 phosphate, and dissolves with difficulty in water.) The solution coagu- 
 lates albumin ; this is a characteristic method of distinguishing meta- 
 from ortho- and pyrophosphoric acid. In aqueous solution, metaphos- 
 phoric acid changes, gradually at ordinary temperature, rapidly by boil- 
 ing, into orthophosphoric acid : 
 
 HPO3 + H2O = H3PO,. 
 
 It is a monobasic acid. Its salts, the metaphosphates, are readily 
 obtained by the ignition of the primary salts of the ortho-acid : 
 
 Nall^PO, = NaP03 -f- H^O. 
 
 When the aqueoussolutions of these salts are boiled, they are converted 
 into the primary salts of ortho})hosphoric acid. With silver nitrate the 
 soluble metaphosphaies give a white precipitate of silver metaphosphate. 
 
 In addition to the ordinary salts of metaphosphoric acid, various modifications of the 
 same exist ; these are derived from the polymeric meta-acids, ri2P.20g, H3P3O,,, H^P^O,2, 
 etc. They are all changed to primary ortho-phosphates by boiling their solutions. [See 
 Tammann, Jr. prakt. Ch. 45 (1892), 417.] 
 
 Phosphorus Pentoxide, PaO^, or Phosphoric anhydricte, is obtained 
 by burning phosphorus in a current of dry oxygen or dry air. 
 
 19 
 
2i8 
 
 INORGANIC CHEMISTRY. 
 
 The following procedure serves for the jjrcparation of it (Fig. 65) : A piece of phos- 
 phorus, placed in an iron dish attached to the glass tube a b, is burned in the glass 
 balloon A. The necessary amount of air is drawn through the vessel by means of an 
 aspirator. The air is lirst ])assed through the bent tube containing pieces of pumice- 
 stone, moistened with sulphuric acid, in order to dry it perfectly. After the j)hosphorus 
 has been consumed, fresh pieces of it are introduced into the little di.sh through a /f>, and 
 the upper end of the tube closed with a cork. The phosphorus pentoxide formed collects 
 partly in A and partly in the receiver (//). 
 
 Phosphorus pentoxide is a wliite, voluminous, flocculent mass. It 
 attracts moisture energetically and delitpiesces in the air. It dissolves in 
 cold water with hissing and yields inetajihosphoric acid. Owing to its 
 great affinity for water it serves as an agent for drying gases, and also for 
 the withdrawal of water in chemical reactions, e. g., in the formation of 
 acid anhydrides from acids (p. 204). 
 
 Tt generally contains oxides of lower oxygen content. It can be freed from these by 
 subliming it over ignited platinum sponge in a current of oxygen. 
 
 Fig. 65. 
 
 Chloranhydrides of the Acids of Phosphorus. — The halogen derivatives of phos- 
 phorus, considered on p. 141, may be viewed as the halogen anhydrides of phos- 
 phorous and phosphoric acids (p. 195). The compounds PCI3, PBrg, and PI3 are 
 derived from phos})horous acid, because they yield the latter acid with water : 
 
 PCI, h 311,0 = II3PO3 + 3IICI. 
 
 The compounds, jdiosphorus pentachloride, PCI5, and phosphorus pentabromide, 
 Plhj, correspond to the normal orthophosphoric acid, P(01I)5, which has not been 
 obtained in a free condition (.see Calcium Phosphate). 
 
 Phosphorus Oxychloride, POCI3, can be viewed as the chloride 
 of the known’ orthophosjihoric acid, PC)(()H)3. It is a colorless liquid, 
 fuming strongly in the air; its siiecific gravity equals 1.68 at 15°. It 
 
COMPOUNDS OF PHOSPHORUS WITH SULPHUR. 
 
 219 
 
 congeals on cooling and then melts at — 1-5°. It boils without, decom- 
 position at 107°; its vapor density corresponds to the formula — 
 
 POCI3 = 153.0. 
 
 Water decomposes it intometa- or orthophosphoric acid and hydrochloric 
 acid : 
 
 rocij + 2H2O = H PO3 + 3HCI 
 POCI3 + 31102 = HgPO^ + 311 Cl. 
 
 It may be obtained by decomposing the pentaehloride with a little water 
 (see above), or by allowing it to deliquesce gradually in moist air. A 
 more practical method consists in distilling the pentaehloride with phos- 
 phorus pentoxide : 
 
 3PCl3 + P203 = 5POCl3; 
 
 or with crystallized boric acid (6 parts to i part) : 
 
 3PCI5 -f 2H3BO3 = 3POCI3 + B2O3 + 6 HC 1 . 
 
 A noteworthy formation of the oxychloride is the union of phosphorus 
 trichloride with oxygen on passing ozonized air through it : 
 
 PCI3 + 03 = POCI3 + O2. 
 
 Potassium chlorate acts in a corresponding manner very energetically 
 upon the trichloride with the production of oxychloride: 
 
 3PCI3 + KCIO3 = 3POCI3 4 KCl. 
 
 The compound PSCI3 is analogous to the oxychloride POCI3. It is obtained by heat- 
 ing the trichloride and sulphur to 130° ; also by the action of phosphorus pentaehloride 
 upon hydrogen sulphide or some metallic sulphides : 
 
 PCI5 + H2S = PSCI3 -f 2HCI. 
 
 Phosphorus sulphochloride, PSCI3, is a colorless liquid of specific gravity 1.6, fuming in 
 the air and boiling at 124-125°. Water decomposes it into phosphoric and hydrochloric 
 acids and hydrogen sulphide : 
 
 PSCI3 -f 4H2O = H3PO4 p 3HCI + H2S. 
 
 The chlorides PO2CI and P203C1^, corresponding to meta- and pyro- 
 phosphoric acids, have also been prepared. (See Chem. Centralblatt, 
 1897, II, I4-) 
 
 COMPOUNDS OF PHOSPHORUS WITH SULPHUR. 
 
 With sulphur, phosphorus affords a number of compounds which are 
 obtained by direct fusion of phosphorus with sulphur. As the union 
 of ordinary phosphorus and sulphur usually occurs with violent explo- 
 sion, red phosphorus should be employed in preparing these compounds. 
 
 The compounds PjS, and PjS^, analogous in constitution to P^O^ and 
 PjOj, are solid crystalline substances, melting at higher temperatures and 
 
220 
 
 INORGANIC CHEMISTRY. 
 
 subliming without decomposition; phosphorus pentasulphide boils at 
 520° and the trisulphide at 540° (cor.). Water changes them to hydro- 
 gen sulphide and the corresponding acids, i)hosphorous and phosphoric. 
 They combine with the alkaline suli)hides to form compounds (e. g., 
 K3PSJ which possess a constitution analogous to that of the salts of phos- 
 phoric and phosphorous acids [see Suli)ho-salts of Arsenic, j). 223; also 
 Glatzel, Z. f. anorg. Chem. 4 (1893), 187]. 
 
 Well-crystallized sulphur phosphides, the composition of which corre- 
 sponds to the formulas P^Sg and PgS^, have been i)repared. The 
 
 supposed compounds P^S and P^S are licpiids which inflame readily in 
 contact with air, but have been proved to be mixtures containing free 
 phosphorus. 
 
 Besides the preceding, we have other phosphorus derivatives which contain nitrogen. 
 These have been little studied, and at present are unimportant. Such compounds 
 are PNjfl (phospham), PNO, PNCI.^. The so-called amid-derivatives, P()Cl2.Nll2, 
 P0C1(NH2)2 and PO(NH2 )s» are produced by allowing ammonia to act upon phosphorus 
 oxychloride, POCI3. In these chlorine is replaced by the amido-group Nlf^. 
 
 Recently, amidophosphoric acid, P0(NIl2) (011)2, corresponding to sulphamic acid, 
 
 S02(NH2)0H, and imidophosphoric acid, the counterpart of imido- 
 
 sulphuric acid, as well as its derivatives, have been more exhaustively studied [Stokes, 
 Jahrb. d. Chem. vi (1896), 89]. 
 
 3. OXYGEN DERIVATIVES OF ARSENIC. 
 
 AS2O3 .... 
 
 Arsenic trioxide. 
 
 AS2O5 HjAsO^. 
 
 Arenic pentoxide. Arsenic acid. 
 
 Arsenic Trioxide, or Arsenioiis anhydride (^Acidum arseni- 
 
 osum), occurs in nature as arsenic “bloom.” It is produced by the 
 burning of arsenic in oxygen or in the air, and by the oxidation of the 
 metal with dilute nitric acid. It is obtained metallurgically on a large 
 scale as a by-product in the roasting of ores containing arsenic. The 
 trioxide thus formed volatilizes and is collected in walled chambers, in 
 which it condenses in the form of a white powder {white arsenic, poison 
 flour). To render it pure, it is again sublimed in iron cylinders, and 
 obtained in the form of a transparent, amorphous, glassy mass {arsenic 
 glass), the specific gravity of which equals 3.74. Upon preservation this 
 variety gradually becomes opaque and porcelaneoiis, acquires a crystalline 
 structure, and its specific gravity decreases to 3.69. Upon dissolving 
 this oxide in hot hydrochloric acid, it crystallizes, on cooling, in shin- 
 ing, regular octahedra. At the same time, this interesting phenomenon is 
 observed : that when the solution of the glass variety crystallizes it phos- 
 phoresces strongly in the dark, while the porcelaneoiis does not exhibit 
 this jiroperty. Arsenic trioxide crystallizes in similar forms of the regu- 
 lar system when its vapors are rajiidly cooled, but upon cooling slowly, 
 it assumes the shape of monoclinic jirisms of sjiecific gravity 4.0 ; 
 therefore it is dimorphous. When heated in the air, it sublimes above 
 
OXYGEN DERIVATIVES OF ARSENIC. 
 
 221 
 
 218°, without melting; under higher pressure, however (in sealed tubes), 
 it melts to a liquid which solidifies to a glassy mass. 
 
 The molecular magnitude of the solid arsenic trioxide, like that of all solids, is not 
 known. When converted into vapor it behaves like sulphur and other solid liquid bodies ; 
 in vapor form at lower temperatures the molecules contain more atoms than at higher 
 temperatures. According to Biltz (Z. f. phys. Ch. 20 (1896), 68), at temperatures rang- 
 ing from 500° to 700° the molecules of As^Og predominate, while between 700° and 1800° 
 the molecules of AsjOg increase in number. Above 1800° the density corresponds to the 
 old formula AsjO,, and for this reason and simplicity’s sake it is retained in the text. It 
 is interesting to observe that at lower temperatures the same group As^ is present in the 
 molecule of arsenious oxide as was observed in the molecule of arsenic (p. 143). 
 
 The trioxide dissolves with difficulty in water ; the solution possesses a 
 sweetish, unpleasant metallic taste, exhibits but feeble acid reaction, and 
 is extremely poisonous. The oxide is very soluble in acids, and probably 
 forms salts with them; at least, on boiling, a solution of the trioxide 
 volatilizes in strong hydrochloric acid, arsenious chloride, AsClg. From 
 this and its feeble acid nature we perceive an indication of the basic char- 
 acter of the trioxide corresponding to the already partly metallic nature 
 of arsenic (see p. 148). 
 
 Nascent hydrogen converts the trioxide into arsine (AsHg) ; but when 
 heated with charcoal it is reduced to the metallic state (pp. 143, 144). 
 Upon heating arsenic trioxide in a narrow glass tube with carbon, the 
 reduced arsenic deposits as a metallic mirror on the sides. Oxidizing 
 agents convert arsenic trioxide into arsenic acid. 
 
 Arsenious Acid, HgAsOg, corresponding to the anhydride, is not 
 known in a free condition. It probably exists in the aqueous solution 
 of arsenic trioxide, but the anhydride separates out upon evaporation. 
 In its salts {arsenites) it is tribasic and usually affords tertiary derivatives : 
 
 AgsAsOg, Mg3(AsOg).^. 
 
 The alkali salts, soluble in water, absorb oxygen from the air and serve 
 as powerful reducing agents, they themselyes becoming arseniates. Silver 
 nitrate forms a yellow-colored precipitate, AggAsOg, with the soluble salts. 
 [See Stavenhagen, Jr. f. prakt. Chem. 51 (1894), i]. 
 
 Other salts exist which are derived from the meta-arsenious acid, 
 HASO2 (p. 201). 
 
 According to Klinger, arsenious acid, like sulphurous and phosphorous acids, is pres- 
 
 yOW 
 
 ent in its derivatives in two forms : (i) the symmetrical As^OH, from which the silver 
 
 ^OH 
 
 OH 
 
 salt, the chloride and the esters are derived, and (2) the unsymmetrical H-As<q 
 or H - As^OH, from which the alkali salts and alkyl arsonic acids (see Organic Chem- 
 
 ^01 1 
 istry) have their origin. 
 
 Arsenic Acid, HgAsO^ ( Acidtim arsenicuni), is obtained by the oxi- 
 dation of arsenic or its trioxide with concentrated nitric acid or by 
 means of chlorine. Upon evaporating the solution rhombic crystals of 
 
222 
 
 INORGANIC CHEMISTRY. 
 
 the forniiila TI^AsO^ -f- sej)aralc out; these deliiiuesce on exposure 
 
 to the air. 'I'liey melt at ioo°, lose their water of crystallization and 
 yield ortho-arsenic acid, H3AsO^, which heated to 140-180° passes into 
 Pyro-arse?iic acid, H^As207 : 
 
 2II3ASO, — . H.AsjO^ -f II2O. 
 
 At 200° this again loses water and becomes Meta-arscnic acid, HAsOg. 
 With water the last two acids become ortho- again ; hence the latter is 
 perfectly analogous to idiosphoric acid. A portion of the arsenic acid 
 is volatilized, i)robably as arsenious chloride, when strong hydrochloric 
 acid solutions of arsenic acid are heated. 
 
 At a red heat the meta-arsenic acid loses all its water and becomes 
 Arsenic pentoxide, AS2O5, a white, glassy mass. Very strong ignition 
 decomj)oses this into arsenic trioxide and oxygen ; in contact with water 
 it gradually changes to arsenic acid. 
 
 Ortho-arsenic acid is readily soluble in water, and is a strong tribasic 
 acid. Its salts — the arseniates — are very similar to the phosphates and 
 are isomorphous with them. With the soluble salts silver nitrate gives a 
 reddish-brown precipitate of tri-silver arseniate, AgjAsO^. 
 
 COMPOUNDS OF ARSENIC WITH SULPHUR. 
 
 Arsenic, like nitrogen and phosphorus, combines in several proportions 
 with sulphur : realt^ar, AS2S2, auripigment or trisulphide, AS2S3, and penta- 
 sulphide, AS2S5. They can be prepared by fusing arsenic and sulphur 
 together. The two containing the more sulphur maybe also obtained by 
 passing hydrogen sulphide into solutions of the arsenic oxides. This indi- 
 cates the metallic nature of arsenic, for this is the universal method 
 of preparing metallic sulphides: 
 
 As.^03 -|- 3H2S = AS2S3 + 3112^^ > AsgOg 5112^ — AS2S5 -|- 5 ll 2 f^* 
 
 Arsenic Trisulphide, AS2S3, occurs in nature as auripigment in bril- 
 liant yellow, leafy, crystalline masses of specific gravity 3.4. It is pre- 
 cipitated from acidulated solutions of arsenious acid or its salts by 
 hydrogen sulphide, as a lemon-yellow amorphous powder. It is insoluble 
 in water, but dissolves readily in ammonium hydroxide, the alkalies and 
 alkaline sulphides. It is worthy of note that it dissolves upon prolonged 
 heating in concentrated hydrochloric acid : 
 
 AS2S3 -|- 6IIC1 = 2ASCI3 -f 3H2S ; 
 
 the arsenious chloride then volatilizes with the hydrochloric acid vapors. 
 
 Arsenic Pentasulphide, AS2S5, separates as a bright yellow powder 
 from the solution of sodium sulpharseniate, Na3AsS4 (see below), upon 
 the addition of acids. Also if a rapid current of hydrogen sulphide be 
 conducted through a slightly acidulated arsenic acid or arseniate solution, 
 heated to 80° ; the arsenic acid is then, according to R. Bunsen, slowly 
 but completely converted into ])entasuli)hide : 
 
 2II3ASO, -f 51128 = AS2S3 -I 8II2O 
 
OXYGEN COMPOUNDS OF ANTIMONY. 
 
 223 
 
 [Ann. Chem. 192 (1878), 305]. Under other conditions arsenic acid 
 can be reduced by hydrogen sulphide entirely or in part to arseniousacid 
 with the separation of sulphur : 
 
 HaAsO^ -f HjS = HgAsOa + IiP -f S, 
 
 which is then either precipitated as trisulphide in the heat or in the 
 presence of much hydrochloric acid it may be volatilized. Consult 
 Brauner and Tomicek, Monatshefte f. Chem. viii (1887), 607; McCay, 
 Z. f. analyt. Chem. 27 (1888), 632, and Piloty and Stock, Ber. 30 
 (1897), 1649, upon the complex and as yet unexplained relations. 
 
 Arsenic Disulphide, As^Sg, also exists. It occurs in nature as 
 realgar y forming the beautiful, ruby-red crystals, of specific gravity 3.5. 
 It is applied as a pigment. It is prepared artificially by fusing arsenic 
 with sulphur in the proportion by weight expressed by its formula. 
 
 Arsenic Sulpho-salts. — Just as the oxygen compounds of arsenic combine with 
 basic oxides and hydroxides to form salts, so the sulpho-salts are formed by combination 
 of the arsenic sulphides with alkaline sulphides : 
 
 'AsjSg -j- 3K2S = 2 K,AsS 3 
 Tripotassium 
 sulpharseuite. 
 
 As.^Sj -f- 3K2S = 2K3 AsS^. 
 
 Tripotassium 
 
 sulpharseniate. 
 
 For the preparation of these sulpho-salts, arsenic sulphide is dissolved in the aqueous 
 solution of the alkaline sulphides or hydrogen sulphide is conducted through the alkaline 
 solution of the oxygen salts : 
 
 KgAsO^ -P 4H2S = K3ASS4 + 4H2O. 
 
 The sulpho-salts of the alkalies and ammonium are easily soluble in water, and when the 
 solution is evaporated they generally separate in crystals. Acids decompose them, arsenic 
 sulphide separating out and hydrogen sulphide becoming free : 
 
 2K3ASS4 -f 6 HC 1 == AS2S5 -f 6 KC 1 4- 3H2S. 
 
 The silver salt, Ag3AsS3, silver sulpharsenite, is the mineral frmistite. 
 
 Antimony, carbon, tin, gold, platinum and some other metals form sulpho-salts similar 
 to those of arsenic (and of phosphorus). 
 
 4. OXYGEN COMPOUNDS OF ANTIMONY. 
 
 The oxygen derivatives of antimony are analogous in constitution to 
 those of arsenic: antimony trioxide, Sb203, and antimony pentoxide, 
 81^205. The similarity of the two metals is indicated here just as in their 
 chlorides. The trioxide does not possess acid but basic ])roperties 
 almost solely; it is true that a few salt-like compounds with such power- 
 ful bases as caustic soda and potash are known, but they are immediately 
 decomposed by water with the sejiaration of antimony trioxide oxide. 
 Its salts with strong acids, in which it ai)])ears as a base, are also easily 
 decomposed by water. The normal hydrate, HgSbOg, corresponding to 
 
224 
 
 INORGANIC CHEMISTRY. 
 
 arsenioiis acid, IIjAsOg, may be obtained from ])otassinm antimonyl tar- 
 trate (see Organic Chemistry). Other hydrates are not definitely known. 
 
 The higher oxidation product, the i)entoxide, Sl)/)^, on the contrary, 
 has an acid nature and yields salts with the bases. The hydrate, IljSbtJ^, 
 or ortho-antimonic acid, ])asses at 175° into vieta-antimonic acid, HSbOg, 
 and at 275° into the anhydride, Sb./)^. Most of the antimoniates are 
 derived from pyro-antimonic acid, H^Sb^O,, which is produced on decom- 
 posing antimony pentachloride with water and drying the precii)itate at 
 100° (p. 148). 
 
 Antimony Oxide (antimony trioxide), SbjOj or Sb^Og, is obtained by burning the 
 metal in the air, or by oxidizing it with dilute nitric acid. By sublimation it may be 
 obtained in two different crystal systems, in regular octahedra and in rhombic prisms. In 
 these forms it occurs also in nature as the isometric senarniontite and rhombic valentinite. 
 It is not isodimorphous with arsenic trioxide, which was long considered to be the case. 
 Its vapor density at 1560° corresponds to the formula Sb^Og, and at higher temperatures 
 it is very probable that it will be like arsenic trioxide, the simpler molecule Sb203. d'he 
 hydrate, Sb(OII)3, is thrown out as a white precipitate on adding dilute sulphuric acid 
 to a solution of tartar emetic. It parts with water quite readily and changes to antimony 
 trioxide also called antimonious acid. 
 
 The latter and the hydrate dissolve in sodium and potassium hydroxide, and, very 
 probably, form salts (NaSb02) which decompose upon evaporating the solution. In this 
 beliQvior the acid nature of antimony hydrate is also seen ; therefore it has received the 
 name of antimonious acid. 
 
 The oxide forms salts with acids, which are derived either from the normal hydrate, 
 HjSbO^, or from the hydrate, HSbOj = SbO.OII {ineia-antinionions acid). In the 
 salts of the first kind we have three hydrogen atoms of the hydrate replaced by acid 
 radicals, or, what is the same, a trivalent antimony atom displacing three hydrogen atoms 
 of the acids : 
 
 Sb03fN02)3 or Sb(N03)3. 
 
 Antimony nitrate. 
 
 In the second variety of antimony salts derived from the hydrate, SbO . OH, hydrogen 
 is replaced by a univalent acid residue, or the hydrogen of the acid is substituted by the 
 univalent group, SbO, known as antimonyl : 
 
 SbO.O.NO^ or N03(Sb0). 
 
 Antimonyl nitrate. 
 
 Of these salts may be mentioned the following : 
 
 Antimony Sulphate, Sb2(S04)3, which separates when a solution of the oxide in 
 concentrated sulphuric acid is cooled. 
 
 Antimonyl Sulphate, (Sb0)2S04, is formed when antimony oxide is dissolved in 
 somewhat dilute .sulphuric acid, and crystallizes in fine needles on cooling. Water decom- 
 poses both, formii>g basic salts ; hence the basic nature of antimony oxide is slight. 
 
 Antimonic Acid, H3SbO^, is obtained upon warming antimony with concentrated 
 nitric acid or by adding antimcrny pentachloride to cold water : 
 
 2SbCl5 -f 8II2O = 2ll3SbO, + loHCl. 
 
 It is most conveniently prej)ared by decomposing its potassium salt with nitric acid. It 
 forms, when dried over concentrated .sulphuric acid, a white powder of the composition 
 indicated by the formula. When air-dried it has the formula — 
 
 2n3SbO^ T II 2 O. 
 
 It is almost insoluble in water and in nitric acid, but reddens blue litmus-paper. It is a 
 weak monobasic acid, the .salts oi' which are mostly insoluble in water. [See Beilstein 
 
COMPOUNDS OF ANTIMONY WITH SULPHUR. 
 
 225 
 See also 
 
 and V. Blaese, Cheni. Centralblatt, 1889, i, 803; Ebel, Ber. 22 (1889), 3044. 
 p. 225.] 
 
 At 4(X)° antimonic acid yields antimony pentoxide, Sb.^Oj, a yellow, amorphous 
 mass, soluble in hydrochloric acid. At 800° the pentoxide breaks down into oxygen and 
 the oxide, Sb204, which can be viewed as antimonyl meta-antimoniate (SbOj. SbO), or 
 
 as a mixed anhydride, | O. It is a white powder, becoming yellow when heated 
 and at the melting point of silver is decomposed into oxygen and the trioxide. 
 
 COMPOUNDS OF ANTIMONY WITH SULPHUR. 
 
 These are perfectly analogous to the sulphur compounds of arsenic, and 
 form sulpho salts with alkaline sulphides, corresponding to the oxygen 
 salts. Acids precipitate antimony sulphide and liberate hydrogen sul- 
 phide from the sulpho-salts. 
 
 Antimony Trisulphide, Sb 2 S 3 , is found in nature as stibnite {Stibium 
 sulphuratu77i nigriwi) in radiating crystalline masses of dark-gray color 
 and metallic luster; its specific gravity is 4 . 7 . When heated it melts and 
 distils. The artificial sulphide obtained by precipitating a solution of 
 the oxide with hydrogen sulphide, is an amorphous red powder. When 
 it is digested with aqueous hydrochloric acid or when carefully heated 
 alone in an atmosphere of carbon dioxide it passes into the black, crys- 
 talline variety. Upon warming, the sulphide dissolves in concentrated 
 hydrochloric acid to form antimony trichloride. 
 
 The compound, Sb 2 S 20 , occurring in nature as red stibnite, can be 
 artificially prepared, and serves as a beautiful red pigment, under the 
 name of a77timony ciTiTiabar. Ke7 77ies 77imerale, employed in medicine, is 
 obtained by boiling antimony sulphide with a sodium carbonate solution, 
 and is a mixture of antimony trisulphide and antimony trioxide. 
 
 Antimony Pentasulphide, S^S^, or gold sulphur (Stibiu77t sulphur- 
 atu77i aurantiacuni), is precipitated by hydrogen sulphide from acid solu- 
 tions of antimonic acid ; it is more conveniently obtained by the precipi- 
 tation of sodium sulphantimoniate, Na^SbS^, with hydrochloric acid: 
 
 2Na3SbS, -f 6 IIC 1 = Sb2S5 + 6NaCl + 3H2S. 
 
 It is an orange-red powder, like the trisulphide; it decomposes on 
 being heated into antimony trisulphide and sulphur. It dissolves to 
 antimony trichloride in strong hydrochloric acid, with separation of sul- 
 phur and liberation of hydrogen sulphide. 
 
 Sodiu77i sulpha7tti77io7iiatey NajSbS^ (Schlippe’s salt), results from boiling 
 pulverized antimony trisulphide with sulphur and sodium hydroxide 
 (p. 223 ). Upon concentrating thesolution it crystallizes in large yellow 
 tetrahedra containing nine molecules of water (SbS^Naj -(- pHjO) ; ex- 
 posed to the air it becomes covered with a brown layer of antimony 
 pentasulphide. It serves principally for the preparation of the officinal 
 gold sulphur. 
 
226 
 
 INORGANIC CHEMISTRY. 
 
 VANADIUM. NIOBIUM. TANTALUM. 
 
 V = 5i.2. Nb = 94. Ta = 183. 
 
 The three rare elements, vanadium, niobium and tantalum, are closely related to the 
 phosphorus grou^). They yield derivatives very much like those of the phosphorus 
 group, but possess a more metallic character. 'I'hey exhibit many characteristics similar 
 to those of chromium, iron and tungsten, with which they are fref[uently associated in 
 their naturally occurring compounds (compare the Periodic System of the Elements). 
 
 Vanadium, observed in 1801 by Del Rio and considered to be chromium until Sefstrotn 
 in 1830 proved it to be a new element, occurs in nature principally in the form of salts 
 of vanadic acid (vanadium lead ore). Creuzot has recently worked the 'I'homas slags for 
 vanadic acid, which originally was present in the iron ores. Vanadium may be obtained 
 free by igniting its chlorides in a current of hydrogen. It is a grayish- white, metallic, 
 lustrous powder, of specific gravity 5.5. It is difficultly fusible, and changes in the air 
 very slowly at ordinary temperature. When heated, it burns to vanadium pentoxide, 
 VjOj. It unites readily with nitrogen to form vanadium mononitride, VN. 
 
 Vanadium trichloride, VCI3, forms red plates, which readily deliquesce in the air; 
 it is not volatile. 
 
 Vanadium oxychloride, VOCI3, results on heating a mixture of vanadium trioxide, 
 V,^03, and carbon in chlorine gas. It is a lemon-yellow liquid, of specific gravity 1.84, 
 and boils at 126°. It fumes strongly in the air and decomposes with water (analogous 
 to phosphorus oxychloride) into vanadic acid and hydrochloric acid. Its vapor density 
 corresponds to the formula VOCI3. 
 
 Vanadium trioxide, V.2O3, is a black powder obtained by heating vanadium pentoxide, 
 V2O5, in hydrogen. It combines with oxygen, to form vanadium pentoxide, VgO^. The 
 corresponding sulphate, V2(SO^)3, combines, like the similarly constituted sulphates of 
 aluminium, iron and chromium, with the sulphates of the alkalies to form alums ( see these ). 
 
 Vanadium pentoxide, V or vanadic anhydride, is a brown mass obtained by fusing 
 the naturally occurring vanadates with niter, etc. It exists in three varieties. It is solu- 
 ble in the alkalies, and forms with the metals salts of vaiiadic, Il3V()4, and metavanadic 
 acids, HVO3. All these compounds are similar in constitution to those of the elements 
 of the phosphorus group. In addition to these, vanadium forms other compounds, con- 
 stituted like those of sulphur and chromium. In this class belong VCI2 (dichloride), 
 the tetrachloride, VCl^, vanadious oxide, VO, vanadium dioxide, VO2, and the oxychlo- 
 ride, VOClg. The tetrachloride, V.C1^, is a red-brown liquid, boiling at 154° ; its vapor 
 density corresponds to its formula. 
 
 Niobium, Nb, and tantalum, Ta, are not well known in the free state. They occur 
 together as niobates and tantalates in a few rare minerals — the columbites and tantalites. 
 The chlorides, NbClg and TaCl^, are volatile and are decomposed by water. Niobium 
 and tantalum unite with potassium fluoride, forming double salts, e.g., 2KFl.NbFl5 and 
 2KFl.TaFl5; also 2KFI . NbOFlg and 2KFI . TaOFlg. When potassium niobium fluor- 
 ide, 2KFI. NTFI5, is heated with sodium, niobium hydride, NbH, is formed. This is a 
 grayish-black powder, which when heated burns to niobic anhydride, Nb205, and water. 
 The oxides, Nb205 and form salts of niobic (HgNbO^) and tantalic (HgTaO^) 
 
 acids with bases. 
 
 4. OXYGEN DERIVATIVES OF THE ELEMENTS OF 
 THE CARBON GROUP. 
 
 'File following normal hydroxides correspond to the halogen deriva- 
 tives, CCl^, SiCl^, GeCl^, and SnCl^, of the quadrivalent elements car- 
 bon, silicon, germanium, and tin (see p. 164): 
 
 IV IV 
 
 C(()II)^ Si(OlI)^ 
 
 Normal Normal 
 
 carbonic acid. silicic acid. 
 
 TV 
 
 Oe(OII), 
 
 Normal 
 
 germanic acid. 
 
 IV 
 
 Sn(OH)4. 
 
 Normal 
 stannic acid. 
 
OXYGEN COMPOUNDS OF CARBON. 
 
 227 
 
 These normal hydrates or acids being true ortho-acids have but little 
 stability, and exist chiefly in their derivatives. By the separation of one 
 molecule of water, they pass into — 
 
 CO3H2 SiOgH^ GeOgHj SnOgllg, 
 
 Or 
 
 C0(0H)2 SiO(OH), GeO(OH)2 SnO(OH)2. 
 
 Carbonic acid. Silicic acid. Germanic acid. Stannic acid. 
 
 These hydroxyl derivatives deport themselves toward the normal just 
 as the meta acids of the elements of the nitrogen group do to the ortho- 
 acids (see p. 201). They constitute the ordinary acids of the quadriv- 
 alent elements carbon, silicon, germanium, and tin, and, as they contain 
 two hydroxyl groujDS, are dibasic. 
 
 Carbon is the lowest member of this group, with the least atomic 
 weight. Among the elements of the other three groups corresponding to 
 it are nitrogen, oxygen, and fluorine : 
 
 C = 12 N = 14.04 O = 16 FI = 19. 
 
 Fluorine and oxygen do not afford any oxygen acids. The normal acids 
 of nitrogen, N(OH)5 and N(OH)3, are very unstable, and pass readily 
 into the meta-acids, NO2 . OH and NO . OH. The normal carbonic acid, 
 C(OH)^, corresponds to this, but is not capable of existing. Indeed, the 
 meta- or ordinary carbonic acid, H2CO3, is also very unstable and at once 
 decomposes, when separated from its salts, into water and carbon dioxide, 
 CO2. Even silicic, germanic, and stannic acids break down readily into 
 water and their anhydrides: 
 
 COg SiOg GeOg SnOg. 
 
 Carbon dioxide. Silicon dioxide. Germanium dioxide. Tin dioxide. 
 
 1. OXYGEN COMPOUNDS OF CARBON. 
 
 Carbon Dioxide, CO2, or carbonic anhydride (generally called car- 
 bonic acid), is produced when carbon or its compounds are burned in air 
 or oxygen. It is found free in the air (in 100 volumes, on an average, 
 0.035 volumes CO2), in many mineral springs (acid springs), and escapes 
 in large quantities from the earth in many volcanic districts. It occurs 
 in the liquid form, enclosed in the cavities of many crystalline minerals 
 (quartz, topaz). It is prepared on a large scale by burning coke or by 
 the ignition of limestone; in the laboratory it may be most conveniently 
 obtained by the decomposition of calcium carbonate (marble or chalk) 
 with dilute hydrochloric acid : 
 
 CaCOg + 2IICI = CaCl.2 + COg + ITgO. 
 
 Calcium Calcium 
 
 carbonate. chloride. 
 
 Carbon dioxide is a colorless gas, with a slightly acid taste. Owing 
 to its weight, the gas may be collected by air disj)lacement, and may be 
 
228 
 
 INORGANIC CHEMISTRY. 
 
 l)oured from one vessel into another filled vvitli air. A liter of carbon 
 dioxide weighs i 965 grams at 0°, 760 mm. pressure, and in 45° latitude 
 at sea-level. By a pressure of 50-60 atmospheres at the ordinary tem- 
 jierature carbon dioxide can be licpiefied, as was first shown by h'araday. 
 'I'he apparatus of Thilorier and Natterer were emjjloyed to this end. At 
 l)resent licpiid carbon dioxide is brought into market enclosed in wrought- 
 iron cylinders, and is used (piite regularly in technical operations. 
 
 Carbon dioxide can only be licpicfied below -1-30.9°; this is its criti- 
 cal temperature (p. 47). Its tension (critical pressure) at this point 
 ecpials 77 atmosi)heres. If licpiid carbon dioxide, enclosed in some suit- 
 able vessel, be allowed to escajie into the air by opening a stop-cock 
 (ordinary pressure), it immediately solidifies (see below) to a white, snowy 
 mass. This is because in the eva[)oration of a jiart of the licpiid .so much 
 heat is withdrawn that the remainder becomes solid. vSolid carbon 
 dioxide is a very poor conductor of heat and vajiorizes very slowly. Not- 
 withstanding its low tem])erature it can be handled without serious result, 
 because it is always surrounded by a gaseous layer, and is, consecpiently, 
 not in immediate contact with the skin. If, however, it be jiressed 
 between the fingers, it will produce painful burns. 
 
 The temperature of the solid carbon dioxide vaporizing in the air under 
 ordinary pressure is about — 78° (its boiling point). When the solid 
 dioxide is mixed with a little ether it forms a paste, and then conducts 
 heat better, and is, therefore, well adapted as a cooling agent. In vacuo 
 its temperature diminishes to — 140°. 
 
 Liquid carbon dioxide is a colorless, very mobile liquid. Its specific 
 gravity is 0.91 at — 1.6°, 0.84 at -|-i5°, and 0.726 at -[-22.2°. Its coeffi- 
 cient of expansion is, consequently, greater than that of the gases ; other 
 gases behave similarly, but only such as are condensed under great pressure. 
 If liquid carbon dioxide, contained in a glass tube, be heated, it expands 
 rapicily and suddenly passes, at the critical temperature, -{-30-9°) ii^to 
 gas. This behavior enables us to determine without difficulty whether 
 the liquids contained (see above) in minerals are liquid carbon dioxide. 
 
 If liquid carbon dioxide, confined in a glass tube, be cooled by a mix- 
 ture of solid, snowy dioxide and ether (see above), it will solidify to a 
 transparent ice-like mass, which will melt at — 65°, according to another 
 authority at — 57°. 
 
 The tension of the solid or liquid dioxide, which at the same time indicates the pres- 
 sure necessary for condensation at various temperatures, is given in the following table : 
 
 Temperaturk. 
 
 Tension. 
 
 Temperature. 
 
 Tension. 
 
 4 30.9° 
 
 77 atmos. 
 
 — 20° 
 
 19.9 atmos. 
 
 20° 
 
 58.0 “ 
 
 0 
 
 0 
 
 1 
 
 10.2 “ 
 
 10° 
 
 46.0 “ 
 
 —60° 
 
 3-9 “ 
 
 0° 
 
 35-4 “ 
 
 — 70° 
 
 2.1 “ 
 
 
 
 —80° 
 
 I.O “ 
 
CARBON DIOXIDE. 
 
 229 
 
 At the temperature of fusion of the solid dioxide ( — 65°) the tension equals about 3.5 
 atmospheres ; the resulting liquid has this tension at this temperature. If the external 
 pressure exerted upon it be less, it cannot exist as a liquid, but must immediately pass 
 into the gaseous state. Herein we observe why the solid dioxide (under ordinary pres- 
 sure) does not melt in the air, but vaporizes at once ; and further, it explains why the 
 liquid dioxide, subjected to the ordinary atmospheric pressure, cannot continue in this 
 state — why it is either gasified at once or changed to the solid form. 
 
 Many other fusible solids behave like the dioxide. If the tension of their vapors at 
 the fusion temperature exceeds that of the external atmospheric pressure they do not melt 
 in the air, because the resulting liquid is immediately transformed into vapor ; they 
 vaporize (sublime) directly, without previous fusion. Such bodies are, e. g., arsenic, 
 arsenic trioxide, AsjOg, camphor, mercurous chloride, HgCl, etc. They can be fused only 
 under increased pressure (in sealed tubes). Again, all solids fusible in the air (under 
 ordinary pressure) may be converted directly into gases by removing the external pres- 
 sure. Thus iodine fuses at 114°, but sublimes in a vacuum without previous fusion. 
 Mercuric chloride, HgClj, fuses at 265°, but not if the external pressure be less than 420 
 mm. Water melts at 0°, its tension at this temperature is 4. 6 mm. If the external pres- 
 sure be less {in vacuo), it will no longer melt, but vaporize at once. The pressure 
 below which solids no longer melt has been called Xhexx critical pressure (Carnelley). It 
 is, of course, understood that this is nothing more than the tension of the substance at its 
 point of fusion. 
 
 Water dissolves an equal volume of the gas at 1 4° ; at 0° it takes up i . 79 
 volumes. This proportion remains constant for every pressure, i. e., at 
 every pressure the same volume of gas is absorbed. As gases are con- 
 densed in proportion to the pressure, the quantity of absorbed gas is also 
 proportional to the former (law of Henry and Dalton). Hence i volume of 
 water absorbs, at 14° and 2 atmospheres pressure, 2 volumes, at 3 atmos- 
 pheres 3 volumes, etc., of carbon dioxide — measured at ordinary pressure. 
 The gas absorbed at higher pressure escapes with effervescence of the liquid 
 when the pressure is diminished; upon this depends the sparkling of soda 
 water and champagne, which are saturated with carbon dioxide under 
 high pressure. Every naturally occurring water, especially spring water, 
 holds carbon dioxide in solution, which imparts to it a refreshing taste. 
 
 As the product of a complete combustion carbon dioxide is not com- 
 bustible, and is unable to support the combustion of most bodies; a glim- 
 mering chip is immediately extinguished in it. In a similar manner it is 
 irrespirable. Although it is not poisonous, in the true sense of the 
 word, yet the admixture of a few per cent, of carbon dioxide to the air 
 makes it suffocating, as it retards the elimination of the same gas from the 
 lungs. 
 
 It is decomposed by the continued action of the electric spark into car- 
 bon monoxide (CO) and oxygen ; upon heating to 1300° it suffers a par- 
 tial decomposition into carbon monoxide and oxygen. It is also decom- 
 posed when conducted over heated potassium or sodium, with separation 
 of carbon ; the potassium combines with oxygen to form potassium oxide : 
 
 CO2 -f- 2K2 C -j- 2K2O, 
 
 which forms potassium carbonate (K.^C03), with the excess of carbon di- 
 oxide. Glowing carbon reduces the dioxide to the monoxide (p. 231). 
 Carbon monoxide is analogously formed on conducting a mixture of the di- 
 oxide and hydrogen (equal volumes) through a tube heated to redness, but 
 
230 
 
 INORGANIC CHEMISTRY. 
 
 between 250° and 300° the opposite reaction takes place, although only 
 to a slight degree : 
 
 CO, + n, ^ CO + 11 , 0 . 
 
 The composition of carbon dioxide is readily determined by burning 
 a weighed quantity of pure carbon (diamond or graphite) in a current 
 of oxygen, and ascertaining the weight of the resulting gas. From the 
 formula COj it follows that in one volume of carbon dioxide there is 
 contained an equal volume of oxygen. We may satisfy 
 ourselves of this by burning carbon in a definite volume 
 of oxygen; after cooling, there is obtained an equal 
 volume of carbon dioxide : 
 
 C + O2 = COj. 
 
 I vol. I vol. 
 
 The experiment is most i)ractically executed by aid of 
 the apparatus of Hofmann pictured in Fig. 66. The 
 sphere-sha})ed expansion of the eudiometer limb of the 
 U-tube is closed by means of a glass stopper, through which 
 two copper wires pass. The one wire bears a combustion 
 spoon at its end, upon which lies the carbon to be burned, 
 while the other terminates in a thin piece of platinum, 
 which is in contact with the carbon. For the performance 
 of the experiment, the [J-tube is filled, to near the bulb, 
 with mercury and the air is expelled from the bulb limb 
 by means of a rapid current of oxygen, the stapj)er made 
 air-tight, the mercury level noted, and then pass the in- 
 duction spark between the platinum wire and the copper 
 Fig. 66. spoon; this induces the burning of the carbon. As the 
 volume of the enclosed gas is greatly expanded by the heat 
 developed, it is advisable, in order to avoid the jumping out of the stopper, 
 to previously reduce the pressure of the gas about two-thirds, by running 
 out mercury. When the combustion is finished the mercury which was 
 drawn off is returned and it rapidly assumes its original position. The 
 same apparatus can also be employed for the illustration of the volume 
 relations observed in the combustion of sulphur and other bodies. 
 
 In dry condition, carbon dioxide, like all anhydrides, exhibits neither 
 basic nor acid reaction. In aqueous solution it colors blue litmus-paper 
 a faint red ; upon drying the paper the red disappears, in consequence 
 of the evaporation of the carbon dioxide. 
 
 We may then regard it as probable that free carbonic acid, H^COj, is 
 contained in the aqueous solution, but this readily decomposes into the 
 dioxide, COj, and water. The salts of carbonic acid are produced by 
 the action of carbon dioxide upon the bases : 
 
 2KOII + CO2 = K2CO3 -f HjO. 
 
 Potassium 
 
 carbonate. 
 
 Carbon dioxide is, therefore, easily absorbed by potassium and sodium 
 hydroxide. On conducting it through a solution of calcium or barium 
 
CARBON MONOXIDE. 23 1 
 
 hydroxide, a white precipitate of barium or calcium carbonate, BaCO^ 
 or CaCOg, is produced. 
 
 Carbonic acid is dibasic, forming primary (acid) and secondary (neu- 
 tral) salts, KHCO3 and K2CO3, called carbonates. As the acidity of car- 
 bonic acid is only slight, the secondary salts, obtained from strong bases, 
 exhibit a basic reaction. Most acids expel the weak carbonic acid from 
 its salts, with decomposition into carbon dioxide and water. 
 
 In the reduction of carbon dioxide we meet an interesting transi- 
 tion or transformation of substances of mineral origin into those of ani- 
 mate nature. According to Lieben nascent hydrogen readily reduces 
 carbon dioxide in the form of dissolved bicarbonates to formic acid (for- 
 mates) : 
 
 HO . COONa + = H . COONa + H^O. 
 
 Sodium Sodium 
 
 bicarbonate. formate. 
 
 And the latest researches show that water and hydrogen, under the influ- 
 ence of the silent electric discharge, reduce carbon dioxide to formic 
 acid : 
 
 CO.2 + H2 = H . COOH and CO2 + H2O = H . COOH + O. 
 
 The reduction of carbon dioxide by the chlorophyl granules of plants 
 in the sunlight is ])articularly important. It is by them that the almost 
 endless number of organic bodies ” contained in plants is built up. In 
 the animal organism, on the other hand, carbon dioxide and water are 
 the chief decomposition products of organic compounds ; consequently 
 the exhaled air is rich in this acid. 
 
 Percarbonic acid, H2C.20g, has been obtained by the electrolysis of acid carbonates just 
 as persuiphuric acid was formed from sulphuric acid. It is due to the union of the ions — 
 
 O-CO-OH 
 
 O-CO-Me (Me = metal). The free acid would have the formula ) 
 
 O-CO-OH’ 
 
 and may be said to be derived from hydrogen peroxide (p. 198 ). Its salts will be 
 described later. 
 
 Carbon Monoxide, CO. Carbon dioxide is the first product in the 
 combustion of coal. But just as soon as the combustion reaches a certain 
 temperature — maximum — the excess of carbon reduces the dioxide and 
 carbon monoxide is formed : 
 
 CO2 -f C = 2CO. 
 
 1 vol. 2 vols. 
 
 This is demonstrated by the following experiment : When dry air is con- 
 ducted over heated coal, carbon dioxide is formed almost exclusively 
 from 400° to about 700°. P'rom this point forward carbon monoxide 
 appears in increasing amount and from 1000° upwards it is almost the 
 sole product. Hence glowing coals at moderate temperature burn with- 
 out flame, but at more elevated temperatures (from 1000° upwards) flame 
 is present (p. 155 ). 
 
 Zinc-dust reacts like carbon : 
 
 COj + Zn = ZnO -f CO. 
 
232 
 
 INORGANIC CHEMISTRY. 
 
 When carbon dioxide is conducted through a glass tube, containing zinc-dust heated to a 
 faint red heat, almost pure carbon monoxide escapes. A more convenient procedure 
 consists in heating pulverized magnesium carbonate and zinc-dust in a gla.ss retort, when 
 CO, containing COj, is eliminated ; .subsequently the former alone e.scapes. Pure mon- 
 oxide is al.so formed upon heating zinc-dust with chalk (in Cfjui valent quantities) : 
 
 Zn + CaCOj = ZnO -f CaO -j- CO, 
 a very convenient method for its preparation. 
 
 The monoxide is prodticed, further, by igniting carbon with different 
 metallic oxides, e. g., zinc oxide : 
 
 ZnO -P c = Zn -f CO. 
 
 Water is similarly decomposed. On conducting aqueous vapor over 
 burning carbon there is produced a mixture of carbon dioxide and 
 hydrogen : 
 
 C -f 2 lbO = COj + 2 ll.,. 
 
 By further action of the heated carbon the carbon dioxide in this equa- 
 tion is reduced to carbon monoxide. This gas mixture is known as water 
 gas and is applied technically. [See A. Naumann, Ber. 25 (1892), i, 
 556; also p. 229.] 
 
 A mixture of hydrogen and carbon monoxide is also produced when 
 the electric arc passes between carbon points under water : 
 
 HjO + C = CO -f Hj. 
 
 For the preparation of carbon monoxide, oxalic acid is warmed with 
 sulphuric acid ; the latter withdraws water from the former, and the residue 
 breaks down into carbon dioxide and monoxide : 
 
 H2C2O, = CO2 + CO + H2O. 
 
 The disengaged mixture of gases is conducted through an aqueous 
 solution of sodium hydroxide, by which the carbon dioxide is absorbed, 
 the monoxide passing through unaltered. Pure carbon monoxide may be 
 prepared by heating yellow prussiate of potassium (see Iron) with 9 parts 
 of concentrated sulphuric acid. The resulting gas is conducted through 
 sodium hydroxide to remove from it traces of carbon dioxide and sulphur 
 dioxide. Pure monoxide is also produced when concentrated formic acid 
 or lead formate is heated with concentrated sulphuric acid: 
 
 CO2H2 = CO + H2O. 
 
 A liter of the gas weighs 1.25078 grams under normal conditions; it 
 is therefore 0.9672 as heavy as air. It is one of the gases which are 
 condensed with difficulty. Its critical temperature is — 141° and its 
 critical pressure is 35 atmospheres. Liquid carbon monoxide solidifies 
 under 100 mm. ])ressure at — 207° and under 4 mm. pressure at — 220° 
 (Olszewsky). It boils at — 190° under 760 mm. j)ressure. It is almost 
 insoluble in water, but is readily dissolved by an ammoniacal solution of 
 cui)rous chloride (CU2CI2) with which it forms a crystalline compound, 
 but this decomposes when its solution is heated and carbon monoxide is 
 
CARBON MONOXIDE. 
 
 233 
 
 again liberated. When ignited, it burns in the air, with a faintly lumi- 
 nous, beautiful blue flame, which distinguishes it from other combustible 
 gases. With air or oxygen, it affords (similar to hydrogen) a very ex- 
 plosive mixture : 
 
 2CO + 02 = 2CO2. 
 
 2 vols. I vol. 2 vols. 
 
 The union of carbon monoxide and oxygen takes place at very high 
 temperatures; hence the burning flame of the gas is extinguished upon 
 cooling. A flame or a spark from a powerful induction coil is necessary 
 to ignite a dry mixture of carbon monoxide and oxygen. When the two 
 gases are moist they are more easily ignited and combustible. This is 
 explained by the fact that carbon monoxide unites with the aqueous 
 vapor and yields the dioxide and hydrogen (CO -|- H2O = CO2 -j- H2), 
 which then combines with oxygen and forms water (pp. loi, 230). 
 
 In consequence of its ready oxidation, it is capable of reducing most 
 metallic oxides at a red heat : 
 
 CuO - 4 - CO = Cu -h CO2 . . . + (31.3 Cal.) 
 
 (37.1 Cal.) (28.5 Cal.) (96.9 Cal.) 
 
 Some noble metals are precipitated from solutions of their salts by car- 
 bon monoxide even in the cold. Thus, palladium and gold are thrown 
 out from their chloride solutions by it. A piece of paper moistened with 
 palladious chloride (PdClg) is blackened by it (delicate test for CO). 
 When carbon monoxide is conducted into a platinic chloride solution 
 platinous chloride separates. 
 
 Carbon monoxide is reduced to carbon with difficulty. Burning bodies 
 are extinguished by it. When heated with potassium it is decomposed 
 with separation of carbon. Under the influence of the silent electric 
 discharge hydrogen reduces it to formic aldehyde: 
 
 CO + H2 = H . COH. 
 
 This is the simplest organic compound, consisting of carbon, oxygen and 
 hydrogen. Water under like conditions changes carbon monoxide to 
 formic acid : 
 
 CO + II2O = H.COOH 
 
 (see p. 231). 
 
 When inhaled, even in slight quantity, it acts very poisonously. Blood 
 containing carbon monoxide is characterized by very distinct spectrum 
 reactions. 
 
 In 1890 L. Mond, in conjunction with other chemists, discovered an 
 exceedingly remarkable property of carbon monoxide, viz., that it com- 
 bined with very finely divided nickel at 25-30° to form a liquid, which 
 volatilized readily — nickel carbonyl, Ni(CO)^. Since then it has been 
 observed that carbon monoxide also combines with other metals of the 
 iron group, forming metal carbonyls. 
 
 Nickel Carbonyl, Ni(CO)4, is a colorless, strongly refracting liquid, 
 which boils at 43° (751 mm. pressure) and becomes crystalline at — 25°. 
 
 20 
 
234 
 
 INORGANIC CHEMISTRY. 
 
 Its vaj)C)rs decom))ose witli ex])losi()n at 6o° ; in the air they burn with a 
 very smoky flame [see Z. f. ])hys. (ill. 8 (1891), 150]. 
 
 Carbon monoxide, being an unsaturated compound, combines like 
 ethylene (]). 154) with two atoms of chlorine to Carbonyl Chloride or 
 Phosgene Gas, COClj, the chloride corresponding to carbonic acid : 
 
 CO 4- Cb = COCl,. 
 
 I vol. I vol. I vol. 
 
 It is obtained by bringing together equal volumes of carbon monoxide 
 and chlorine in direct sunlight (hence the name from light, and 
 yevvdio, I produce) ; also when the gases are conducted over ignited plati- 
 num sponge or animal charcoal. It can also be made by conducting 
 monoxide into antimony pentachloride, SbCl^. 
 
 This compound is important in the color industry, and other con- 
 venient methods for its jireiiaration may be found by consulting Erd- 
 mann, Ber. 26 (1893), II, 1990. 
 
 It is a colorless, suffocating gas, the density of which agrees with the 
 molecular formula COCb- Phosgene can easily be liquefied by chilling ; it 
 boils at -|-8° and has the specific gravity 1.43. Water decomposes it into 
 hydrogen chloride and carbon dioxide : 
 
 COCI2 -f HjO = COj + 2HCI. 
 
 Amido-derivatives. — The following are amido-derivatives of carbonic acid which are 
 fully treated in Organic Chemistry and will receive but mere mention here : Carbamic 
 Nil NH 
 
 acid, (^0<Cqj4^» carbamide or urea, They bear the same relation to 
 
 carbonic acid that sulphamic acid and sulphamide sustain to sulphuric acid : 
 
 NfT 
 
 CO<oH^ 
 
 NH 
 SO <-^ 2 
 
 CO< 
 
 NH, 
 
 NH, 
 
 SO 
 
 3 
 
 2 
 
 and can be obtained in a corresponding manner from carbonyl chloride and ammonia. 
 Analogous bodies are derived from hydrazine — the hydrazides, e. g., carbohy dr azide. 
 
 Mention may also be made of nitrogen carbonyl, prepared 
 
 by Curtius and Heidenreich. It is the counterpart of phosgene but is obtained from 
 hydrazoic acid. It is a very explosive, exceedingly volatile, crystalline compound. In 
 this body also the group N3 deports itself like the halogens. [See p. 133 and Ber. 27 
 (1894), 2684; Jr. prakt. Ch. 52 (1895), 454.] 
 
 COMPOUNDS OF CARBON WITH SULPHUR. 
 
 Carbon Bisulphide, CS.^, is formed, like the dioxide, by the direct 
 union of carbon and sulphur; if vapors of the latter are led over ignited 
 carbon the escaping bisulphide vajiors are condensed in a cooled receiver. 
 Pure carbon bisulphide is a colorless, mobile liquid, of faintly ethereal, and 
 in no sense disagreeable, odor; it solidifies at — 116°, and refracts light 
 
CYANOGEN COMPOUNDS. 
 
 235 
 
 strongly. Its specific gravity equals 1.29 at 0°. It is very volatile, boils 
 at 47°, and burns with a blue flame to carbon dioxide and sulphur dioxide. 
 When a mixture of carbon bisulphide vapors and oxygen is ignited, a 
 violent explosion ensues : 
 
 CS, + 3O, = CO, + 2SO,. 
 
 I vol. 3 V'ols. I vol. 2 vols. 
 
 In nitric oxide, the vapors burn with an intensely brilliant flame. On 
 directing a strong current of air upon carbon bisulphide in a porcelain 
 capsule (which conducts heat poorly), so much heat is absorbed by the 
 evaporation that the residual liquid solidifies to a white, snow-like mass 
 which contains water. Carbon bisulphide is insoluble in water ; but mixes, 
 in every proportion, with alcohol and ether. It dissolves iodine with a 
 violet-red color, and serves as an excellent solvent for sulphur, phosphorus, 
 caoutchouc, and the fatty oils. On conducting carbon bisulphide vapors 
 over heated zinc-dust, all the sul})hur unites with the zinc, forming zinc 
 sulphide, while the carbon separates as soot: 
 
 CS, + 2Zn = 2ZnS -(- C. 
 
 Most metals react in a similar manner. 
 
 Carbon bisulphide may be viewed as the anhydride of sulpho carbonic 
 acid, H2CS3. The salts of this acid are obtained by the solution of car- 
 bon bisulphide in alkaline sulphides (see Sulpho-salts, p. 223) : 
 
 CS, T KjS = K^CSj. 
 
 On adding hydrochloric acid to the solutions of these salts the sulpho- 
 carbonic acid separates as a reddish-brown oil. This decomposes 
 readily. 
 
 The sulphur compound corresponding to carbon monoxide is not known : there exists, 
 however, one containing both oxygen and sulphur — carbon oxysulphide, COS. This sub- 
 stance is produced (in small quantity, because it decomposes at about the same tempera- 
 ture) when a mixture of sulphur vapors and carbon monoxide gas is passed through red- 
 hot tubes. 
 
 It is most readily obtained from potassium sulphocyanide, CN. SK (see Organic Chem- 
 istry) by the action of dilute sulphuric acid. Carbon oxysulphide is a colorless gas, with an 
 odor reminding one of hydrogen sulphide. It is present in some sulphur springs. The gas is 
 very readily inflammable and burns with a blue flame to carbon dioxide and sulphurous acid. 
 
 It is soluble in an equal volume of water, decomposing gradually (more rapidly in the 
 presence of alkalies) into the dioxide and hydrogen sulphide : 
 
 COS -f 11,0 = CO, + HS,. 
 
 CYANOGEN COMPOUNDS. 
 
 Of the compounds of carbon mention will also be made of those of cyanogen, as they 
 are of importance in inorganic chemistry. 
 
 Nitrogenous carbon compounds heated with potassium hydroxide yield potassium cya- 
 nide, KCN, which with ferrous oxide forms the so-called yellow prussiate of potassium, 
 K^Fe(CN)p. All the other cyanogen derivatives may be prepared from these two com- 
 pounds, They all contain the group CN, called cyajtogcn. It is similar to the groups 
 OH, Nil,, CH3, and is a univalent radical (pp. 167, 171). In chemical behavior the 
 cyanogen group is very similar to the halogens and the group N3 of hydrazoic acid ; with 
 
236 
 
 INORGANIC CHEMISTRY. 
 
 tlie metals it forms metallic cyanides (KCN, AgCN) very similar to the haloid salts. 
 Hydrogen cyanide is evolved when the cyanides are heated with sulphuric acid : 
 
 2KCN 1 HjSO^ = KjS(\ + 2IICN. 
 
 Hydrogen Cyanide, IICN, is a colorless, mobile liejuid, of peculiar odor, and boiling 
 at 27°. Like the halogen hydrides, it is an acid, forming salts with metals and bases, 
 and is known as hydrocyanic or prussic acid. Loth it and its salts are very powerful 
 poisons. If the CN group is separated from its salts it doubles itself, yielding dicy- 
 anogen or free cyanogen, CjNj (N C-C N), because, like the other univalent groups 
 (as CII3, see p. 168), it cannot exist in a free condition. 
 
 The heats formation of the simplest carbon compounds (from the diamond) above cited 
 correspond with the symbols : 
 
 (C,0)^-26.3 (C0,0) = 68.3 (C, 0 ,) = 94.3 (CO^.Aq) - 5 - 8 . 
 
 If an element combine with another according to multiple proportions, there usually 
 occurs, in the union of the first atom, a greater disengagement of heat than with the fol- 
 lowing atom. The numbers above, on the contrary, show that the union of the second 
 atom of oxygen with carbon (CO,n) sets free 68.3 Cal. ; that of the first atom (C, 0 ), 
 however, only 26.3 Cal. This can only be explained by the fact that, for the vapor- 
 ization and disaggregation of the solid carbon molecules, heat is necessary. If we 
 assume that the direct union of the first atom also disengaged 68.3 Cal., it would 
 follow from this that, in the dissociation of 12 parts by weight of carbon into gaseous free 
 atoms, 42.0 (— 68.3 — 26.3) Cal. were absorbed. 
 
 2. OXYGEN COMPOUNDS OF SILICON. 
 
 Silicon Dioxide, Si02 {^Silicd), the only well-known oxide of silicon, 
 is widely distributed in nature as rock-crystal, quartz, sand, etc. Quartz 
 occurs in nature crystallized in figures of the hexagonal system, with the 
 specific gravity, 2.6; these crystals are colorless, or colored by impurities. 
 
 Quartz is a very widely distributed mineral. The following varieties, distinguished by 
 color and general appearance, exist : Rock-crystal, colorless ; smoky quartz, brown ; 
 morion, almost black ; citrine, yellow ; amethyst, violet, which on heating becomes yel- 
 low. Common quartz is cloudy, slightly transparent, white or colored. It is an essen- 
 tial constituent of granite, gneiss, quartz-porphyry, sandstone and other rocks. Special 
 varieties of it are: fibrous quartz (tiger’s-eye, Africa), the cat’s-eye (Ceylon), etc. The 
 gray hornstone, the green chrysoprase and red or brown jasper consist of compact quartz, 
 etc. 
 
 Tridyinite (rptSv/wt, triplets) has the same composition as quartz and crystallizes in 
 the hexagonal system. It is found in volcanic rocks. Its specific gravity is 2.3. 
 
 Silica may be obtained in an amorphous form by igniting the silicic 
 acid which sei)arates from silicates, but in fusions it has been obtained in 
 both crystalline forms. The natural anhydride always contains impuri- 
 ties; the purest is the colorless rock-crystal. 
 
 Silicon dioxide is insoluble in water and in all acids; but is decom- 
 l)osed by hydrolluoric acid with the formation of silicon fluoride (SiFlJ 
 and water (j). 162). Strong ignition with sodium or ])otassium reduces 
 it in j)art to metallic silicon. The reduction is readily effected by the 
 
OXYGEN COMPOUNDS OF SILICON. 
 
 237 
 
 use of metallic magnesium or aluminium [Gattermann, Ber. 22 (1889), 
 186; Winkler, ibid. 23 (1890), 2652; Jahrb. f. Chem. v (1895), ^^] • 
 
 Si 02 + 2Mg = 2MgO + Si. 
 
 The dioxide prepared artificially dissolves when boiled with potassium or 
 sodium hydroxide ; the natural only when finely divided [Jahrb. f. Chem. 
 VII (1897), 85]. By fusion with the hydroxides or carbonates of the alka- 
 lies all varieties of silica yield a glassy mass (water-glass) soluble in water 
 and containing silicates (K^SiO^ or KjSiOg). Upon the addition of 
 hydrochloric acid to the aqueous solution of the potassium or sodium salt, 
 a very voluminous, gelatinous mass separates; this is probably norjnal 
 or orihosilicic acid, H^SiO^ : 
 
 Na^SiO^ + 4HCI = qNaCl + H^SiO^. 
 
 Orthosilicic acid is also formed in the decomposition of silicon fluoride 
 by water (p. 163). It becomes a white amorphous powder having the 
 composition H2Si03 or 3Si02.2H20 when washed with water and dried 
 in the air. By strong ignition water is gradually expelled until finally 
 the anhydride remains. The freshly precipitated acid is somewhat solu- 
 ble in water, more readily in dilute hydrochloric acid and in caustic soda. 
 On adding a solution of sodium silicate to an excess of dilute hydro- 
 chloric acid the silicic acid remains dissolved. From the hydro- 
 chloric acid and sodium chloride solution we can obtain a perfectly pure 
 aqueous solution of silicic acid by dialysis by proceeding in the fol- 
 lowing manner : 
 
 Pour the hydrochloric acid solution into a wide cylindrical vessel, a. a. whose lower open- 
 ing is covered with animal bladder or parchment paper, and then suspend the vessel (dialy- 
 ser) in another, b, containing pure water (Fig. 67). Osmosis now sets in. The sodium 
 
 Fig. 67. 
 
 chloride and hydrochloric acid pass through the parchment paper into the outer water, b, 
 while on the other hand, water passes from the outer vessel into the dialyser ; the 
 parchment paper is not permeable to silicic acid. This alternate diffusion of the different 
 particles occurs until the outer and inner liquids show the same quantity of diffusible sub- 
 stances. Upon introducing the dialyser into a fresh portion of water, the dialysis com- 
 mences anew. Finally, after repeated renewal of the external water, the dialyser will 
 contain a perfectly pure silicic acid solution, free from sodium chloride and hydrochloric 
 acid. The solution may be concentrated by evaporation ; it then readily passes into a 
 
238 INORGANIC CHEMISTRY. 
 
 gelatinous mass. 1 lie same occurs inslantancously in dilute .solutions if a trace of sorlium 
 carbonate be added or carbon dioxide be led into it. 
 
 Like sodium chloride, all cry.stallizable soluble substances diffu.se through jiarchment. 
 d'hese are known as crystalloids, to distinguish them from the non-dilfusible colloids. 
 To the latter belong gum, gelatine, albumin, starch, glue (koaAo, hence the name col- 
 loid), and especially most of the sub.stances which occur chiefly in vegetable and animal 
 organisms. Like silicic acid these colloids exist in lifjuid, soluble, and .solid gelatinous 
 conditions. Many other substances (like ferric and aluminium oxides) which ordinarily 
 are insoluble in water, can be brought into acjueous solution by dialysis. 
 
 For “osmotic pressure” and the determination of molecular weights 
 see Solutions. 
 
 We have already seen that acids like sulphuric, jthosphoric, and arsenic, 
 are capable of forming anhydro- or poly-acids by the union of several 
 molecules and the elimination of water (j). 214). Silicic acid is particu- 
 larly inclined to that form of condensation. It forms a large number 
 of poly-silicic acids, Si./)3(OH)2, Si30^(0H)^, Si^CJ^OH),, etc., derived 
 from the normal and ordinary acid, according to the common formula: 
 
 ;;/Si(OII)^ — 
 
 These poly-acids are not known free ; it appears, however, that many 
 amorphous forms of silica occurring in nature, as agate, chalcedony, 
 opal, which lose 5-15 i)er cent, of water by ignition, represent such poly- 
 acids. The natural silicates are the salts of such acids. Agate consists 
 of numerous layers of variously colored chalcedony, quartz and 
 amethyst. The majority are derived from the acids H2Si205, H^SigOg, 
 H2Si30^, H^Si^Ojo, and others. Only a few silicates are obtained from the 
 normal acid, e. g., chrysolite, Mg2Si04. 
 
 Corresponding to carbon bisulphide, CS2, is 
 
 Silicon Disulphide, SiS2, which may be made by heating amorphous silicon with 
 sulphur, or by conducting sulphur vapors over an ignited mass of silica and carbon. It sub- 
 limes in shining, silky needles, which water changes to silicic acid and hydrogen sulphide. 
 
 Germanium, Tin and Lead belong to the group of carbon and sili- 
 con. They yield the derivatives MeOg and Me(OH)^ (pp. 226, 227), but 
 as they form lower oxides, GeO and SnO, which manifest a perfectly 
 basic nature and unite with acids to form salts, and as these elements are 
 thereby allied to the metals, they will be discussed under that division. 
 
 TITANIUM. ZIRCONIUM. THORIUM. 
 
 Ti = 48.1. Zr = 90.6. Th = 232. 
 
 The same relation that vanadium, niobium, and tantalum show to the elements of the 
 I)ho.sj)horus group is manifested by the three elements titanium, zirconium and thorium 
 for the silicon group (.see Periodic System, p. 246). 
 
 P = 31.0 V = 51.2 Si = 28.4 Ti = 48.1 
 
 As — 75 Nb = 94 (le =72 Zr = 90.6 
 
 Sb 120 Ta = 183 Sn = 118.5 Th = 232. 
 
TITANIUM — ZIRCONIUM. 
 
 239 
 
 In all their deportments they strongly resemble tin ; they possess, however, a more metal- 
 lic character in their derivatives. They are quadrivalent, affording compounds of the 
 form MeX^, in which X represents univalent elements and groups ; those of the form 
 MeX2, corresponding to the stannous derivatives, are unknown. The hydroxides, 
 Me(OH)^ and MeO(OHj2, have a stronger basic nature than stannic acid and form stable 
 salts with acids ; the basicity increases successively with the atomic weights. Corre- 
 sponding to this, the acidity of the hydrates, i. e., their capability of exchanging hydrogen 
 for metals, gradually diminishes. Thorium hydroxide, Th(OH)^, is not able to form 
 metallic salts. 
 
 TITANIUM. 
 
 Ti = 48.1. 
 
 This metal occurs in nature as titanium dioxide (rutile, anatase, brookite) and in titan- 
 ates (perofskite, CaTiOg, menaccanite, FeTiOg). Free titanium is a gray, metallic 
 powder, obtained by Berzelius on heating potassium titanium fluoride (K2TiFlg) with 
 potassium. Wohler, however, was the first to recognize it as metal. It burns when 
 heated in the air, and decomposes water on boiling. It dissolves in dilute hydrochloric 
 and sulphuric acids, with evolution of hydrogen. 
 
 Titanium Chloride, TiCl^, is formed, like silicon chloride, by conducting chlorine 
 over an ignited mixture of the dioxide and carbon. A colorless liquid, of specific gravity 
 1.76, fuming strongly in the air (with decomposition into hydrochloric and titanic acids), 
 and boiling at 136°. It solidifies at — 25°. The vapor density corresponds to the molec- 
 ular formula TiCl^. It behaves like tin tetrachloride with water. A compound — hydro- 
 fliiotitanic acid, HgTiP’g — and its salts are known. They correspond to hydrofluosilicic 
 acid. 
 
 Titanic Acid, H^TiO^ (first obtained pure by H. Rose), separates as a white, amor- 
 phous powder, on adding ammonium hydroxide to the hydrochloric acid solution of the 
 titanates. When dried over sulphuric acid it loses i molecule of water and becomes 
 TiO(OH)2. Titanic acid, like silicic and stannic acids, forms poly-acids. The hydrates 
 dissolve in alkalies and strong acids, to form salts. On igniting the hydroxides we get 
 
 Titanium Dioxide, Ti02, which may be procured crystallized as rutile, brookite, and 
 anatase. When ignited in a stream of hydrogen it changes to the oxide '1^20,. Titanium 
 dioxide is almost insoluble in acids ; it is only dissolved by hydrofluoric acid. It forms 
 titanates upon fusion with the alkalies. 
 
 The hydroxides, TiO^H^, Ti03H2, etc., conduct themselves as feeble bases with strong 
 acids and afford salts with them (<?. g., TiO. SO4), which are decomposed by water. The 
 alkaline titanates (K2TiOg) are very unstable. Other titanates occur in nature, e. g., 
 CaTiOg, MgTiOg, and the so-called titanic iron, FeTiOg. 
 
 Titanium also forms derivatives after the types, Ti.^Og and TiO, c. g., TigClg and Ti2Cl4. 
 The sesquioxide compounds are usually green or violet in color, while those of the mon- 
 oxide form are black or brown. 
 
 Titanium yields various compounds with nitrogen. When the dioxide is heated in 
 ammonia gas, a dark-violet powder of the composition TiN2 results. The compound 
 TigCN^ — the so-called cyan-titanium nitride — is sometimes found in copper-red, metallic 
 cubes, in blast-furnace slag, when iron ores, containing titanium, have been fused. 
 
 ZIRCONIUM. 
 
 Zr = 90.6. 
 
 Zirconium is very rare in nature, but is generally found in silicates, and especially as 
 zircon, ZrSiO^. Free zirconium, like carbon, is known in three allotropic forms. On 
 fusing potassium zirconium fluoride with aluminium L. Troost obtained it in brilliant, 
 hard, steel-gray, fiat leaflets, of specific gravity 4.15, which burned in the oxyhydrogen 
 flame to the dioxide. L. Troost also i.solated the graphitic form. The amorphous variety 
 was observed by Berzelius, who was the first chemist after Klaproth, its discoverer, to 
 engage in a study of zirconium. lie obtained it by the method used for titanium and 
 got it as a black powder which burnt readily. Zirconium tetrachloride, ZrCl^, and fluoride, 
 ZrFl^, are very similar to the corre.sponding titanium compounds. 
 
 Zirconic acid, Zr(OH)^, is not definitely known. Metazirconic acid, ZrO(OH)2, is 
 
240 
 
 INORGANIC CHEMISTRY. 
 
 [)recipitatecl on adding ammonia water to acid solutions, in the form of a white, volumi- 
 nous mass, which yields zirconium dioxide, Zr( ).^, u|)on ignition. A colorless, transparent 
 crystalline modification is known — isomorphous with rutile and cassiterite. Metazircotiic 
 acid is insoluble in caustic potash and soda ; it only forms zirconates, Na^Zrt )., and Na^Zrt 
 when fused with alkalies and alkaline carbonates. Water decomjjoses tiiese salts witii 
 the liberation of alkali, h'reshly precipitated metazirconic acid, ZrO(On)2, dissolves 
 readily in strong acids. Zirconium dioxide is soluble in hot, concentrated sulphuric acid, 
 forming Zr(S()^)2, which can also be obtained from a(jueous solution in crystals, 
 
 Zr(SOj2.4lI.p. . 
 
 In late years zirconium dioxide has become of great imi)ortance for illuminating pur- 
 poses. It is, therefore, i)rei)ared in large amounts from zircon. 'I'he finely pulverized 
 mineral is mixed with carbon and while being gently heated chlorine is conducted over 
 the mixture when silicon chloride, SiCl^, and zirconium chloride, ZrCl^, are produced ; or 
 the mineral is decomposed with hydrofluoric acid. IJerzelius noticed the light-emission 
 power of the oxide when it was heated. Caron (1868) determined its value for the tem- 
 perature of the oxyhydrogen flame and d'essie du Mothay endeavored to utilize it practi- 
 cally. Linnemann (1885) made the oxide into thin plates and soon thereafter W. Kochs 
 succeeded by a peculiar process in shaping it into any desired form, d'he durability of 
 the zirconia light-giving bodies in the oxyhydrogen flame, the fact that they do not decom- 
 pose in the air, that the light which they diffuse is intensely white (its continuous spec- 
 trum includes the P'raunhofer lines from A to H) render it superior to the Drummond 
 lime light. It is extensively used for projection lamps, for photography and in optical 
 researches. 
 
 THORIUM. 
 
 Th = 232. 
 
 Until recently thorium was only known as a constituent of certain rare northern min- 
 erals. Berzelius first discovered it in the thorite from Arendal, later Wohler found it in 
 pyrochlore, and Kersten in monazite. Small amounts of it are present in many orthites. 
 Its application to illuminating purposes has created a demand for large quantities of it 
 and at present it is isolated from a rich monazite deposit in McDowell County, N. C. 
 That mineral contains 5-6 per cent, of thorium oxide. 
 
 Free thorium has been obtained as a gray, crystalline powder of specific gravity i i.o by 
 reduciirg potassium thorium fluoride or potassium thorium chloride with metallic sodium 
 or potassium. Its specific heat is 0.0276. It burns, when heated in the air, to dioxide, 
 thoria, Th02. The metal is easily soluble in concentrated sulphuric acid, nitric acid, and 
 hydrochloric acid. It is not soluble in the alkalies. The tetrachloride, ThCl^, results 
 from the action of hydrogen chloride upon metallic thorium ; it melts at a white heat and 
 sublimes in white needles. Its vapor density corresponds to the formula ThCl^. It is 
 soluble in water and separates in crystals of the formula ThCl^ . 8H2O. It forms crystal- 
 line double salts with the alkaline chlorides. The fluoride, ThFl^, is a white powder, 
 insoluble in water. The hydroxide, Th(0H)4, is precipitated as a white jelly on adding 
 ammonia water to thorium salts. The hydroxide, ThO(OH)2, is not known. On heat- 
 ing thorium hydroxide, Th (011)4, the dioxide, ThOg, is obtained. The latter is not 
 soluble in dilute acids ; it is, however, converted into thorium sulphate, Th(S04)2, when 
 acted upon with concentrated sulphuric acid. The sulphate dissolves readily in cold 
 water and crystallizes from the same in colorless, brilliant needles of the formula 
 '141(804)2. When its aqueous solution is heated almost all of the sulphate sepa- 
 rates as a woolly mass, which on cooling, slowly dissolves. The hydrate and dioxide do 
 not form .salts with the alkalies. The i.somorphism of thorium dioxide with uranium 
 dioxide, and thorium .sulphate with uranium sulphate is very interesting. 
 
 Thorium dioxide is even more valuable for lighting purposes than zirconium dioxide 
 because it gives forth a blinding light at the temperature of the Bun.sen flame — light rays 
 of every refrangibility are emitted. Auer von Websbach has utilized this property in a 
 gas-light. Fine-meshed material of cotton or linen .saturated with a .solution of thorium 
 nitrate and a little cerium nitrate, surrounds the colorless flame of a Bunsen burner. 
 When the vegetable material has burned away a mantel of thorium dioxide and cerium 
 oxide remains, which has the exact form of the original substance, and emits a bright, 
 clear, soft light. 'I'he ordinary “ gh)w-mantle.s ” contain about 98-99 percent of thoria 
 
BORON. 
 
 241 
 
 and 2-1 per cent, cerium oxide ; it is exactly this mixture which affords the desired white 
 light. The intense emission power Bunte believes is due to the catalytic action of finely 
 divided cerium oxide, which causes a rapid combustion and consequently a high tem- 
 perature. Hence in the Welsbach mantle the thorium oxide simply acts as a stable, 
 porous carrier for the cerium oxide. 
 
 BORON. 
 
 B = II. 
 
 This element is generally classed with the metalloids, and stands 
 isolated among them ; it forms the transition from these to the metals, 
 which is manifest from its position in the periodic system. On the one 
 side, especially when free, it resembles carbon and silicon ; on the other, 
 it approaches the metals, beryllium, aluminium, and scandium (see the 
 Periodic System of the Elements). As recently observed, it affords a 
 gaseous hydride, but it is not very stable, and, like stibine, may be easily 
 decomposed into its constituents. Its oxide, B.^03, although really of an 
 acid nature, approaches such metallic oxides as aluminium oxide, AI2O3, 
 which functionates both as base and acid. Boron is trivalent, and yields 
 only compounds of the form BX3. 
 
 It is found in nature as boracic acid and in the form of borates, like 
 borax (sodium salt), boracite (2Mg3BgOj5 -)- MgCl2 ; Stassfurt). Boron 
 exists in two allotropic modifications: amorphous and crystalline. 
 
 Mois.san made aynorphotts boron by igniting a mixture of freshly fused boric acid with 
 magnesium powder, free from iron. By a rather circuitous treatment of the fusion the 
 boron (with 2 per cent, impurities) is obtained as a black powder of specific gravity 2.45. 
 The product obtained heretofore by other methods, and believed to be amorphous boron, 
 was a mixture of boron, iron boride, boron nitride, etc. The boron of Moissan ignites 
 at 700° in the air ; at higher temperatures sulphur, chlorine, bromine, nitrogen, silver 
 and platinum combine readily with it. It has a great affinity for oxygen, hence acts 
 as a powerful reducing agent. When rubbed together with lead dioxide explosion 
 occurs. Oxygen acids are comparatively easily reduced by boron ; potassium perman- 
 ganate, silver nitrate and ferric chloride suffer this change in the cold. 
 
 The crystalline variety may be obtained by igniting boron trioxide with aluminium. 
 The boron, separated by the aluminium, dissolves in the excess of the latter, and crystallizes 
 from it on cooling ; upon dissolving the aluminium in hydrochloric acid the boron remains 
 in shining, transparent, quadratic crystals, which are more or less colored, and have a 
 specific gravity of 2.63. The crystals are not pure boron, but contain aluminium and car- 
 bon. In their luster, refraction of light, and hardness, they resemble the diamond. 
 Cry.stalline boron is more stable than the amorphous ; it does not oxidize upon ignition, 
 and is only slightly attacked by acids. Fused with potassium and sodium hydroxide both 
 modifications yield borates. 
 
 Boron Hydride, probably BH3, is only known mixed with hydrogen. It results 
 when hydrochloric acid acts upon magne.sium boride. A solid boron hydride is produced 
 on heating boron trioxide with sodium, or borax with magnesium pow^der. It has not 
 been obtained pure. 
 
 Boron Trichloride, BCI3, may be prepared by heating boron in 
 chlorine, or conducting a stream of the latter over an ignited mixture of 
 the trioxide and carbon (see Silicon and Aluminium Chlorides) : 
 
 + 3C + — -BCL -P 3CO. 
 
 21 
 
242 
 
 inor(;anic chemistry. 
 
 It is a colorless liquid, of specific gravity 1.35, and boiling at 17°. Its 
 vapor density corresponds to the molecular formula BCl,. The liquid 
 fumes strongly in the air and is decomposed by water into boric and 
 hydrochloric acids : 
 
 I3CI3 + 311,0 = B(OII)3-f3lICl. 
 
 The trichloride also results from the action of the pentachloride of 
 phosphorus upon boron trioxide : 
 
 b-Ps + 3 bCb = 2BCI3 + 3POCI3. 
 
 Boron Fluoride, BFI3, is similar to silicon fluoride, and is produced 
 according to the same methods, by the action of hydrofluoric acid upon 
 the trioxide, or by warming a mixture of the trioxide and calcium fluo- 
 ride with sulphuric acid : 
 
 B,03 + 3CaFl, + 3lI,SO, = 3CaSO, + 311,0 + 2BFI3. 
 
 It is a colorless gas, fuming strongly in the air, of specific gravity 66 
 (O, = 32), and may be condensed to a liquid under strong pressure. It 
 dissolves very readily in water (700 volumes in i volume), producing 
 Hydrogen Borofluoride, BFl^H (::r= BFI3. HFl), which remains in 
 solution : 
 
 4BFI3 -f 311,0 = 3HBFh + H3BO3. 
 
 The reaction is analogous to the formation of hydrofluosilicic acid 
 from silicon fluoride (see p. 163). Hydrogen borofluoride is a monobasic 
 acid, only known in its solution and in its salts. 
 
 Boric Acid, H3BO3 = B(OH)3, occurs free in nature and in 
 salts. In some volcanic districts, especially in Tuscany, the steam 
 escaping from the earth (fumeroles, etc.) contains small quantities 
 of it. These vapors condense in small natural pools, or are conducted 
 into walled basins. By concentration of the aqueous solution, boric 
 acid separates. To prepare pure boric acid, precipitate a hot solu- 
 tion of borax with nitric acid. The acid seiiarates in colorless, shining 
 scales ; it dissolves in 25 parts of water at 14°, or in 3 parts at 100°. The 
 solution shows a feeble acid reaction with litmus; turmeric paper, 
 moistened with it, is colored red-brown after drying. On boiling the 
 solution, boric acid escapes with the steam. An alcoholic solution 
 of the acid burns with a green flame. These reactions afford a ready 
 means for its detection. 
 
 When heated to 100°, the acid loses one molecule of water, and passes 
 into the anhydro- or meta-acid, HBO,, which at 140° is converted into 
 tetraboric acid, H,B^O^. When ignited, boric anhydride or Boron tri- 
 oxide, B.,03, is produced. This is a fusible, glassy mass, of specific 
 gravity 1.8, and is slightly volatile at a very high heat. Water dissolves 
 the anhydride to boric acid. 
 
 It is a very weak acid ; and can be ex])clled from its salts by most 
 other acids. By fusion it removes most acids from their salts, in con- 
 sc(|uence of the difficult volatility of its anhydride. 
 
 Salts of normal boric acid, B(OII)3, are not known, while the ethers, 
 
PERIODIC SYSTEM OF THE ELEMENTS. 
 
 243 
 
 B( 0 . CH3)3, are. The salts of metaboric acid, BO . OK, can be obtained 
 crystallized, but they are very unstable. They are decomposed by car- 
 bon dioxide with production of salts of tetraboric acid : 
 
 4NaB02 -b CO 2 = Na^B.O^ + Na 2 C 03 . 
 
 The latter, from which the ordinary borates are derived (see Borax), 
 may be viewed as an anhydro-acid, produced by the union of four mole- 
 cules of trihydric boric acid (compare p. 238): 
 
 4B(0H)3-5H20 = H2B,0,. 
 
 On heating amorphous boron in a stream of nitrogen or ammonia, or by igniting a 
 mixture of the trioxide and carbon in nitrogen gas, there is formed Boron nitride^ BN. 
 This is a white amorphous powder, which gives forth an extremely intense greenish-white 
 light when heated in a gas flame. Boric acid and ammonia result when steam at 200° 
 is conducted over the nitride : 
 
 BN + 3H2O = B(0H)3 4- NH3. 
 
 In the heat of the electric furnace carbon unites with boron to Boron carbide, BgC, 
 which is very similar to carborundum, but is harder than the latter. It consists of black, 
 brilliant crystals, of specific gravity 2.5. It is exceedingly stable (Moissan). 
 
 PERIODIC SYSTEM OF THE ELEMENTS. 
 
 In the preceding pages we have studied the non metals and their com- 
 pounds with hydrogen, the halogens and oxygen. In their entire 
 behavior they arrange themselves in four natural groups, the members of 
 which show unmistakable family similarities. It is only in the case of 
 boron and of hydrogen and the new constituents of air that allied mem- 
 bers have not been found. Attention has been repeatedly called to the 
 gradation in the similarity of the group members, which seems to increase 
 with the rise of the atomic weight, and also to the relations of the various 
 groups to one another — relations observed among the metals. All these 
 points appear more striking and more regularly, if the elements be con- 
 sidered in the arrangement which the periodic system assigns them. 
 
 From the time the first atomic weight determinations were made, regu- 
 larities between the atomic weights and the properties of the elements 
 were thought to exist. One of the first chemists to surmise this was J. 
 W. Ddbereiner, whose views were clear and definite in contrast to the 
 fantastic, mathematical sophistries of many of his successors, especially 
 those of the Frenchman Chancourtois, who recently has been unjustly 
 heralded as the forerunner of D. Mendelejeff and Lothar Meyer, the 
 founders of the periodic system.* 
 
 *See Dobereiner, Pogg. Ann. 15 (1829), 301 ; Lothar Meyer, Ann. Chem. Suppl. 
 7 (1871), 354; Mendelejeff, ibid., Supjjl. 8 (1872), 133; re])rinted in No.s. 66 and 
 68 of Ostwald’s Klas.siker der exakten Wissen.schaften. For tlie history of tlie sys- 
 tem see Mendelejeff, Princiifles of Chemistry (1891), 683, 692; L. Meyer, Modern 
 Theories of Chemistry (1883) ; also K. Seubert, Z. f. anorg. Chem. 9 (1895), 334. 
 
244 
 
 INORGANIC CHEMISTRY. 
 
 Arranging the elements according to increasing atomic weight we 
 observe that similar elements return after definite intervals, 'rims they 
 arrange themselves in several })eriods: 
 
 
 
 
 
 I. 
 
 I.i 
 
 P)e 
 
 
 P C 
 
 N 
 
 0 
 
 LI 
 
 
 
 
 
 
 
 2. 
 
 Na 
 
 Mg 
 
 A 1 Si 
 
 P 
 
 S 
 
 Cl 
 
 
 
 3- 
 
 K 
 
 Ca 
 
 Sc 
 
 Ti V 
 
 Cr 
 
 Mm 
 
 be 
 
 Ni Co 
 
 Cu 
 
 Zn 
 
 Ca (ie 
 
 As 
 
 Se 
 
 4- 
 
 Rb 
 
 Sr 
 
 Y 
 
 Zr Nb 
 
 Mo 
 
 — 
 
 Ru 
 
 Rh Pd 
 
 Ag 
 
 Cd 
 
 1 11 Sn 
 
 Sb 
 
 d'e 
 
 5- 
 
 Cs 
 
 Ba 
 
 La 
 
 (Ce Nd 
 
 Pr) 
 
 — 
 
 — 
 
 
 
 
 
 
 6. 
 
 — 
 
 — 
 
 Yb 
 
 — Ta 
 
 W 
 
 — 
 
 Os 
 
 Ir Pt 
 
 All 
 
 Hg 
 
 Tl Pb 
 
 Pi 
 
 
 
 7- 
 
 
 
 
 
 
 
 
 
 
 
 
 
 Hydrogen has no family; it alone would form the first period, and is 
 therefore omitted, as are also helium, argon, and the other recently dis- 
 covered constituents of the air, because their chemical nature is not 
 known; the positions of tellurium and iodine have also been exchanged, 
 on the presumption that the atomic weight assigned tellurium at present 
 is too high. 
 
 The period syste?n, or the law of periodicity, which is indicated by this 
 arrangement, declares that the properties of the ele^nents are periodic func- 
 tions of their atomic weights. 
 
 The first two series, lithium (Li) to fluorine (FI), and sodium (Na) to 
 chlorine (Cl), present two periods of seven members each, in which the 
 corresponding (above and below) members exhibit a great but not com- 
 plete analogy. Sodium resembles lithium ; magnesium, beryllium; chlo- 
 rine, fluorine, etc. Then follow two periods, consisting of seventeen ele- 
 ments each : potassium (K) to bromine (Br), and rubidium (Rb) to iodine 
 (I). The series 5 and bare incomplete, and together probably constitute a 
 period. In the seventh series there are as yet but two elements : thorium 
 = 232 and uranium == 239.5. Thus result three great periods, whose cor- 
 responding members exhibit an almost complete analogy: the elements 
 K, Rb, Cs; Ca, Sr, Ba; Ga, In, T 1 ; As, Sb, Bi, etc., are so similar that 
 they remind us of the homologous series of the carbon compounds (com- 
 pare p. 153), and, therefore, can be designated as homologous elements. 
 It is only in the third great period (series 5 and 6) that the middle 
 members exhibit any variations. 
 
 Now on comparing the three great periods with the two small ones, we 
 discover that the first members are analogous to each other; K, Rb, Cs, 
 resemble Na and la ; Ca, Sr, Ba, resemble Mg and Be. Then the simi- 
 larity gradually lessens, disappears apparently in the middle members, 
 and only aj)j)ears again toward the end of the periods: I and Br resemble 
 chlorine and fluorine; Teand Se, sul|)hur and oxygen ; Bi, Sb, As, phos- 
 phorus and nitrogen, etc. The character or the function of the three 
 great periods is therefore other than that of the two small periods. But 
 in all five periods we can detect a gradual, regular alteration in the prop- 
 erties of the adjoining heterologous This is particularly mani- 
 
 fest in the measurable physical jiroperties, all of which show a maximum 
 or minimum in the middle of the periods (of both the great and the 
 small ). 
 
 'I'he ‘^ame regularity exhibits itself even in chemical properties, in the 
 
PERIODIC SYSTEM OF THE ELEMENTS. 
 
 245 
 
 two small ])eriods, esj)ecially in the valence of the elements in their com- 
 pounds with hydrogen or the hydrocarbon groups CH.^, etc. (desig- 
 
 nated by R; compare ]). 247). The hydrogen valence rises and falls 
 periodically with the specific gravity : 
 
 I II III IV III II I 
 
 NaR MgR2 AIR3 Silb bHg SII^ CIH. 
 
 On the other hand, the maximum valence of the elements with refer- 
 ence to oxygen increases gradually if chlorine be regarded as septivalent 
 in perchloric acid (pp. 171, 173) : 
 
 I IT III IV V VI VII 
 
 Na 20 MgO AI2O3 Si 02 SO3 
 
 The chemical valence expresses itself somewhat differently in the three 
 great periods. In them we have a double periodicity ; thus, e. g., with the 
 salt-forming oxides: * 
 
 I II III IV V VI VII 
 
 K2O CaO SC2O3 Ti02 V2O5 Cr03 Mn207 
 
 II II III IV V VI VII 
 
 CuO ZnO Ga203 Ge02 AS2O5 (8003) (Br20d. 
 
 In consequence of this double periodicity, the first seven and the last 
 seven members of the two great periods, with respect to their valence 
 (and consequently also their compounds), resemble the seven members 
 of the two small periods. To bring out this double periodicity and 
 analogy, the first seven and last seven members of the great periods are 
 divided into two series, and arranged under the corresponding seven 
 members of the small periods. 
 
 In this way the three middle members of the great periods (which are 
 found between the dotted lines of the table, p. 244) come to stand apart, 
 as they have no analogues. In this manner arises the following table, in 
 which the seven (or ten) vertical columns include homologous elements : 
 
 Li 
 
 Be 
 
 B 
 
 C 
 
 N 
 
 0 
 
 FI 
 
 
 Na 
 
 Mg 
 
 A 1 
 
 Si 
 
 P 
 
 S 
 
 Cl 
 
 
 K 
 
 Ca 
 
 Sc 
 
 Ti 
 
 V 
 
 Cr 
 
 Mn 
 
 Fe Ni CO 
 
 Cu 
 
 Zn 
 
 Ga 
 
 1 Ge 
 
 As 
 
 Se 
 
 Br 
 
 
 Rb 
 
 Sr 
 
 Y 
 
 Zr 
 
 Nb 
 
 Mo 
 
 — 
 
 Ru Rh Pd 
 
 
 Cd 
 
 In 
 
 Sn 
 
 Sb 
 
 Te 
 
 I 
 
 
 Cs 
 
 Ba 
 
 La 
 
 Ce 
 
 (Nd,Pr) 
 
 — 
 
 — 
 
 
 — 
 
 — 
 
 Yb 
 
 — 
 
 Ta 
 
 W 
 
 — 
 
 Os Ir Pt 
 
 Au 
 
 Hg 
 
 T 1 
 
 Bi 
 
 Pb 
 
 — 
 
 — 
 
 
 In the table (on p. 246) we have presented the same grouping of the 
 elements, together with their atomic weights, given in round numbers. 
 
 When the periodic grouping of the elements was first presented, the atomic weight of 
 indium and that of uranium had to be doubled, if they were to occupy the places indi- 
 
 *The elements are given here with their highest oxygen valence. Perbromic anhy- 
 dride, Pr20-, and selenic anhydride, SeO.,, have not been prepared, although some of 
 their derivatives are known. Iron, cobalt and nickel have been omitted because their 
 highest oxides are not well defined. 
 
The Periodic Systef?i of the Elefne^its. 
 
 246 
 
 INORGANIC CHEMISTRY 
 
 VIII 
 
 Group. 
 
 ISIO 
 
 
 On 
 
 \r^ 
 
 0 
 
 u 
 
 0 
 
 tr> 
 
 1 : 
 p^ 
 
 
 0 " 
 
 'Pa 
 
 Ov 
 
 VT) 
 
 Ya 
 
 Rh 103 
 
 1 - 
 }m 
 
 
 0 
 
 Pa 
 
 MD 
 
 vr> 
 
 0 
 
 Ph 
 
 Ru 102 
 
 Os 191 
 
 
 \TI 
 
 Group. 
 
 
 On 
 
 E 
 
 Cl 35-4 
 
 0 
 
 CO 
 vri ^ 
 
 0 ^ 
 
 r 
 
 1 1 
 
 0 
 
 \r% 
 
 C^ 1 
 
 1 
 
 \T 
 
 Group. 
 
 e’o” 
 
 IS 
 
 VO 
 
 0 
 
 N 
 
 ro 
 
 m 
 
 vr> 
 
 Mo 96 
 (?)Te 127 
 
 1 
 
 1 T 
 
 CO 
 
 1 
 
 Ov 
 
 rn 
 
 cs 
 
 Li 
 
 V ’ 
 
 Group. 
 
 
 
 ro 
 
 pH 
 
 t/) 
 
 p 
 
 0 
 
 (S 
 
 'i- 
 
 CV^Q 
 ^ C/2 
 
 Y 
 
 00 _ 
 
 r E w 
 
 1 
 
 IV 
 
 Group. 
 
 ^6 
 
 PP 
 
 u 
 
 00 
 
 M 
 
 ’(h 
 
 r) 
 
 00 
 
 H 
 
 Zr 90 
 
 Sn 1 18 
 
 ~"n 1 0 
 
 0 
 
 E ^ 
 
 OJ 1 
 
 U 1 
 
 cs 
 
 rn 
 
 cs 
 
 H 
 
 1 
 
 II III 
 
 Group. Group. 
 
 CO 
 
 pq 
 
 < 
 
 0 
 
 -tO 
 
 0 
 
 m 
 
 Y 89 
 
 In 1 14 
 
 1 La 138 
 
 Yb 173 
 
 T 1 204 
 
 - '-1 
 
 li 
 
 Ov 
 
 o 
 
 pq 
 
 '+ 
 
 rt 
 
 b/) 
 
 S 
 
 i 
 
 Ca 40 
 
 1 Zn 65 
 
 Sr 87 
 
 Cd 112 
 
 Ba 137 
 
 Hg 200 
 
 - '-i 
 
 I 
 
 Group. 
 
 iS* 
 
 E 
 
 ro 
 
 N 
 
 vO 
 
 OVr^ 
 
 ro^ 
 
 E 
 
 Rb 85 
 
 Ag 108 
 
 Cs 133 
 
 Au 197 
 
 1 
 
 
 1 
 
 H-Compounds. 
 Highest salt-forming 
 oxides. 
 
 t/i ^ "5 -5 'B 'B'B'B 
 
 .U „ N tovO 00 0 
 
 <v 
 c/) 
 
 t/5 
 
 ^ ^ ^ j> P 
 
 <u 
 
 Oh 
 
 1 
 
PERIODIC SYSTEM OF THE ELEMENTS. 
 
 247 
 
 cated for them by the system. All such alterations have been confirmed by recent inves- 
 tigations. The series of gold, iridium, platinum and osmium did not acsord with the 
 system, which rather demanded the following : Osmium, iridium, platinum and gold. 
 This requirement was satisfied by a redetermination of the atomic weights of these ele- 
 ments by Seubert [Ann. Chem. (1891) 261, 272]. Hence, the periodic system offers a 
 control for the numbers of the atomic weight, while formerly they appeared to be irregu- 
 lar, and, at the same time, accidental. We are consequently justified, until we have 
 more evidence to the contrary, in assuming that the determinations of the atomic weight of 
 tellurium have placed that value too high. If it should finally be proved to be greater 
 than that of iodine, then the periodic system would be seriously affected in its founda- 
 tion ; it would then lose its claim to being a natural law — for this would not tolerate an 
 exception. 
 
 Mendelejeff, on the basis of the periodic system, predicted in a manner similar to 
 that shown by Leverrier in his precalculation of neptune, the existence of new, not yet 
 known, elements which correspond to unoccupied, free places or gaps in the table. 
 In fact, three such gaps have been filled by the discovery of gallium, scandium, and 
 germanium ; their properties have shown themselves to be perfectly accordant with 
 those deduced from the periodic system. [It may be stated here that Dobereiner, on 
 
 the basis of the law of Triads, — Cl, Br, I ; Li, Na, K ; Ca, Sr, Ba, etc.: = Br, 
 
 etc., — predicted two elements, which should form a triad with fluorine, and one a 
 complete series with phosphorus and arsenic, but the prediction has not been realized.] 
 At present, only the first homologue of manganese (with atomic weight of about 100) is 
 wanting. In the very incomplete series of period V, it is probable that the poorly in- 
 vestigated elements mentioned on p. 27 will find place (see p. 244). 
 
 The entire character of a given element is determined to a very high 
 degree by the law of periodicity; hence, all physical and chemical 
 properties of the same are influenced by its position in the system. 
 These relations we will examine more closely in the individual groups of 
 the metals, and here confine ourselves to a notice of some general rela- 
 tions, and the connection of atomic weight with the chemical valence of 
 the elements. 
 
 The relation of metalloids to metals is shown with great clearness in 
 the periodic system. The first members of all periods (on the left side) 
 consist of electro-positive metals, forming the strongest bases, the alka- 
 lies — Cs, Rb, K, Na, Li, and metals of the alkaline earths — Ba, Sr, Ca, 
 Mg, and Be. The basic character diminishes successively, in the follow- 
 ing heterologous members, and gradually passes over into the electro- 
 negative, acid-forming character of the metalloids FI, Cl, Br, I. Here 
 is observed that, in the periods following each other, with higher atomic 
 weights, the basic metallic character constantly exceeds the metalloidal. 
 The first period comprises five metalloids (B, C, N, O, FI), the second 
 only four (Si, P, S, Cl), the fourth and fifth periods each only three (or 
 two) metalloids (As, Se, Br, and Sb, Te, I), which, at the same time, 
 become less negative. With the metalloidal nature is combined the 
 power of forming volatile hydrogen compounds. Similar volatile deriva- 
 tives are also afforded by the metalloids with the univalent hydrocarbon 
 groups (as CH3, CjH^, CgH^, etc.), which resemble hydrogen in many 
 respects. Such metallo-organic compounds, in which the elements show 
 the same valence as in the hydrogen compounds, are also produced by 
 the metals adjacent to the metalloids : 
 
 II III IV III II I 
 
 MgfCIb), Al(CIl3)3 Si{CH3h P(CH3)3 S(CIl3), CICIL. 
 
248 
 
 INORGANIC CHEMISTRY. 
 
 Their Stability gradually dimill islics with the increasing basic nature of 
 the metals; hence, in the three large jieriods, this power extends only to 
 Zn, Cd, and Hg. 
 
 In consequence of the opposite character of the two ends of the jieriods, 
 there are in the table representing the double ])eriodicily of the great 
 periods (pp. 244 and 246) two sub-groups each, with the seven vertical 
 groups ; on the left with the more positive, basic, and on the right with 
 the more negative, metalloidal elements. Thus in group VI, in addition 
 to O and S (belonging to the small periods) stands the more basic sub- 
 group Cr, Mo, W, and the metalloids Se and Te ; in group II stand the 
 strong basic metals Ca, Sr, Ba, and the less basic heavy metals Zn, Cd, 
 Hg. The elements of group VIll form the gradual transition from the 
 latter to the former. 
 
 Periodicity of Chemical Valence. — Groui) I of the table com- 
 prises the univalent metals, group II the bivalent. In group III are the 
 trivalent metalloid, boron, and the tri valent metals Al, Sc, Y, and Ga, 
 In, Tl. In the quadrivalent carbon group the valence arrives at its maxi- 
 mum ; trom here it gradually decreases with increasing atomic weight ; 
 the nitrogen group is trivalent, the oxygen group is bivalent, that of the 
 halogens univalent. This valence is derived from the compounds with 
 hydrogen and hydrocarbons (compare p. 247), or where such do not 
 exist, as in the case of boron and many metals, from the halogen com- 
 pounds : 
 
 
 
 
 IV 
 
 III 
 
 II 
 
 I 
 
 
 
 
 CH, 
 
 NH3 
 
 UH2 
 
 FIH 
 
 I 
 
 II 
 
 III 
 
 IV 
 
 III 
 
 II 
 
 I 
 
 LiCl 
 
 BeCq 
 
 Bcq 
 
 ccq 
 
 NCI3 
 
 0CI3 
 
 Fb 
 
 NaCl 
 
 MgCb 
 
 AICI3 
 
 Sicq 
 
 PCI3 
 
 seq 
 
 cq. 
 
 The elements of the first four groups are not capable of yielding higher 
 compounds with the halogens.* On the other hand, as we have seen, 
 the higher analogues of nitrogen and other metalloids can unite with a 
 larger number of halogen atoms (see pp. 169, 1 70). The higher valence of 
 these elements is more manifest in the more stable oxygen compounds. 
 On bringing together the highest oxides of the seven groups capable of 
 forming salts (salt-building oxides), we get this series: 
 
 I II III IV V VI VII 
 
 LijO BeO COj SO3 (ip/). 
 
 The elements of the first four groups in their oxygen compounds 
 exhibit, consequently, the same valence as in the compounds with hydro- 
 gen (or hydrocarbon radicals) and the halogens; in the last three series, 
 however, there is noticed a constant increase of valence for oxygen. f 
 
 * Recently lialide derivatives of c;esium and rubidium of the formulas CsMj and 
 CsMr, liave been prej)ared (see Oesium). 
 
 f 'I'lie comix)und sulphur heptoxide (jn 188), evidently does not fit in here. It 
 
 v(uy probably belongs to the (rue peroxides with the chain -()-()-, so that its compo- 
 sition does not give a definite answer as to the valence of sulphur. 
 
PERIODIC SYSTEM OF THE ELEMENTS. 
 
 249 
 
 Besides these highest oxides, remarkable for their greater stability, the 
 elements of the last three groups afford lower oxides, returning in this 
 manner to the hydrogen valence : 
 
 III 
 
 IV 
 
 V 
 
 P303 
 
 S 02 
 
 bOs 
 
 
 II 
 
 III 
 
 
 SCb 
 
 (CIA) 
 
 I 
 
 CI2O 
 
 PH 3 
 
 SHj 
 
 cm. 
 
 The hydroxyl compounds of the elements of the seven groups are anal- 
 ogous to the oxides in constitution. They afford the following series, 
 expressing the maximum valence (compare p. 169) : 
 
 I II III IV V VI VII 
 
 Na(OH) Mg(OH )2 Al(OH )3 Si(OHh P(OH )5 S(OH)e 0(011)^. 
 
 The hydroxyl compounds of the elements of the first four groups exist in 
 free condition, excepting that of carbon, C(OH)^, which is only repre- 
 sented in its derivatives. . The strong basic character of the hydroxides 
 of group I (NaOH) diminishes, step by step, in the succeeding groups, 
 down to the weak acid hydrate, Si(OH)^. The hydrates of the last three 
 groups are of acid nature, and mostly unstable or not known. By the 
 elimination of one, two and three molecules of water they yield the 
 ordinary acids: 
 
 V VI VII 
 
 P0(0H)3 S02(0H)2 C103(0H). 
 
 Phosphoric acid. Sulphuric acid. Perchloric acid. 
 
 The non-saturated hydroxides behave in the same way : 
 
 III IV V 
 
 ?(0H)3 S(0H), C1(0H)3 
 
 III 
 
 C1(0H)3 
 
 I 
 
 . . Cl(OH). 
 
 Sulphurous acid, S0(0H)2, is derived from the hydrate, S(OH)j 
 chloric acid, ClOg.OH, from the hydrate, Cl(OH)5; and chlorous acid, 
 CIO. OH, from the hydrate, Cl(OH)3. The hydrates, P(0H)3 and 
 ClOH, are very unstable, and the first appears to pass readily into 
 HP0(0H)2 (compare p. 215). 
 
 It has been already shown in the case of periodic, sulphuric, and nitric 
 acids, how the so-called hydrates with water of crystallization (regarded 
 as molecular compounds) are explained by the acceptance of the exist- 
 ence of such hydroxyl derivatives. The same may be done for many salts 
 with water of crystallization. 
 
 Thus, we see, and in the following pages will find it more extensively 
 developed, that the relations of valence of the elements have their com- 
 plete expression in the periodic system, are regulated by it, and hence 
 we must conclude that, in fact, the valence is not only a projierty at- 
 taching to the elements per se, but is influenced also by the nature of the 
 
250 
 
 INORGANIC CHEMISTRY. 
 
 combining elements; the hydrogen valence is constant, the valence of 
 oxygen and the halogens, on the contrary, varies according to definite 
 rules. Valence, therefore, is a relative function of the elements (j). 169). 
 
 The periodic system, the natural division of the elements, apjiarently 
 emphasizes the fact that the atoms of the chemical elements are aggrega- 
 tions or condensations of one and the same iirimordial substance — the 
 unity of matter, which corresponds to the unity of force already known 
 tons. It is only by assuming a jirimordial substance that the periodic 
 dependence of the properties of the chemical elements iiiion the magni- 
 tude of the atomic weights can be comjirehended. 
 
 It was suggested that hydrogen was this original substance, because it 
 appeared as if all atomic weights had to be expressed in whole numbers 
 — that is, multiples of the atomic weight of hydrogen, taken as unit 
 (Front 1815, Meinecke 1818). More accurate determinations, made 
 with exceeding care by Stas, demonstrated this not to be true in all 
 cases [J. S. Stas, Untersuchungen iiber die Gesetze der chemischen Pro- 
 portionen, iiber die Atomgewichte und ihre gegenseitigen Verhaltnisse, 
 Deutsch von Aronstein, Leij)zig, 1867]. The inquiry was then made as 
 to whether the variations of the atomic weights from whole numbers 
 could not be occasioned by definite amounts of ])onderable ether enter- 
 ing or escaping in the chemical transposition of bodies. The experi- 
 ments of H. Landolt on this point gave a negative reply, thus cutting 
 off the last avenue which remained open to the hypothesis of Prout and 
 Meinecke [Ber. 26 (1893), ii, 1820]. 
 
 It must be borne in mind, however, that it is not only the atomic 
 weight which influences the properties of an element, but that the molec- 
 ular weight and the energy of the molecule exercise great influence in 
 this direction. The occurrence of allotropic modifications of an element 
 can only be explained on this assumption. These, in their general 
 deportment, often differ more from one another than two different but 
 related elements. Phosphorus in its varieties is a most striking example. 
 Similar relations have been noticed with carbon — soot, graphite, dia- 
 mond, and with oxygen — ordinary oxygen and ozone. It is only the 
 power of uniting a definite number of different atoms which remains in 
 all varieties of one element. Thus, red phosphorus on oxidation yields 
 the same oxides and acids as the yellow, and by burning the diamond the 
 same carbonic acid is produced as with soot and graphite. The valence 
 which one element shows with reference to other elements, is, there- 
 fore, primarily a function of the atomic weight. Its physical properties, 
 on the other hand, are also dependent upon the molecular weight and the 
 energy of the molecule. 
 
 The la 7 v ttnderlyifig the periodic system will only be recognized perfectly 
 and in its eiitirety ivhcn it becomes possible to deduce the properties of the 
 different varieties of a )i element fro 7 n the atomic 7 veight of the latter. 
 
 Again, these allotrojiic modifications indicate that bodies can be formed 
 by a different arrangement of similar atoms, and it is only by keen scien- 
 tific observation that varieties of one and the same substance can be 
 detected, d’he elements could similarly consist of varieties or allotropic 
 forms of some one ])rim()rdial form of matter. 
 
THE METALS. 
 
 Although there is no sharp line of demarcation between metals and 
 non-metals, yet these two classes of bodies are characteristically distinctive 
 in their entire deportment, as may be plainly seen in the periodic system 
 of elements. In physical respects the character of metals is determined 
 by their external appearance and by their ability to conduct heat and 
 electricity; chemically, it shows itself chiefly in the basicity of the oxygen 
 compounds ; yet we see that with the increase of the number of the oxygen 
 atoms, the basic character gradually diminishes and becomes acidic. 
 Gaseous compounds of the metals with hydrogen are not known. 
 
 PHYSICAL PROPERTIES OF THE METALS. 
 
 At ordinary temperatures all the metals excepting mercury are solid, 
 slightly volatile bodies. They are opaque, and only a few, like gold, 
 permit the passage of light to a limited extent when beaten into thin 
 leaflets. In compact mass they exhibit metallic luster and mostly possess 
 a whitish-gray color ; gold and copper are, however, brilliantly colored. 
 In powder form almost all the metals are black. 
 
 Most of them crystallize in the forms of the regular system ; only a 
 few, showing a metalloidal character, are not regular. Thus antimony 
 and bismuth crystallize in the hexagonal system, and tin is quadratic. 
 
 The specific gravities of the metals vary greatly, from 0.59 to 22.5, as 
 seen from the following arrangement : 
 
 Ij'thium, 
 
 0-59 
 
 Germanium, 
 
 5*47 
 
 Ixad, 
 
 11-37 
 
 Pota.ssium, 
 
 0.87 
 
 Zinc, 
 
 7-1 
 
 Palladium, 
 
 II-5 
 
 Sodium, 
 
 Rubidium, 
 
 0.97 
 
 Tin, 
 
 Iron, 
 
 7-3 
 
 Thallium, 
 
 II. 8 
 
 152 
 
 7.8 
 
 Mercury, 
 
 13-55 
 
 Calcium, 
 
 1-57 
 
 Cobalt, 
 
 8.5 
 
 Gold, 
 
 19-3 
 
 Magne.sium, 
 
 1.74 
 
 Copper, 
 
 8.9 
 
 Platinum, 
 
 21.5 
 
 Aluminium, 
 
 2.6 
 
 Bismuth, 
 
 9.8 
 
 Iridium, 
 
 Osmium, 
 
 22.4 
 
 Barium, 
 
 3-75 
 
 Silver, 
 
 10.5 
 
 22.5 
 
 In general the specific gravities of the metals, and also those of the 
 metalloids, increase with the atomic weights; they stand more especially 
 in sharp periodic dependence with reference to the latter. The first 
 members of all periods possess low specific gravities; the latter grow 
 
 * Liquid ; solid = 14. 2. 
 251 
 
252 
 
 INORGANIC CHEMISTRY. 
 
 gradually until the middle of the period, when the maximum is attained, 
 and then they again decrease (p. 245). These relations show themselves 
 more fully if, instead of the specific gravity, we comi)are the siiecific 
 volumes or atomic volumes; i. e., the cpiotients from the atomic weights 
 (A) and specific gravities (d) : 
 
 ^ = specific volume. 
 
 These quotients express the relative value of the atoms (in solid or 
 liquid state), 'hhus the atomic volume of lithium = 11.9, that 
 of potassium = 45-o^ i- potassium atom occiqn'es a sjiace 
 
 3.8 times larger than that of the lithium atom. The periodic alterations 
 of the atomic values are in opposition to those of the specific gravities, as 
 the former are obtained by the division of the atomic weights by the 
 specific gravities. Therefore, the atomic volumes decrease gradually, 
 commencing with the first members of the periods, attain a minimum in 
 the middle of the periods, and then increase again up to the last mem- 
 bers. On the other hand, we find that with the homologous elements 
 (the vertical series) an increase in the atomic volumes almost invariably 
 occurs as the atomic weights increase. 
 
 The metals whose specific gravities are less than 5 are termed light 
 7 netals, the rest heavy 7 netals. The naturally occurring compounds of the 
 heavy metals are termed ores. 
 
 Most metals are very 77 talleable and tenacious, hence can be beaten into 
 thin plates and leaves, and drawn out into wires ; gold and silver are the 
 most malleable. A few, like bismuth and tin, possess a metalloidal char- 
 acter, are brittle, and may be pulverized like antimony and arsenic. 
 
 Heat will fuse all metals, although some require high temperatures. 
 The fusing points of the most important of them are the following : 
 
 Mercury, . . . , 
 
 . . . . —39° 
 
 Aluminium. . . . 
 
 • . . 650° 
 
 Rubidium, . . . , 
 
 . . . .+38° 
 
 Germanium, . . . 
 
 . . . 900° 
 
 Potassium, . . . , 
 
 . . . . 62° 
 
 Silver, 
 
 . . . 960° 
 
 Sodium, , . , , 
 
 . . . . 96° 
 
 Gold, 
 
 . . . 1060° 
 
 Tin, 
 
 . . . . 230° 
 
 Copper, 
 
 . . . 1090° 
 
 Bismuth, . . . , 
 
 . . . . 265° 
 
 Cast iron, .... 
 
 . . . 1150° 
 
 Cadmium, . . . , 
 
 . . . . 320° 
 
 Wrought iron, . . 
 
 . . . 1600° 
 
 Lead, 
 
 . . . . 330° 
 
 Platinum, .... 
 
 . . . 1775° 
 
 Zinc, 
 
 ... 420° 
 
 Iridium, .... 
 
 . . . 1950° 
 
 Ruthenium, osmium and also chromium, molybdenum, uranium, tung- 
 sten and vanadium melt with greater difficulty than platinum. 
 
 A greater volatility also corresponds to the greater fusibility. Mer- 
 cury boils at 357°; potassium and sodium at about 700°; cadmium at 
 770°; zinc about 950°, and the difficultly fusible metals may be also 
 volatilized in the electric furnace. 
 
 Metallic borides, silicides and carbides (p. 163) withstand temperatures 
 which volatilize i)latinum, calcium and carbon (Moissan). 
 
SPECIFIC HEAT — ATOMIC HEAT. 253 
 
 All these physical properties bear a periodic dependence to the atomic 
 weights, as will be more plainly indicated in the indvidual groups. 
 
 Until lately the highest temperature attainable for scientific and technical uses was 
 about 2000°. The idea of using the electric arc to extend the range of temperature is 
 not new, but the actual achievement was first accomplished by the French chemist, H. 
 Moissan, by the construction of the electric furnace. The principle involved in the latter 
 is to expose a substance to the heat of the electric arc in a space as small as possible and 
 surrounded by fire-resisting material. The action is then due to heat and not to elec- 
 tricity. Moissan in this way obtained a temperature of 3500° C. The furnace model 
 first used by him in 1892 is 
 sketched in Fig. 68. It con- 
 sists of two blocks of unslaked 
 lime. The lower contains a 
 groove for the carbon elec- 
 trodes E, E, and about the 
 middle is a depression in 
 which the material is placed 
 and upon which the heat of 
 the arc is to act. The material 
 can also be exposed in a small 
 carbon crucible. The upper 
 portion is slightly arched above the arc. As the lime soon melts on the surface and 
 becomes soft the arch acts as a heat reflector. For the various forms of electric furnaces 
 consult H. Moissan, Der elekt. Ofen, 1897 ; also Liebetanz. Calciumkarbid- und Acety- 
 lentechnik, pp. 38-79. Compare Z. f. anorg. Ch. 22 (1896), 293. 
 
 SPECIFIC HEAT. ATOMIC HEAT. 
 
 Of all physical properties of the elements, from a chemical standpoint, 
 their heat capacity is the most important, as it can serve for the determi- 
 nation of the atomic weights. The specific heat of a substance is the 
 quantity of heat which increases its mass unit i° in temperature. Heat 
 quantities are measured in either large or small calories (p. 66), and 
 according as one or the other is chosen, the expression for the mass unit 
 will be I kilogram or i gram. As water, of all substances, possesses the 
 greatest specific heat, and as this also serves as unit; therefore, for all 
 other bodies the specific heat represents the fraction of a calorie which 
 raises their mass unit i°. 
 
 The atomic heat of an element is the product of its specific heat (H) 
 and atomic weight (A). Dulong and Petit (i8i8) discovered the remark- 
 able fact that the solid elements possess approximately the same atomic heat. 
 It is about 6.4 : 
 
 A. H = 6.4. 
 
 For silver A = 107.93 H = 0.057 ; A . H = 6. i ; for lead A = 
 206.9, H = o.o3i,A.H = 6 4; for potassium A = 39.15, H = 0.166, 
 A. H= 6.5 — /. e., 107.93 gi'ams of silver, 206.9 grams of lead, and 39.15 
 grams of potassium are heated by approximately 6.4 small calories 
 through 1° of temperature, or 107.93, 206.9 and 39.15 kilograms are 
 raised 1° by 6.5 large calories. The equation A . H = 6. 4 also indicates 
 that the specific heats of two elements are to each other approximately as 
 their atomic weights — the greater the atomic weight, the smaller the 
 specific heat. 
 
 It is only in the case of a few of the elements that the atomic heat 
 
254 
 
 INORGANIC CHEMISTRY. 
 
 varies widely from the mean 6.4, e. g., sulphur is 5.7, phosphorus 5.9, 
 silicon 4.6, germanium 5.6, while carbon is 2-2.8, boron 2.6 and beryl- 
 lium 3.7. These elements have low atomic weight and are either metal- 
 loids or resemble metalloids. 
 
 These variations from the mean are in part explained by the change in 
 s])ecific heat with the temperature. It was known before that these show a 
 slight increase with the temperature, but it is only recently (1875) that 
 Weber has proved that the increase is very considerable for the elements 
 C, B and Si, which, at medium temperatures, possess a remarkably low 
 atomic heat ; that, beyond a definite temperature, the atomic heat becomes 
 tolerably constant, and then almost agrees with the law of Dulong and 
 Petit. According to Nilson, beryllium shows a similar deportment : 
 
 
 H 
 
 A 
 
 H X A 
 
 Diamond, graphite, above 900°, 
 
 0.459 
 
 12.00 
 
 5-5 
 
 Boron, above 600°, 
 
 0-5 
 
 I I 
 
 5-5 
 
 Silicon, above 200°, 
 
 0.204 
 
 28.4 
 
 5 8 
 
 Beryllium, above 257°, 
 
 0.58 
 
 9.1 
 
 5 3 
 
 From this close agreement of the found atomic heat of the metals with 
 the mean, it follows, without doubt, that there does occur a regularity, 
 and we must conclude that the slight variations, apart from the inaccuracy 
 of the observations and the impurities of the substances used, are influ- 
 enced by physical causes of an unknown nature. Hence, the specific heat 
 may serve for the derivation of the atomic weight of the elements; ^/le 
 atomic weight is equal to the eo?istafit 6.4 divided by the found specific heat : 
 
 The atomic weights derived from the specific heat — the so-called ther- 
 mal atomic weights — agree in almost all instances with those obtained 
 from the vapor density of the free elements or their volatile compounds. 
 Where no volatile compounds of an element are known, the specific heat 
 is the only certain means of fixing the actual atomic weight. The equiv- 
 
 I 
 
 alent weight, 38 (InCl), of indium was fixed with great accuracy by 
 analysis ; it was, however, unknown whether the atomic weight was 
 double or triple that quantity. The specific heat of indium was 
 found to be 0.0569, from which the atomic weight would be 
 = 112, a number closely ap})roaching the trebled equivalent weight of 
 indium, 114 (= 38 X 3)- From this it follows that the true atomic 
 weight of indium is 114 and that indium is trivalent (InClg). 
 
 In their solid com])ounds the elements retain the specific heat ; hence the molecular 
 heat is nearly e(|ual to the sum of the atomic heats of the elements constituting the 
 molecules — law of h'. Neumann and II. Ko])]). I lence the atomic heat of elements not 
 known in solid condition may be derived from the molecular heat of their compounds. 
 In this manner the following atomic heats are found : For nitrogen, 5-0 j for chlorine, 
 C ; for oxygen, 4 ; for (luorine, 5 ; for hydrogen, 2.3 ; for carbon, 1. 8. 
 
CHEMICAL PROPERTIES OF THE METALS. 255 
 
 In the free state the gaseous elements usually have a slighter atomic heat, as seen from 
 the following table : 
 
 
 A 
 
 H* 
 
 A X H 
 
 Oxygen, 
 
 16 
 
 0. 156 
 
 2.5 
 
 Hydrogen, 
 
 1 .01 
 
 2.405 
 
 2.4 
 
 Nitrogen, 
 
 14.04 
 
 0.172 
 
 2.4 
 
 Chlorine, 
 
 3545 
 
 0.093 
 
 3 3 
 
 ISOMORPHISM. 
 
 As irrdicated in the preceding pages, the atomic weights of the elements may be 
 derived directly from the specific heat of solids, while from the gas density of the volatile 
 compounds we get the molecular weights, and from the latter, indirectly, ascertain the 
 atomic weights (compare p. 80). A third, although less general and certain, means of 
 determining the atomic weight is afforded by isomorphism. By this is understood the phe- 
 nomenon observed by Mitscherlich (1819). Characteristics of the isomorphism of two 
 bodies are : 
 
 (l) Similarity in crystalline form, whereby the symmetry properties agree perfectly and 
 the geometrical constants nearly coincide (see pp. 31, 38) ; (2) the power of forming 
 within certain definite limits mixed crystals in any desired quantity by weight without 
 altering the crystal form; (3) the power of overgrowth, i. e., crystals of the one sub- 
 stance being able to continue their growth in supeisaturated solutions of the other com- 
 pound. From Mitscherlich’ s discovery we can conclude from the isomorphism of two 
 compounds that they are similarly constituted and have an equal number of atoms in the 
 molecule. Accordingly the quantities of the elements, occurring in isomorphous mixtures, 
 would be to each other as their atomic weights, and this would make it possible to estab- 
 lish an incorrect atomic value. For example, the metals calcium, strontium and barium 
 do not afford volatile derivatives. Their atomic weights could not be deduced from their 
 thermal capacity, and it was the isomorphism of many of their compounds with those of 
 magnesium that determined the same ; the quantities of these elements, replacing 24.36 
 parts by weight of magnesium (l atom), were accepted as the true atomic weights. 
 
 In the present state of chemistry we attach but secondary importance to isomorphism as 
 a method of determining atomic weights because it is frequently very difficult to demon- 
 strate that isomorphism is actually present, for the coincidence in geometrical constants 
 requires to be only approximate, the ability to yield mixed crystals occurs in every imag- 
 inable degree and bodies possessing not the slightest chemical similarity do show over- 
 growths. It is indeed true that bodies similar chemically often exhibit like crystalline 
 form, and the power of forming mixed crystals also increases the greater the similarity in 
 chemical properties. 
 
 CHEMICAL PROPERTIES OF THE METALS. 
 
 Alloys. — Many metals may be so mixed by fusing them together that a 
 solid, homogeneous mass will result upon cooling ; the constituents cannot 
 be separated then by mechanical means. Such a mixture is based upon a 
 molecular interpenetration ; it cannot be distinguished from a mere 
 mechanical mixture and it bears the name alloy \legere, put together). 
 
 * By constant volume. 
 
256 
 
 INORGANIC CHEMISTRY. 
 
 A great deal of heat is frequently developed just as in a chemical union 
 by fusing metals or by bringing together those already melted. Yet alloys 
 in general cannot be considered chemical compounds; their composition, 
 even in the case of those which crystallize, is rarely according to the 
 atomic ratio. They should rather be regarded as solid solutions of one 
 metal in another or of metallic comi)ounds in a metal, if, following 
 Ostwald, a homogeneous mixture of various substances (showing no vari- 
 ability) is defined, without reference to its state of aggregation, as a 
 solution. 
 
 Many molten metals may be mixed just like water and alcohol in every 
 l)roportion with one another, e. g., tin and lead ; others in a less degree, 
 similarly to ether and water, e. g., zinc and lead, d'he second class of 
 alloys occurs more frequently. Usually metals which are similar chemic- 
 ally alloy more readily with one another than those which exhibit marked 
 differences in their chemical behavior. 
 
 Many alloys can be obtained crystallized, yet in actual practice crys- 
 tallization is avoided because it diminishes the solidity, ductility, etc. 
 
 As a rule, the density of an alloy cannot be calculated from its constit- 
 uents, but this is possible with a mere mechanical mixture. Many alloys 
 exhibit a very definite condensation (contraction), e. g., silver-gold, cop- 
 per-tin ; others a distinct expansion: copper-silver, antimony-tin, tin- 
 lead, while others do not manifest either of these properties: copper-gold, 
 antimony-bismuth. 
 
 The color of alloys does not correspond to the quantity of their con- 
 stituents, for copper alloys, with 30 per cent, of tin, are white in color, and 
 pure yellow with the same amount of zinc. The bright color disappears 
 with 60 per cent, of zinc. Aluminium changes the color of copper 
 readily; the same is true of nickel, as may be seen in the coin, contain- 
 ing about 25 per cent, of nickel and 75 per cent, of co])per. Gold possesses 
 slight coloring power. Gold and silver alloys having 30 per cent, of 
 gold possess a pure silver-white color. 
 
 The melting point of an alloy is generally lower than would be expected 
 from the melting points of its constituents. An alloy of 8 parts of lead, 15 
 parts of bismuth, 4 parts of tin, and 3 parts of cadmium melts at 68”, 
 although each of these metals melts above 200”. Potassium (m. p. 
 62°) and sodium (m. p. 96”) unite in the ratio of their atomic weights 
 to an alloy which is liquid, like mercury, at the ordinary temperature. 
 
 Quite frequently molten alloys manifest the tendency on slow cooling 
 to separate into several compounds which differ in composition, fusibility, 
 etc., from one another. Under these conditions pure metals also sepa- 
 rate from many alloys. This process, very important for the preparation 
 of metals, is designated segregation (see Silver). 
 
 Alloys often resist acids much better than single metals; however, they 
 frequently dissolve in acids even when the one constituent is insoluble in 
 them. Thus an alloy of 56. 5 i)cr cent, of copper and 43.5 ])er cent, of zinc 
 is scarcely attacked by nitric acid, and in gold-silver alloys, where there 
 is a prejiondcrance of gold, the silver is not dissolved out by nitric acid. 
 If the alloy contains 3 ])arts of silver to one part of gold the nitric acid 
 will dissolve the silver (cpiartation). 
 
CHEMICAL PROPERTIES OF THE METALS. 
 
 257 
 
 Mercury is able to dissolve almost all metals, forming alloys known as 
 aftialgains (from imlayiia, a soft plaster), which are generally crystallizable. 
 In chemical respects hydrogen is a metal, and very many of the metals 
 combine with it. Palladium, potassium and sodium yield the com- 
 pounds K^H2 and Na^H.^, which deport themselves as alloys. That 
 antimony yields a gaseous product (SbHg), is due to its pronounced 
 metalloidal character. The ability of individual metals of the platinum 
 and iron groups to permit the passage of hydrogen at a red heat depends, 
 probably, upon a chemical attraction; hydrogen first dissolves and is 
 then evaporated again. 
 
 Metallic Carbides. — These are produced by dissolving carbon in 
 the molten metal or by reducing metallic oxides with carbon. The elec- 
 tric furnace with its high temperatures has been applied in this direction. 
 The researches of Moissan and the technical application of calcium car- 
 bide and silicon carbide (carborundum) have given great theoretical and 
 practical importance to the carbides in general (pp. 154, 163). Molten 
 gold, bismuth, lead and tin do not dissolve carbon, and copper and silver 
 very little. The melted platinum metals take it up in large amounts, and 
 upon cooling it separates as graphite. Many other metals, under like 
 conditions, form carbides which crystallize beautifully. Carbides may 
 be arranged in two series: 
 
 1. Those decomposed by water at the ordinary temperature. Here 
 belong the carbides of potassium, lithium, Li2C2, calcium, barium, stron- 
 tium (Me"C2), aluminium, Al^Cg, manganese, MnC, and beryllium. 
 
 2. Those which are not affected by water : molybdenum carbide, CMOg, 
 chromium carbide, CrgCg, etc. The carbides decomposable by water 
 yield hydrogen, methane, acetylene, and also liquid and solid hydrocar- 
 bons. Moissan thinks that this behavior affords hints as to the origin 
 of petroleum (see Calcium- and Aluminium Carbide, also p. 252). 
 
 Halogen Compounds. — The metals unite directly with the halogens 
 to form compounds, which are not decomposed by water at ordinary tem- 
 peratures, and, in general, are very stable ; on the other hand, the halogen 
 compounds of the metalloids (excepting those of carbon) are easily 
 decomposed by water. These compounds are also produced by the action 
 of the haloid acids upon the free metals, their oxides, hydroxides, and 
 carbonates, whereby they plainly characterize themselves as salts of the 
 haloid acids. A third procedure for the formation of chlorides and bro- 
 mides, essentially analogous to the first, is based upon the simultaneous 
 action of carbon and chlorine, or bromine upon the oxides (see Aluminium 
 Chloride and Silicon Chloride). 
 
 The following types of halogen derivatives exist and show the different 
 valences of the metals : 
 
 I II III IV V VI 
 
 KCl ZnClj InClg SnCl, TaClg WClg. 
 
 22 
 
258 
 
 INORGANIC CHEMISTRY. 
 
 OXIDES AND HYDROXIDES— HYDRATES. 
 
 I'he affinity of the metals for oxygen varies. Some of them oxidize in 
 moist air and decompose water, even at ordinary temperatures. Such are 
 the so-called alkalies and alkaline earths (the potassium and calcium 
 groups). Their oxides dissolve readily in water and form strong basic 
 hydroxides or hydrates, KOH, Ca(OH)2, which are usually not decom- 
 ])Osed by ignition. 
 
 Other metals (the so-called heavy metals) oxidize and decompose water 
 only at higher temperatures; their oxides are insoluble in water, and 
 generally afford no hydroxides, as the latter ui)on heating readily decom- 
 ])Ose into oxides (anhydrides) and water: 
 
 Zn(OII)2 = ZnO + H^O. 
 
 They are of a less basic nature, and their soluble salts usually exhibit 
 acid reaction. Finally, some metals, as gold and platinum (the noble 
 metals), are incapable of combining directly with oxygen. Their oxides, 
 obtained in another way, decompose readily under the influence of heat 
 into metal and oxygen. The universal method for the preparation of 
 insoluble oxides and hydroxides of the heavy metals depends upon the 
 precipitation of the solutions of their salts by alkaline bases: 
 
 HgCl., + 2KOII = 2KCI -h HgO + HjO 
 
 II II 
 
 CuSO* + 2KOH = K2SO4 + Cu(OH)j. 
 
 The different valences of the metals are most clearly seen in their oxygen derivatives, 
 which form salts. We have the following eight forms or types of the highest salt-pro- 
 ducing oxides (see p. 244), corresponding to the eight groups of the periodic system of 
 the elements : 
 
 I II III IV V VI VII VIII 
 
 K2O MgO AI2O3 Sn 02 Bi 205 CrOg Mn 207 OSO4. 
 
 The hydroxides of the first two forms possess a strong basic character and only yield 
 salts with acids. In the hydroxides of the two succeeding forms there is shown an acid- 
 like character together with the predominating basic character. Hence they dissolve in 
 alkalies and form salt-like derivatives with bases, in which hydrogen is replaced by metals, 
 e. g., Al(ONa)3. These higher (normal) hydrates are not very stable, give up water and 
 })ass into meta-hydrates, which retain the acid character. Thus, from Al(OH)3is derived 
 AlO.OH, which yields salt-like compounds, e. g., AlO . ONa ; from Sn(OH)4 are 
 derived stannic acid, SnO(OII).2, and its .salts, as SnO(ONa)2. The oxides of the next 
 fundamental forms, 1^2^^ Ma^O^, are of an acid nature. The corresponding highest 
 hydroxides do not exist ; but the salts of the anhydro-acids, derived from them, are known : 
 
 IIBiOs Ik^CrO^ HMnO^. 
 
 Bismuthic Chromic Permanganic 
 
 acid. acid. acid. 
 
 In comj)osition these acids correspond to the following metalloid acids: 
 
 IINO3 IICIO,. 
 
 Nitric acid. Sulphuric acid. Perchloric acid. 
 
 Salts of osmium tetroxide (OsO^) are not known. 
 
OXIDES AND HYDROXIDES — HYDRATES. 
 
 259 
 
 T.ike the metalloids, the metals of the last four series form lower oxides and hydrates 
 in which they exhibit a lower valence : 
 
 SnO BiaOs MnO 
 
 Sn(OH)2 Bi(OH)3 Mn(OH)2. 
 
 These lower oxides and hydroxides have a basic character, and in their whole deport- 
 ment resemble the corresponding compounds of the metals of the first three groups. 
 
 The metals of the first two groups have higher oxygen compounds, 
 called peroxides, e. g., Na202, BaOj. These do not form corresponding 
 salts. By the action of dilute acids hydrogen peroxide is produced : 
 
 BaO, -f 2HCI = BaClj -f H^Oj. 
 
 In consequence of this reaction, it is very probable that in the peroxides, 
 the oxygen atoms are arranged in a chain-like manner as in hydrogen 
 peroxide : 
 
 Na — O. 
 
 Na — O'^ 
 
 Sodium peroxide. Barium peroxide. 
 
 When concentrated acid acts upon them, oxygen is evolved, and salts 
 of the lower oxides result; heated with hydrochloric acid, chlorine is 
 generated : 
 
 BaOj + 4HCI = BaClg + 2H2O T Clj. 
 
 Ordinarily, all higher oxides which evolve chlorine with hydrochloric acid are termed 
 peroxides, e. g., Pb02, lead peroxide, and MnO.^, manganese peroxide. However, these 
 latter compounds do not possess the structure of true peroxides. Lead dioxide, Bb02, 
 is wholly analogous to tin dioxide, SnO.^, and is capable of combining with bases ; it 
 also yields a tetrachloride, PbCb, which it is true readily breaks down into the dichloride 
 and chlorine ; therefore we infer that the two oxygen atoms are in direct union with 
 quadrivalent lead. So manganese is probably quadrivalent in manganese peroxide. 
 The difference between these oxygen compounds and the true peroxides is shown by 
 their inability to form hydrogen peroxide. Persulphuric and percarbonic acids belong 
 to the true peroxides. 
 
 Finally, some univalent metals are capable of forming oxides contain- 
 ing four atoms of metal, e. g., K^O, Ag^O ; these compounds are termed 
 quadrant oxides or suboxides. 
 
 
 Salts. — By the action of bases upon acids, salts and water result : 
 NaOH + HNO3 = NaN03 + H2O, 
 
 whereas oxides and acid anhydrides, e. g., CaO and CO 2 , in the absolute 
 absence of water do not combine, or at least only with difficulty. 
 
 d'hese salts are also i)roduced like the halides by the action of metals 
 upon the acids : 
 
 H2SO, j Zn = ZnSO, f IL, ; 
 
INORGANIC CHEMISTRY. 
 
 260 
 
 most of the oxygen acids suffer simultaneous reduction liy tlie hydrogen 
 evolved, d'his hapiiens with suli)huric acid if it be used in a concentrated 
 form, so that, e. g., in the action of zinc on concentrated .sul])huric acid 
 both sulphurous acid and hydrogen sulphide are produced ; 
 
 II2 + 2II/); 
 
 4ll2 + II,S + 41120 . 
 
 Sulphurous acid, nitric acid and chloric acid are also reduced in dilute 
 solutions (see pp. 184, 186, 203). Phosphoric acid, in contrast to arsenic 
 and antimonic acids, is very stable in the presence of nascent hydrogen 
 ( pp. 148, 216). The salts of oxygen acid, like those of the haloid acids, are 
 to be regarded as derived from the acids by replacement of their hydrogen 
 by metals. The acids themselves can be viewed as hydrogen salts. 'This 
 was the conception of IT Davy (1815), of P. I.. Dulong (1816), and par- 
 ticularly of J. Liebig (1838), from chemical considerations, and also a 
 little later (1839) of J. Fr. Daniell, from electrolytic reasons. 
 
 As we have seen, the poly basic acids yield the primary, secondary, ter- 
 tiary, etc.g salts by the replacement of one or several hydrogen atoms by 
 metal (see p. 172). In the same manner n-series of salts are derived from 
 n acid bases (which contain n-OH groups) : 
 
 .OH 
 
 rOH 
 
 /OH 
 
 (NO, 
 
 Bi^OH 
 
 OH 
 
 BifN 03 
 
 Bi \ NO, 
 
 ^OH 
 
 iN 03 
 
 \no: 
 
 [no,. 
 
 Such salts in which not all the hydroxyl groups of the polyacid hydrox- 
 ide are replaced by acid residues are called basic : 
 
 Basic lead nitrate. Basic zinc chloride. 
 
 Besides these basic salts there exist some of another form. We saw 
 
 that the polybasic acids can combine to form poly- or anhydro-acids ; 
 similarly, the polyhydric bases form polyhydrates : 
 
 SO,<OH 
 
 S02<0H 
 
 Cu<OH 
 
 Cu<OH 
 
 Pb< 
 
 Pb< 
 
 Pb< 
 
 OH 
 
 O 
 
 O 
 
 on, 
 
 from which basic salts are obtained (see Copper and Lead) by replace- 
 ment of hydroxides by acid residues. 
 
 By the replacement of the hydrogen atoms in the polybasic acids or of 
 the hydroxyl groups in the polyhydric bases by various radicals we get the 
 so-called mixed or double salts : 
 
 K 
 
 NIL 
 
 H 
 
 Potassium ammonium 
 phosphate. 
 
 f‘<Cu 
 
 SO.<K 
 
 Potassium copper 
 sulphate. 
 
 S0,< 
 
 so,^ 
 
 K 
 
 A1 
 
 Potassium aluminium 
 sulphate. 
 
 ’‘''<co 
 
 I’l><c| " 
 
 MgClj. KCl 
 
 NO3 
 
 Cl, 2 
 
 AuCb.KCl PtCh.2KCl. 
 
ACTION OF METALS UPON SALTS AND ACIDS. 
 
 261 
 
 Just as the fluorides of boron and silicon, BFI3. KFl, SiFl^. 2KFI, are 
 derived from peculiarly constituted atomic acids, HBFl^, H.^SiFl^^ (pj). 
 170, 226), so also many of the so-called double chlorides,* e. g., PtCl^.- 
 2KCI, must be viewed as atomic in their structure (see p. 271 and under 
 Platinum). 
 
 ACTION OF METALS UPON SALTS AND ACIDS. 
 
 We have seen that the metals by solution in acids are able to form salts, 
 h\ drogen being simultaneously evolved. 
 
 The metals deport themselves in the same manner with the salts. Zinc 
 introduced into a solution of copper sulphate is dissolved to sulphate and 
 metallic copper is deposited : 
 
 Zn -f- CuSO^ = ZnSO^ -f- Cu. 
 
 Herein is shown the perfect analogy between acids and salts : 
 
 Zn + H^SO, = ZnSO, + H^. 
 
 In chemical nature hydrogen is a metal. Hence the acids maybe viewed 
 as hydrogen salts. 
 
 The displacement of a metal from its salts by others appears to be in- 
 fluenced by its electrical deportment in so far as it occurs, in general, in 
 accordance with its position in the electric tension series. Indeed the 
 more electro-positive (basic) metals replace the electro-negative (less 
 basic). In the following series each metal throws out from solution those 
 preceding it : Au, Pt, Ag, Hg, Cu, Pb, Sn (Fe, Zn). Iron and zinc pre- 
 cipitate almost all the heavy metals from solutions of their salts. The 
 strongly positive potassium and sodium are able to displace all other 
 metals. This is very evident from their action upon fused haloid salts — 
 a reaction which frequently serves for the separation of the metals in a 
 free condition : 
 
 AICI 3 + sNa = A1 + sNaCl. 
 
 In its electrical deportment hydrogen stands near zinc ; like the latter, it 
 must, therefore, displace all more negative metals. If this does not happen, 
 the cause must be sought in the volatility of hydrogen ; in fact, we know 
 that hydrogen, under powerful pressure, is capable of separating gold, 
 silver and some other metals from their salt solutions. 
 
 In addition to electrical deportment, the nature of the solution influences the precipi- 
 tation of one metal by another. Thus, lead separates tin from its chloride, wSnCq, while, 
 on the other hand, tin throws out lead from the solution of its oxide in alkalies. This 
 recalls the fact that iodine is displaced in hydriodic acid and the metallic iodides by chlo- 
 rine ; and that conversely iodine replaces chlorine in its oxygen derivatives (p. 278). 
 
 Mylius and Fromm (Her. 27 (1894) i, 630) claim that when positive metals act upon 
 dilute salt solutions of negative metals, alloys are produced, because in the moment of 
 their precipitation metals are capable of uniting at the ordinary temperature. Generally 
 these alloys are y)orous, l)lack, and either amorphous or crystalline. The following crys- 
 talline compounds were obtained : Cu.^Cd, AuOkj, CiqSn and lead-platinum. Zinc pre- 
 
 * Consult American Chemical Journal, li, p. 291. 
 
262 
 
 INORGANIC CHKMIS'I’RV. 
 
 cipitates a zinc- silver alloy from dilute silver solutions, which is converted into white crys- 
 talline silver by a concentrated silver solution, and also hy strong acids : 
 
 Ag„Zn I 2AgN()3 = Ag(„ t 2) + /^(NOa)^. 
 
 The other alloys are similarly decomposed hy acids, the negative element being set free. 
 
 ELECTROLYSIS OF SALTS. 
 
 Conductors are divided into two classes, depending upon their behavior 
 in the passage of the electric current. Conductors of the first class are 
 not decomposed in the passage of the current : the metals, peroxides and 
 carbon. Conductors of the second class carry the current with simul- 
 taneous decomposition : “ Electrolysis ” — salts, acids and bases, fused or 
 in aqueous solution. They are also termed electrolytes (p. 92). In 
 most instances the conductivity of the first class diminishes with rise in 
 temperature, while it increases with the second class. 
 
 On subjecting a salt in a fused or dissolved condition to the action of 
 an electric current, it is decomposed, so that the metal separates at the 
 negative pole and the acid group or halogen at the positive : 
 
 NaCl = Na 4 Cl 
 CuSO, = Oi + SO,. 
 
 The liberated acid residues, e. g., SO,, cannot exist in a free condition ; 
 they break down into oxygen and an acid oxide, which, with the water 
 of the solution, again forms the acid : 
 
 SO, -f H^O = H^SO, + O. 
 
 Thus, in the electrolysis of such salts, the metal and oxygen separate 
 out — the former at the negative, the latter along with free acid at the 
 ])ositive pole (see p. 92 note). 
 
 All neutral salts are similarly decomposed. If, however, the metal 
 contained in the salt acts upon water when free, manifestly a secon- 
 dary reaction must occur at the negative pole. The real electrolytic 
 decomposition of potassium sulphate would then take place according to 
 the following equation : 
 
 K3SO, = K, + (S^,). 
 
 d’he sei)arated potassium decomposes the water with formation of potas- 
 sium hydroxide and the disengagement of hydrogen : 
 
 2K -f 211,0 2K0II + 11.,. 
 
 Therefore, hydrogen and potassium hydroxide occur as definite decom- 
 j)osition ])roducts, at the negative ])ole (the kathode); at the positive 
 (the anode), however, we have oxygen and sulphuric acid. On coloring 
 the liquid exposed to the electrolysis with a little violet-syrup, that part 
 
ELECTROLYSIS OF SALTS. 
 
 263 
 
 at the positive pole will become red, owing to the acid formed, while that 
 at the negative pole will be colored green by the base. That the electro- 
 lytic decomposition of potassium sulphate and similar salts proceeds in 
 the manner given, may be proved experimentally by using mercury as the 
 negative electrode ; then the separated potassium will combine with the 
 mercury and form an amalgam, which will act gradually upon the water. 
 
 The bases and acids are decomposed in this manner by the current. 
 Thus, as Davy (1807) showed, molten caustic potash, KOH, breaks down 
 into K and OH; the first separates in metallic form upon the negative 
 pole (and gradually acts upon the fused KOH with the liberation of 
 hydrogen), while at the positive pole water and oxygen appear — produced 
 from the OH-ions : 
 
 2 ( 0 H) = H^O + O. 
 
 Acids when electrolyzed behave like hydrogen salts. Hydrochloric 
 acid, for example, breaks down into hydrogen and chlori-ne. Sulphuric 
 acid (depending upon its strength) is decomposed in aqueous solution 
 into the ions 2H and SO^ or H and HSO^. The anion SO^ (or HSOJ 
 is immediately converted by the water into sulphuric acid and oxygen : 
 
 2HSO, + 2H2O = H2SO, -h O. 
 
 However, under certain conditions, an appreciable amount of persul- 
 phuric acid is produced : 
 
 2HS0^ = H.^S^Og. 
 
 [See p. 188 and Richarz, Ber. 21 (1888), 1673.] 
 
 The electrolysis of acid sulphates and of acid carbonates can also pro- 
 ceed in the direction last indicated; they break down into kations H 
 and anions MeSO^ and MeCOg, from which arise persulphates and per- 
 carbonates: MegSgOg and MegCgOg ; see also p. 231. 
 
 In the thirties of the present century Michael Faraday demonstrated that 
 when a compound is decomposed by the electric current the quantity of 
 the material so broken down is proportional to the quantity of electricity 
 which has passed through it. He also discovered the following law bear- 
 ing his name : 
 
 When the electric cu7'rent passes through differefit deco7nposable conduc- 
 tors a7'ranged m series^ the qua7itities of the substances separated si77iulta- 
 77eously are to each other as the ratio of their che7nical equivale7its. 
 
 Thus in the simultaneous decomposition of hydrochloric acid, dilute 
 sulphuric acid and ammonia* (pp. 77, 98), equal volumes of hydrogen are 
 liberated, while at the positive pole i volume of chlorine, volume of 
 oxygen, and ^volume of nitrogen appear, after the solutions have become 
 saturated with the gases. The quantities decomposed by electrolysis, 
 therefore, bear the following relation : 
 
 H0SO4 NH3 
 ’ 2 ’ 3 • 
 
 *The hydrochloric acid sliould contain 219 grains in a liter, d'his can be olitained by 
 mixing 200 grams of fuming hydrochloric acid, of specific gravity 1.185, with 128 grams 
 of water. P'or the electrolysis of ammonia use i volume of concentrated ammonia water 
 with 12.5 volumes of a saturated salt solution. 
 
264 
 
 INORGANIC CHEMISTRY. 
 
 In tlie same way, equal quantities of chlorine are set free from all metallic 
 chlorides, while the (juanlities of the j)recipitated metals agree with the 
 values according to which they enter chemical action, their equivalent 
 weight. The quantities of the different salts, decomposed ])y electrol- 
 ysis, stand in the following relation : 
 
 ^ Cu^l-J ShCla FcClo FeCls S11CI2 S11CI4 CuSOj Hk(CN) 2 Hfi:2(N03)2 
 
 ”'^* 2 ' 2 * 3 * 2 * 3 * 2 * 4 * 2 * 2 * 2 ' 
 
 Therefore, 31.8 parts of cojiper are deposited for the 35.45 parts of 
 chlorine in cupric chloride, CuCl.^, but from cuprous chloride, Cu./'l^, 
 we obtain 63.6 parts of copper; from mercuric cyanide, Hg(CN).^, we 
 obtain 100.15 parts of mercury, and from mercurous nitrate, }ig(N03), 
 200.3 parts of mercury, etc. 
 
 Tliese relations can be ol)litcratecl in that reductions of the electrolyte at the negative 
 pole and its oxidation at the positive may occur through the decomposition products. 
 Thus, in the electrolysis of cuprous chloride, cupric chloride instead of chlorine will be 
 found at the anode : 
 
 CuCl 4- Cl = CuCh, 
 
 and in the electrolytic decomposition of cupric chloride cuprous chloride will appear at 
 the kathode : 
 
 CuCb + Cu = 2CuCl. 
 
 Some obscurity of the ratios acquired by the law also arises if the salt is hydrolytically 
 decomposed, e.g., if water converts it partly into hydroxide or basic salt, and free acid : 
 
 FeCh + 3H,0 = Fe(OH)3 + 3110 ; 
 
 the free acid is also subject to electrolytic decomposition. Further, the proportions 
 change iCthe electrodes are attacked by the substances liberated upon them (p, 77). 
 For example, if in the electrolysis of copper sulphate the positive electrode consists of 
 copper then oxygen will not be liberated here (SO^ = SO3 -}- O), but just as much 
 metallic copper will dissolve as separates upon the kathode : 
 
 Cu + SO4 = CuSO, ; 
 
 consequently the copper content of the solution remains unchanged. 
 
 H. Helmholtz accordingly expresses Faraday’s law as follows; 
 
 The same quantity of electricity acting upon different electrolytes liberates 
 equal valences or combines the7n in sojue new way. 
 
 Causs and Weber referred most electrical and magnetic magnitudes to length (cm.), 
 mass (gram) and time (sec.). 
 
 'I'he units derived in this way are the absolute measurements. The unit quantity of 
 electricity in this measurement precipitates, aecording to Kohlrausch, 0.01118 gram of 
 silver. 'I’he tenth part of this quantity, or that which will precij)itate 0.001 1 18 gram — 1. 1 18 
 mg. of silver, has been selected as the “practical unit” and is called the “coulomb” 
 
 (coul.) ; 300 mg. of silver deposited iq^on the kathode indicate therefore that ^^^^^ = 268 
 coni, have passed through the solution. If this took place in ten minutes, then in one 
 second ^ = 0.447 coul. would have passed through the liejuid. 'I'he current strength 
 would then have been 0.447 miipere. 'I'he current strength (amp.), therefore, represents 
 
SOLUTIONS. 
 
 265 
 
 the number of coulombs which pass through the entire cross-section of the current’s 
 course. A current of one ampere deposits or separates : 
 
 Mg. Silver. Mg. Copper. 
 In one second, . . 1.118 0.3284 
 
 In one minute, . . 67.08 i 9 - 7 o 
 
 In one hour, . . 4025 1182 
 
 C.c. Electro- 
 lytic Gas C.c. Hydrogen 
 
 AT 0° AND AT 0° AND 
 
 Mg. Water. 760 Mm. 760 Mm. 
 
 0.0933 0.1740 o. n6o 
 
 5.60 10.44 6.960 
 
 335-9 626 417 
 
 The electro-chemical equivalents are the quantities whicli are separated out or converted 
 into other compounds in the unit of time. It is possible in accordance with Faraday’s 
 law to calculate readily from the silver and copper values those of other elements 
 or atomic groups. In general, a current of i ampere deposits in i sec. 0.01036, in i 
 
 min. 0.6215, ^ hour 37.29 mg. — equivalents of any substance ~ 
 
 0.01036 ; Ag = 107.93 ; “ — 31-^) » hence, 35.45 . 0.01036 — 0.3673 mg. chlorine in a 
 second, etc. It may also be mentioned that the practical unit of resistance, the ohm, is a 
 column of mercury 106.3 cm. in length and i .square mm. in section, and that the volt, the 
 practical unit of electromotive force, is the power which in a closed circuit of an ohm 
 resistance produces a current of i ampere. 
 
 Faraday’s law requires that when equal quantities of electricity pass through different 
 electrolytes, the same number of valences must be liberated or converted into new com- 
 pounds at each pole. Therefore, the decomposition of aqueous hydrochloric acid is 
 accomplished by the consumption of the same quantity of electricity which- would be re- 
 quired by equivalent quantities of hydrobromic acid, hydriodic acid, etc. We know on 
 the other hand that very different amounts of heat are necessary to effect the decompo.si- 
 tion of equivalent quantities of these compounds. Hence we might expect that their 
 electrolytic decomposition would demand a varying expenditure of electricity. The 
 endothermic bodies should be most easily decomposed, while those with the greater heat 
 of formation and heat of decomposition would be most difficult. 
 
 Furthermore, so long as the electric force, active in the conductor, is not equal in 
 strength to the affinity, decomposition does not occur, but when these forces are equal, 
 then many molecules will at once be broken down. This contradicts all experience 
 which, according to R. Clausius, may be explained as follows [Fogg. Annalen (1857) 
 loi, 338; die mechanische Warmetheorie (1879) Ed. ii, 164] : The current does not 
 decompose the electrolyte into its constituents. In consequence of the heat motions of 
 its molecule the electrolyte is already partially dis.sociated and the current merely separates 
 the already free ions spacially. Williamson, at an early date, had expre.s.sed a similar 
 view in an unusually able paper on the theory of the ether-formation [Ann. Chem. 
 Fharm. (1851) 77, 37]. He thought that in a mass of molecules the latter continually 
 exchanged the atoms and atom groups, of which they consisted, with one another. For 
 example, in hydrochloric acid the molecule of hydrogen chloride did not constantly con- 
 sist of the same hydrogen atom and chlorine atom, but that there was rather a mutual 
 exchange of corre.sponding atoms among the molecules. S. Arrhenius has in late years 
 expanded and developed these ideas into the theory of electrolytic dissociation, which 
 will receive attention later, p. 268. 
 
 SOLUTIONS. 
 
 Most liquids can take up gases, other liquids and solids in such a man- 
 ner that a new liquid, homogeneous both physically and chemically — a 
 solution — is produced (p. 256). No deep-seated chemical changes occur 
 between the solvent and the substance dissolved, otherwise the process 
 
 23 
 
266 
 
 INORGANIC CHEMISTRY. 
 
 could not be defined as solution in the sense indicated. If, for example, 
 carbon dioxide is conducted into and absorbed by caustic potash, a solu- 
 tion of carbon dioxide does not result, but rather one of potassium car- 
 bonate. In the same way we cannot say that a metal dissolves in an 
 acid, because it is first changed, with the evolution of hydrogen, into a 
 salt which dissolves. Gases can mix to any degree with one another, but 
 the solvent power of liquids for gases or other liquids or solids is gener- 
 ally limited and also dependent upon jiressure and temperature. Aqueous 
 solutions are alone important here, and only they will be considered 
 now. 
 
 The absorption of gases (Henry, 1803) or gas mixtures (Dalton, 1807), 
 as mentioned on p. 229, is proportional to tlie pressure under which the 
 gas exists. The gas dissolved by a liquid can be expelled from it by 
 raising the temperature, by lowering the pressure, or by introducing 
 another gas, e. g., carbon dioxide, by passing air through the solution, or 
 by shaking the latter with air. Certain gas solutions are distinguished 
 by the fact that under a definite pressure they will distil without change 
 in their composition at a constant temiierature ; (.see Hydrochloric Acid, 
 p. 59, Hydrobromic Acid, p. 62). 
 
 Many metallic salts separate in union with water — water of crystalliza- 
 tion — from their solutions. This can frequently be noticed in their 
 change of color. Anhydrous copper sulphate, CuSO^, is white, while 
 its aqueous solution is blue, and on evaporation blue hydrous CuSO^ . sH^O 
 crystallizes out. This salt loses a portion of its water, on exposure to 
 the air, at the ordinary temperature — it effloresces. It becomes anhydrous 
 above 200°. Hence, we can conclude that such a hydrate existed in the 
 solution. In some cases the metallic salt, depending on the physical 
 conditions at which the solution or crystallization occurred, unites with 
 varying amounts of water, which is recognizable by its color. The com- 
 pounds of cobaltous chloride, CoCl.^, with one and two molecules of water 
 of crystallization are blue in color like the solution from which they 
 separate, whereas the red compound with six molecules of water, 
 CoClg + 6H2O, is obtained from red-colored solutions. As a rule, 
 the solubility of salts increases with the rise in temperature and is 
 definite for every temperature {saturated soliitiofi). Thus, i part of 
 potassium nitrate dissolves in 4 parts of water at 15°, and in 0.3-0. 4 
 parts at the boiling temperature. When such a hot saturated solution 
 cools, the salt crystallizes out. Sodium chloride is a remarkable excep- 
 tion to this, as it dissolves very little more in hot than in cold water: 
 100 ])arts of water at the ordinary temperature take up 36 parts, and when 
 boiling 39 jiarts of the salt. The solubility with other salts increases up 
 to a definite temperature and then it diminishes. It is thus with sodium 
 sulphate, Na.^SO^ ; 100 parts of water at 0° dissolve five parts of the salt, 
 and the most, namely 55 jiarts, at 34° ; at 60° only 45 and at 100° about 
 42 parts (see Glauber’s Salt). When a salt dissolves it absorbs heat and 
 the water cools. However, if the salt, like anhydrous copper sulphate 
 or calcium chloride, (kaCl.^, can take u]) water, i. e., able to form a 
 hydrate, then heat is evolved in its solution, because the quantity of 
 heat, set free in the formation of the hydrate, exceeds that necessary to 
 
SOLUTIONS. 267 
 
 dissolve it. Therefore, hydrated calcium chloride, CaCl^ -f 6H2O, dis- 
 solves in water with the lowering of temperature. 
 
 THEORY OF DILUTE SOLUTIONS. 
 
 Many chemists regard solution as a chemical process induced by the chemical attrac- 
 tion between the solvent and the substance dissolved, and leading finally to very unstable 
 and undetermined chemical compounds (Berthollet, Mendelejeff, Berthelot), while others 
 believe it to be physical, e. g., Dossios, who considers it a sort of diffusion, which may be 
 compared in every respect to vaporization. Neither of the two views alone is sufficient 
 to explain satisfactorily the process of solution [see Z. f. anorg. Ch. 6 (1894) 392]. 
 
 While we cannot be certain as to the forces which bring about solution, yet since 
 1885 J. H. van’t Hoff has by his pioneer researches shown that the condition of a sub- 
 stance in dilute solution is similar to the gas condition. The equation deduced on 
 page 12 1 : 
 
 pv = RT 
 
 embodies the law of gases (Boyle and Gay-Lussac) and that of Avogadro, and as 
 proved by van’t Hoff it also answers for dilute solutions if “osmotic pressure” be 
 substituted for gas pressure [Z. f. phys. Ch. i (1887) 487; 3 (1889) 198; Ber. 27 
 (1894) 6], 
 
 Osmotic pressure is the pressure exerted against a diaphragm, separating a solu- 
 tion from a solvent, and w'hich is permeable to the solvent but not to the dissolved 
 substance (see p. 237). Changes brought about by osmotic pressure play a very 
 important role in the living organism and particularly in plants, therefore they 
 were first more accurately studied by botanists — W. Pfeffer in 1877 and H. de Vries 
 in 1884. 
 
 The action of osmotic pressure may be demonstrated as follows : A cylinder of about 
 TOO c.c. capacity is filled with a syrupy sugar solution and then made air-tight with a 
 covering of animal membrane. When this vessel is immersed in a vessel of water, the 
 membrane gradually forms a curved surface, because the water penetrates through it to 
 the sugar solution. The experiment can also be conducted as follows: An open 
 cylinder is closed at one end with a membrane and then filled with a concentrated sugar 
 solution, the other end of the cylinder being provided with a stopper which carries a 
 narrow glass tube. When the cylinder is immersed to the stopper in water, the latter 
 penetrates the membrane and as a consequence the liquid rises gradually in the tube to 
 a definite height ; thus a means, though not very exact, is afforded for the measurement 
 of the magnitude of the osmotic pressure. 
 
 Pfeffer and de Vries made measurements of osmotic pressure and found the following 
 relations : 
 
 1. Osmotic pressure is proportional to the concentration or inversely proportional to 
 the volume in which a definite amount of substance is dissolved. 
 
 2. Osmotic pressure increases with constant volume in proportion to the absolute tem- 
 perature. 
 
 3. Quantities of dissolved substances, which are to one another as their molecular 
 weights, exert equal osmotic pressure when dissolved in equal volume at like temperatures. 
 
 One of van’t Hoff’s services was the recognition that these three laws of osmotic pres- 
 sure correspond to the laws controlling gases. The first is the counterpart of Boyle’s 
 law, the second, of the law of Gay-Lussac and Dalton, and the third corresponds to 
 Avogadro’s law. 
 
 Let us apply these relations to a solution which contains one part by weight of cane 
 sugar in lOO parts by weight of water. In accordance with the specific gravity of such 
 a solution it would contain i gram of sugar in 100.6 c.c. Pfeffer’ s experiments show 
 that this solution exerts at 0° an osmotic pressure which would hold a mercury column 
 of 49.3 cm. in height in equilibrium. If it is desired to ascertain whether these relations 
 can be expressed by the equation p . v = 84800 T, the units mentioned on p. 1 22 must be 
 chosen and the pre.ssure in grams per square centimeter, the volume in cubic centimeters 
 
268 
 
 inorc;anic chemistry. 
 
 for the gram-molecule and the temperature in absolute numbers must be considered in the 
 calculation.* 
 
 In this particular instance the value for p would be 49.3 . 13.6 670.48 ; for '1' the 
 
 value 273. Cane sugar, CjjIIj./),,, has the molecular weight 342.18 ; tlie gram-molecule 
 of the solution in <|uesti()n is then contained in 342.18 . 100.6 — 34400 c.c. (v) approxi- 
 mately. Hence 670.48 . 34400 - - R . 273 and R 84500, a value which accords 
 within the limits of error with that of the gas constants. 
 
 Conseciuently dissolved substances exert the same pressure, as o.smotic pressure, which 
 they would exert at like temperature and like volume in the gas form. 'I’heir mole- 
 cules are, as in the case of gases, .so far removed from one another that only the proper- 
 ties dej)en(lcnt upon their number manifest themselves, while those dej)ending uix)n com- 
 position and chemical structure ofttimes disappear absolutely. Properties of this kind, 
 which can assume like values for chemically com[)arable quantities of the most different 
 substances — h)r molecules, for atoms of bodies which are exceptions to Avogadro’s law 
 (P- 79 )> ions, as we shall .soon .see, are termed, following Ostwald, colligative proper- 
 ties [co/ligare, to link together). The f|uantities of two substances, which in dilute .solu- 
 tion or in the gas form manifest colligative properties of ecjual value, will be to one an- 
 other as the ratio of their molecular weights if the exce])tions indicated be not considered. 
 It would then follow : that if at one time n-gram-molecules of a substance A be di.ssolved 
 in a definite volume of a solvent, and then n-gram-molecules of the substance P>, in both 
 cases certain properties of the original .solvent would be similarly altered. The freezing 
 point will fall and the boiling point will rise regularly in both instances ; the tension of 
 the two solutions will be the same ; they pos.sess eipial osmotic jiressure, ?. e., they are 
 isotonic {^Igoc, equal ; Tovn^, tension). It is in this manner possible, upon comparing solu- 
 tions of a substance of unknown molecular weight with those of one of known molecular 
 weight, to determine the unknown molecular weight: (l) by isotony (Pfeffer, de Vries) ; 
 (2) by lowering of tension (Raoult) ; (3) by the rise in boiling point ( Reckmann-Arrhe- 
 nius) ; (4) by lowering of the freezing point (van’t IIoff-Raoult-Eykmann). The.se 
 methods are described in Richter’s Organic Chemistry. To show how their properties 
 change in proportion to concentration, mention may be made of the following : A .solution 
 of I part of common salt in 100 parts of water begins to freeze at — 0.6°, one of 2 : 100 
 at — 1.2°, of 4 : 100 at — 2.4°, of 14 ; too at — 8.4°, i. e., at — 0.6 X Relow the 
 
 last-named temperature the proportionality does not hold for anhydrous sodium chloride, 
 but for the hydrate NaCl 2 H.pj. In the case of other .salts the hydrates play this role 
 at higher temperatures: thus, Nal 4H2O, MgSO^ 7H2O, etc. [Rudoidf and de 
 Coppet]. 
 
 THEORY OF ELECTROLYTIC DISSOCIATION. 
 
 The laws just mentioned serve only for very dilute solutions, just as the laws for gases 
 correspond with greater accuracy in the behavior of those gases which are far removed 
 from the vapor condition (p. 122). And even very dilute solutions show deviations from 
 van’t Hoff’s theory when the solvent is water and the dissolved .substance is an electrolyte 
 — a conductor of the second class (p. 262). The salts, acids and bases are electrolytes 
 in aqueous .solution, while the non-electrolytes are the organic compounds, with the ex- 
 cei)tion of the pronounced acids, bases and salts ; also the solutions of all substances in 
 benzene, carbon bisulphide, ether, chloroform, etc. The alcoholic solutions constitute a 
 transition to the electrolytes. 
 
 An example will illustrate the.se important relations. While the aqueous solutions of 
 ether, glycol, sugar, urea and similar non-conducting organic compounds, which contain 
 the molecular weight of the respective sub.stances expressed in grams per liter, freeze at 
 — 1.8°, the freezing point of the corresponding solutions of potassium iodide, sodium 
 chloride and silver nitrate falls to — 3-6°, double the first ; and that of .such a solution of 
 sodium sulphate to — 5.4°, treble the first. The values of other colligative properties of 
 these solutions manifest similar variations from the van’t Hoff theory : rise in the boiling 
 
 * If the |)rcssiire be given in atmospheres and the volume in liters, then R naturally 
 acquires another value. As the molecular volume (p. 98), at l atmosphere and 0°, 
 equals 22.4 liters, then R would e(|ual 0.082 (because p ::= l, v = 22.4, T = 273) ; 
 lienee in general [> . v 0.082 'f (liter-atmosphere). 
 
SOLUTIONS. 
 
 269 
 
 point, decrease in the vapor tension, and the osmotic pressure is also double or treble 
 that required by the theory. Dilute solutions of electrolytes consequently behave as if 
 they contained more molecules of the dissolved substance than the corresponding solu- 
 tions of non-electrolytes. 
 
 Svante Arrhenius’ theory of electrolytic dissociation [Z. f. phys. Ch. i (1887) 631 ; 
 see also Planck, ibid. 576] offers an interpretation for this abnormal deportment. 
 According to it the electrolytes — which Hittorf designates salt-like bodies inasmuch as 
 the acids are salts of hydrogen and the bases of hydroxyl — do not exist as such in aqueous 
 solution but have broken down entirely or in part into their ions (p. 265) — NaCl into 
 + — + — + — + — 
 
 Na and Cl, AgNOg into Ag and NO.^, Na2S04 into 2Na and SO4, and KOH into K and OH. 
 The anions are charged with negative and the kations with positive electricities which, in 
 accord with the law of Faraday, are equally large for equivalent quantities of different 
 ions, e. g., the ion SO4 contains twice the quantity of electricity of the ion Cl. 
 
 The ions regarded as independent mass-particles for comparison’s sake play the part of 
 molecules and accordingly influence the values of the colligative properties. This would 
 explain the exceptions, cited above, to the theory of solutions; the molecule NaCl, 
 + — _ -f 
 
 resolved into the ions Na and Cl, acts like two molecules; Na^vSO^ separated into 2Na 
 
 and SO4 behaves like three molecules of a non-electrolyte. The determination of the 
 lowering of the freezing point, etc., can therefore serve in determining the degree of dis- 
 sociation. Dissociation and electric conductivity run parallel because the conduction of 
 the current is influenced only by the presence of free ions and their quantity, whereas 
 molecules which are not dissociated do not participate in conducting the current. 
 Accordingly the conductivity actually increases, if molecular quantities be considered, 
 with growing dilution and attains its maximum when all the molecules of the electrolyte 
 are dissociated. 
 
 The neutral salts are most strongly dissociated and particularly those with univalent 
 ions, e. g.^ NaCl, AgNOg, KI, NH4Br. Usually more than half of the salt is present 
 in the form of free ions in aqueous solutions at medium concentrations. Salts with poly- 
 valent ions are less completely dissociated, and in the case of the mercury haloids the 
 dissociation is extremely slight. The degree of dissociation with acids and bases corre- 
 sponds to what is commonly called their ‘‘strength” ; the strongest bases and acids are 
 most completely dissociated. A gradual dissociation occurs with the polybasic acids : 
 
 + — 
 
 thus sulphuric first breaks down into the ions H and HSO4 ; the univalent anion HSO4 
 
 — -I- 
 
 in turn dissociates into SO4 and H. 
 
 The strong acids would then be : Hydrochloric, hydrobromic, hydriodic, nitric, 
 chloric, perchloric and sulphuric, as well as the polythionic acids. Strong bases : Alka- 
 lies, alkaline earths and the oxide of thallium. In solutions of medium concentration 
 these compounds are dissociated more than half. 
 
 Moderately strong acids: Phosphoric, suljDhurous and acetic. Moderately strong 
 bases : Ammonia, magnesia and silver oxide (dissociation does not exceed 10 per cent.). 
 
 Weak acids and bases, the dissociation of which is in part scarcely measurable : Car- 
 bonic acid, hydrogen sulphide, hydrogen cyanide, silicic acid, boric acid — the hydroxides 
 of the other bivalent and trivalent metals. 
 
 The formation of a salt from a base and an acid proceeds, according to this theory, in 
 aqueous solution as follows : The base breaks down into the metallic kation and the 
 anion OH, the acid into th« kation H and the anion — the acid residue j)reviously com- 
 bined with it. Because water j)ossesses an extremely slight degree of dissociation, being 
 a non-conductor (a gram equivalent of its ions is contained in about 10,000,000 liters), 
 it will be produced whenever the ions H and OH meet, so that in the present case, if we 
 have started with equivalent quantities, only the ions of the salt will remain in solution : 
 
 Na + OH -f Cl + H = Na + Cl + H^O + 13-7 Cal. 
 
 K + OH + N^ +^H = K -f NO3 -f H2O -f 13.7 Cal. 
 
 Ca 4- 2OH j 2CI -f 211 = Ca -f 2CI 4 - 2H2O 4- 27.4 Cal. 
 
270 
 
 INORGANIC CHEMISTRY. 
 
 'I'he common and essential thing in these changes is the prodnctioti of water from the 
 ions 11 and (Jll ; the other ions remain unchanged as long as the concentration of tile 
 solution is not altered. A fact in liarmony with this is that upon fieufraiizin^ equivalent 
 quantities of strong bases with strong acids an equal heat modulus (thermal value) is 
 obtained ; hence, 
 
 (II acp, OH aq.) = 13.7. 
 
 This fact, for which no reason could be previously observed, now seems in a similar man- 
 ner to be demanded by and also to conhrm the theory just discussed. 
 
 d'he salt-like bodies according to the theory of electrolytic dissociation appear to be 
 binary in their constitution. Berzelius’ old electro chemical theory also made them con- 
 sist of electro-positive and electro negative parts — they were dualistically constituted. 
 The salts of the oxygen acids consisted of the acid anhydride and the metallic oxide, e.g., 
 — 
 
 potassium sulphate of SO.,. — whereby they were brought in opposition to the haloid 
 
 salts, which facts did not justify, d'he electro-chemical theory of Arrhenius proceeding 
 from positive and negative constituents has a dualistic character, attaches itself thereby to 
 the idea of Daniell and of Liebig (p. 260), and corresponds to the views to which Wil- 
 liamson and Clausius were led, the former for chemical and the latter for electrolytic 
 reasons (p. 265). d'his theory at first appeared strange and met great oj^position, but 
 it is now almost universally accepted and it is generally admitted that we are indebted to 
 it for many great and surprising discoveries in the domain of chemistry and physics which 
 had gone unobserved. It appears almost absolutely necessary for young chemists to 
 make themselves conversant with this theory and to that end the publications cited on 
 pp. 49, 66 will be found helpful. The following will also prove especially valuable : 
 ( Istwald, Foundations of Analytical Chemistry ; Liipke, Elements of Electro chemistry ; 
 Le Blanc, Electro-chemistry ; Lob, The Elements of Electro chemistry ; Windisch, 
 Determination of the Molecular Weights. 
 
 TRANSPOSITION OF SALTS. 
 
 When two salts in solution or fusion come together, a chemical action will frequently 
 occur. Claude Louis Berthollet (Es.sai de statique chimique, 1803) endeavored to ex- 
 plain the resulting phenomena by referring them to purely physical causes, and excluded 
 chemical affinity. 
 
 In the opinion of Berthollet, four salts always arise in the solution of two. For 
 example, on mixing solutions of copper sulphate and sodium chloride, there exist in solu- 
 tion copper sulphate, sodium sulphate, copper chloride, and sodium chloride : 
 
 nCuSO^ -j- mNaCl yield 
 
 (n — x)CuSO„ (m — 2x)NaCl -)- xNa2SO„ -j- xCuCl2. 
 
 That copper chloride is really present in the solution together with the sulphate, 
 follows, from the fact that the blue color of the latter acquires a greenish color, peculiar 
 to the copper chloride, by the addition of sodium chloride ; other phenomena are not 
 noticeable at first. Suppose one of the four salts formed in the solution is insoluble or 
 volatile, the reaction will occur somewhat differently. Upon adding barium chloride to 
 the copper sulphate solution four salts will be formed at the beginning just as in the first 
 case. The barium sulphate produced separates, however, in consequence of its insolu- 
 bility, the ecjuilibrium of the four salts will be disturbed, and new quantities of copper 
 sulphate and barium chloride act upon each other until the transposition is complete : 
 
 CuSO, -f BaCl2 = BaSO^ -f CuCl.,. 
 
 'I'he chemical transposition may, therefore, be explained by the insolubility of the barium 
 sulphate. On adding hydrochloric acid, or soluble chlorides, to the solution of a silver 
 salt all the silver is precipitated as chloride, becau.se the latter is in.soluble. 
 
 'I'ake another example. On adding sulphuric acid to a solution of potassium nitrate 
 there is apparently no pcrce])tible alteration ; yet four compounds, KNO3, KIISO4, H2SO4 
 and UNO,,, are [)re.sent in tlie .solution. This was proved by the thermo-chemical invest!- 
 
TRANSPOSITION OF SALTS. 
 
 271 
 
 gationsof Julius Thomsen, and from determinations made by W. Ostwald on the changes 
 in volume and density which are connected with transpositions. The two acids distribute 
 themselves upon the base. The proportion or degree to which this occurs is dependent 
 upon the quantity of potassium nitrate and sulphuric acid in a volume unit, from external 
 circumstances, such as the temperature, and upon the nature of reacting substances. The 
 more sulphuric acid there is in proportion to the nitric, the more sulphate will there be 
 formed. This is a case of mass-action^ the theory of which was developed by Guldberg 
 and Waage (p. 95). On heating or evaporating the nitrate solution containing sulphuric 
 acid a new condition arises : the volatility of nitric acid. Hence it follows that on 
 evaporation the transposition proceeds to completion in the sense of the equation : 
 
 2KNO3 -f 2H2SO4 = KHSO^ + 2HNO3. 
 
 These relations are thus explained by the theory of electrolytic disso- 
 ciation. Neutral salts in dilute solutions do not as a rule act upon one 
 another because, like the salts which can arise from them by transposition, 
 they are equally dissociated. Thus a dilute solution of potassium 
 chloride contains the ions K and Cl, that of sodium nitrate, the ions of 
 Na and NO,. If these solutions are mixed no change occurs. A solu- 
 tion of exactly the same properties would be obtained on mixing corre- 
 sponding amounts of potassium nitrate and sodium chloride. 
 
 If a substance, less soluble and less dissociated under the prevailing conditions, can 
 be formed from the ions, then a transposition, apparently an exchange, will occur. It is 
 in this fashion that the transpositions described above are to be explained. The neutral- 
 ization of bases by acids is to be accounted for in this way : water, which is but slightly 
 dissociated, is formed from the ions H and OH, whereas the ions which belong to the 
 salt remain (p. 269). The expulsion of a weak acid from its salts by a stronger acid 
 is similarly completed. While the neutral salts are approximately equally dissociated, 
 feeble acids on the other hand are changed but little. If therefore a strong acid be 
 added to the solution of such a salt its hydrogen atoms will come together with the 
 anions of the salt and will combine more or less completely with them to an acid which 
 is not dissociated ; there will remain the kation of the salt and the anion of the added 
 acid, e. g.^ hydrochloric acid and sodium borate: 
 
 — + — + -l- — 
 
 Cl + H + BO2 + Na = HBO2 + Na -f Cl. 
 
 Salts of weak acids or of feeble bases are also hydrolytically decomposed by the action 
 of water, i, e., they break down into base and acid, of which the weak portion exists 
 undissociated in the solution. The base or acid dissociated to a greater extent may be 
 recognized by the fact that its solution reacts acid (with feeble base) or alkaline (with 
 feeble acid). 
 
 The reactions employed to detect substances depend, according to the theory of elec- 
 trolytic dissociation, chiefly upon reactions of ions. All compounds, for example, which 
 in aqueous solution yield the anion Cl, show the reaction of hydrochloric acid, so far as 
 they produce a precipitate of silver chloride with silver nitrate. When chlorine does not 
 appear alone as an ion, but as a part of .such, this reaction does not take place. The 
 compound Na.2PtClg, the .sodium salt of hydrochlorplatinic acid, rich in chlorine, does 
 not yield a precipitate of silver chloride with silver nitrate, because in arpieous solution 
 it dis.sociates into the anion PtClg and the kations 2Na. The color of the solution is 
 al.so materially affected by the ions. The following therefore is accordingly explained : 
 
 When ferric chloride and potassium fluoride meet in solution in equivalent amounts a 
 whole series of characteristics belonging to ferric chloride solutions disappear : the solu- 
 tion is colorless ; iodine is not set free from potassium iodide even after the addition of 
 an acid ; potassium .sulphocyanide, salicylic acid and substances which otherwise detect 
 ferric chloride with great accuracy, fail to show anything. Formerly there was no ex- 
 planation for the deep-seated transposition which had evidently occurred in the liquid, 
 but it may be found according to the theory of electrolytic dissociation in the fact that 
 
272 
 
 INORGANIC CHEMISTRY. 
 
 ferric chloride and iwlassiuin fluoride arc coinj)lelely transposed to ferric fluoride anfl 
 potassium chloride : 
 
 FeCl., I 3KKI 3KC:1 1 Feld,, 
 
 and ferric fluoride is not dissociated. 'I'he trivalent iron ions, u|)on which the reactions 
 recorded above are dependent, are no lonj^er present in the solution ; therefore, the reac- 
 tion cannot occur. This transposition completes itself .so fully that it can be applied in 
 the quantitative determination of fluorides by working with ati excess of ferric chloride and 
 then determining the portion, not converted into fluoride, with potassium iodide, upon 
 which ferric fluoride does not act (Knobloch). Similar relations will be cited in the 
 following pages. 
 
 I. GROUP OF THE ALKALI METALS. 
 
 Potassium, 39-15 
 
 Rubidium, 85.4 
 
 CcTesium, 133 
 
 Lithium, 7-03 
 
 Sodium, 23.05 
 
 (Ammonium, NH^ = 18.07) 
 
 The metals of this grotij^ are decidedly the most pronounced in metallo- 
 basic character, and this constitutes a visible contrast with the elements 
 of the chlorine group, the most energetic among the acid-forming metal- 
 loids. 
 
 The alkali metals in physical and chemical proj)erties exhibit great 
 similarity. They oxidize readily in the air, decompose water violently, 
 even in the cold, with the formation of strong basic hydroxides, which 
 dissolve readily in water and are called alkalies (caustic potash, caustic 
 soda) ; hence the name alkali metal {al kaljiin, Arabic, meaning the ash 
 of sea and beach plants, and the extract from the same). They are not 
 decomposed by ignition. Their chemical energy increases with increas- 
 ing atomic weight (more correctly atomic volume), sodium is more ener- 
 getic than lithium, potassium more than sodium, and rubidium more than 
 potassium. Caesium has not been studied in the free condition, but, judg- 
 ing from its compounds, it possesses a more basic character than rubidium. 
 We saw in other analogous groups (of chlorine, oxygen, phosphorus, car- 
 bon, and similar elements), that the metalloidal, electro-negative character 
 diminishes, and the basic increases with the increasing atomic weight. 
 
 The specific gravities increase simultaneously with the atomic weights; 
 but as the increase of the latter is greater than that of the former, the 
 atomic volumes (the quotients p. 252), are always greater. The in- 
 
 creasing fusibility and volatility correspond to the increase of the atomic 
 volumes ; rubidium distils at a red heat, while lithium volatilizes only 
 with difficulty : 
 
 
 Li 
 
 Na 
 
 K 
 
 Rb 
 
 Cs 
 
 Atomic weight, 
 
 Specific gravity (15°), - - 
 
 Atomic volume, .... 
 
 J^'usion teitiperature, . . . 
 boiling temperature, . . . 
 
 7-03 
 
 0.59 
 
 1 1.9 
 
 180° 
 
 23-05 
 
 0.97 
 
 23-7 
 
 96° 
 
 742° 
 
 39-15 
 
 0.87 
 
 45 
 
 62.5° 
 
 667° 
 
 ^ 5-4 
 
 1-52 
 
 56.1 
 
 3 ^- 5 ° 
 
 133 
 
 1.88 
 
 70.7 
 
 26.5° 
 
 270° 
 
POTASSIUM. 
 
 273 
 
 Although the alkali metals exhibit a great similarity in their chemical 
 deportment, we discover more marked relations between potassium, rubid- 
 ium and caesium upon the one hand, and lithium and sodium on the 
 other, which accords with their position in the periodic system of the 
 elements (p. 246). Especially is this noticed in the salts. The first 
 three metals form difficultly soluble tartrates and chlorplatinates (see 
 Platinum). Their carbonates deliquesce in the air, while those of sodium 
 and lithium are stable under similar circumstances ; the last is, indeed, 
 rather insoluble in water. The phosphates deport themselves similarly; 
 lithium phosphate is very difficultly soluble. It must be remarked that 
 the normal carbonates and phosphates of all other metals are insoluble in 
 water. In lithium, then, which possesses the lowest atomic weight, it 
 would seem the alkaline character has not reached its full expression, and 
 it in many respects approaches the elements of the second group, espe- 
 cially magnesium, just as beryllium approaches aluminium. The elements 
 of the two small periods (lithium and sodium) are, indeed, similar, but 
 not completely analogous, while the homology of the three great periods 
 finds expression in potassium, rubidium and caesium. 
 
 Of the thermo-chemical relations of the alkali metals only the heat of formation of some 
 hydrates will be given here : 
 
 (Li, O, H, Aq) = 117.4 (Na, O, H, Aq) ^ 111.8 (K, O, H, Aq) = 116.4. 
 
 On comparing these values with the heat of formation of water ( Hg, O) = 68.36 Cal., 
 we immediately perceive why it is so readily decomposed by the alkali metals. All 
 metals, disengaging more than 68.3 Cal. in the formation of their oxides, Me^O, or their 
 hydroxides, MeOH, decompose water, and the energy will be greater, the greater the 
 difference of heat. The insolubility of the oxides constitutes an obstacle to the action ; 
 this, however, may be removed (see Aluminium) by addition of neutral solvents. Con- 
 versely, all oxides, affording less heat in their formation, are easily reduced by hydrogen 
 (pp. 92, 95). 
 
 1. POTASSIUM. 
 
 K = 39-15- 
 
 In nature, potassium is found principally in silicates, viz. : feldspar 
 and mica. By the disintegration of these frequently occurring minerals, 
 potassium passes into the soil, and is absorbed by plants ; the ashes of the 
 latter consist chiefly of different potassium salts. The chloride and sul- 
 phate are also found in sea-water, and in large deposits in Stassfurt, at 
 Magdeburg, and in Brunswick and in Galicia, where they were left by the 
 evaporation of the water of inclosed seas. 
 
 Metallic potassium was first obtained by Davy, in the year 1807, by the 
 decomposition of the hydroxide, by means of a strong galvanic current. 
 At present it is prepared by igniting an intimate mixture of carbon and 
 potassium carbonate : 
 
 K.COj T 2C = 2K + 3CO. 
 
274 
 
 INORGANIC CHEMISTRY. 
 
 Such a mixture may be made l)y the carl)Ouizatiou of organic potassium 
 salts, <f. , crude tartar. It is tlieu ignited to wliitc lieat, in an iron 
 
 retort, and tlie escaping vapors collected in receivers of ]>eculiar construc- 
 tion, filled with rock oil, 'I'he latter, a mixture of hydrocarbon, serves 
 as the best means of preserving potassium, which would otherwise oxidize 
 in the air, and dccom[)ose other licpiids. Potassium carbon monoxide 
 (KC())p is a by-product in the preparation of the metal ; see Richter’s 
 Organic Chemistry. 
 
 In a fresh section, potassium shows a silver-white color and brilliant 
 metallic luster. At ordinary temperatures it is soft, like wax, and may 
 be easily cut. It crystallizes in octahedra, and has a specific gravity of 
 0.87 at 15°. It melts at 62° and boils at about 667°, and when raised to a 
 red heat, is converted into a greenish vapor. It oxidizes in the air, and 
 becomes dull in color ; heated, it burns with a violet flame. It decom- 
 poses water energetically, with formation of iiotassium hydroxide and the 
 liberation of hydrogen. If a piece of the metal be thrown upon water, 
 it will swim on the surface with a rotary motion ; so much heat is dis- 
 engaged by the reaction that the generated hydrogen and the potassium 
 inflame. Finally, a slight explosion usually results, whereby pieces of 
 potassium are tossed here and there; it is advisable, therefore, to execute 
 the experiment in a tall beaker glass, covered with a glass plate. Potas- 
 sium combines directly and very energetically with the halogens. 
 
 On conducting hydrogen over metallic potassium heated to 300-400°, 
 potassium hydride, K2H (or K^H2), results. This is a metallic, shining, 
 brittle compound, which, upon stronger heating (above 410°), more 
 readily in vacno, is again decomposed. The sodium hydride, Na^H2, 
 obtained in the same way, does not ignite spontaneously, but in other 
 respects is very much like potassium hydride. 
 
 The influence of heat and pressure in the formation and decomposition of these com- 
 pounds is very noteworthy. If, for example, potassium hydride be heated it melts, but 
 otherwise remains unchanged. Above 2DO° (in a vacuum) it sustains a partial decom- 
 position (dissociation), which gradually increases as the temperature rises. If the heating 
 should take place in a closed vessel provided with a manometer, it will be observed that 
 the decomposition at a given temperature will continue until the liberated hydrogen has 
 acquired a definite tension — until it exerts a definite pressure. For potassium hydride, 
 this tension, at 330°, equals 45 mm. The decomposition will then cease, but will pro- 
 ceed further at the same temperature if the hydrogen gas be removed, until the pressure 
 of 45 mm. is again reached. This pressure is therefore called the dissociation tension. 
 In this manner a complete decomposition of the hydride may be effected at the tempera- 
 ture given above. If, however, the di.sengaged hydrogen is not removed, but be added 
 to the completely or jrartially decomjrosed hydride, and the pressure be raised to 45 mm. 
 (at the tem])erature 330°), the jiotassium hydride will be re-formed. Consequently, both 
 the decomposition and the formation of a body can follow, dejrending uj)on whether the 
 external partial ])ressure be lowered or increased. Similar phenomena occur at higher 
 temperatures, tlie corres|)ondiiig pressure, of course, increasing by regular steps. 
 
 'I'he tension of dissociation is independent of the relative (juantity of the dissociated 
 body and of the s[)ace which the di.sengaged gas can occupy, whereas in solutions and 
 absori)li()ns (ammonia by charcoal) the pressure at one and the same temperature 
 increases with (he (|uantity of the absorbed gas. 
 
 All exothermic compounds behave like potassium and .sodium hydrides when they are 
 decomposed into their eoin|)()nents ; if the pressure be raised above the tension of disso- 
 ciation the components reunite — the compounds are re-formed. The decomposition of 
 
POTASSIUM HYDROXIDE. 
 
 275 
 
 the endothermic compounds (potassium chlorate into chloride and oxygen) is quite dif- 
 ferent (pp. 30, 94). It proceeds with heat disengagement (KCljO.^ = — '9.8 Cal.), 
 corresponding to the chemical affinities, and is only induced by application of external 
 heat. It is independent of external pressure, and there is no reunion of the decomposi- 
 tion products upon increasing the external pressure or upon lowering the temperature. 
 
 We must not omit mentioning the great analogy between the phenomena of dis- 
 sociation and the vaporizing of liquids. Like dissociated bodies, liquids exhibit at all 
 temperatures a definite vapor tension. If the pressure above the liquid be diminished 
 the evaporation will continue until the vapor tension is regained, but if the pressure be 
 increased then a corresponding portion of the vapor will be condensed. 
 
 Oxygen compounds of potassium are not known in a pure state. The 
 metal is not attacked by pure, dry oxygen below 60-80°. Above this 
 temperature it burns in the gas, if its surface be renewed, to a yellow 
 mass which, according to Erdmann and Kothner, consists of the sesqui- 
 oxide, K2O3, and superoxide, KO2. The oxide K2O, from which the 
 potassium salts are derived, is not definitely known. 
 
 Potassium Hydroxide, o,r Caustic potash {^Kalium causticuni), 
 KOH, is obtained by the action of potassium or its oxide upon water. For 
 its preparation, potassium carbonate is decomposed by calcium hydroxide 
 (slaked lime) ; 
 
 K2CO3 + Ca(OH)2 = CaC03 + 2KOH. 
 
 The solution of i part of potassium carbonate in 10-12 parts of water is boiled with i 
 part of slaked lime in an iron pot until a filtered portion does not effervesce, when hydro- 
 chloric acid is added ; i. e., until there is no longer any carbonic acid present. On 
 standing awhile, the insoluble calcium carbonate subsides, and the liquid becomes clear. 
 The solution of potassium hydroxide is then poured off, evaporated, the residue melted in 
 a silver dish (which it does not attack), until the hydroxide begins to volatilize in clouds, 
 when it is poured into moulds [K. causticuni ficsum). The caustic potash, prepared in 
 this way, is not entirely pure, but contains potassium chloride and other salts. To 
 obtain a product that is chemically pure, fuse potassium nitrate with copper filings, and 
 evaporate the aqueous extract of the fusion in silver vessels. 
 
 One of the most brilliant achievements of modern chemical industry is the electrolytic 
 decomposition of the chlorides of the alkali metals into free chlorine and 7 uetal ; the latter at 
 the moment of its liberation is converted by water into its hydroxide. The practical solu- 
 tion of this problem has made this operation a technical process since 1890. The Stass- 
 furt potassium chloride has been thus made to yield pure caustic in solution and in solid 
 form ; also chlorine, bleaching lime and hydrogen (pp. 51, 278 ; also under Soda). 
 
 Potassium hydroxide forms a white, crystalline mass which fuses rather 
 easily, and volatilizes undecomposed at a very high temperature. Ex- 
 posed to the air it deliquesces, as it absorbs water and carbon dioxide. 
 It is very soluble in alcohol, and esjiecially in water. The solution 
 {Liquor kalii caustici) possesses a strong alkaline reaction, saponifies the 
 fats, and has a corrosive action upon the skin and organic tissues. At 
 low temperatures the hydrate KOH -f- 2H2O crystallizes out from con- 
 centrated solutions. 
 
 The haloid salts of potassium are obtained by the direct union of the 
 halogens with potassium, and by the saturation of the hydroxide or car- 
 
276 
 
 INORGANIC CHEMISTRY. 
 
 bonatc witli lialoid acids, d'hey arc readily solulile in water, have a 
 salty taste, and crystallize in cul)es. When heated tliey melt, and are 
 somewhat volatile. 
 
 Potassium Chloride (A". chloratuj?i), K.C 1 , occurs in Stassfurt in 
 large deposits, as sylvite, and c(;mbined with magnesium chloride exists 
 di?, carnaili^e {iAgC\^, KCl -f 6H,/)). 
 
 The opening up and the practical yields of the Stassfurt salt beds containing carnal- 
 lite ; kieserite, MgS( )^ . 1 1 .^( ) ; tachhydrite, CaCl.^ . 2MgCl2 + 12II./); kainite, MgSO^.- 
 K.^SO^. MgCl.^ hll.^O; boracite, 2Mg3H^r)|. j b? ^ii^l ^ 1 ''^’ brotnides (p. 53) have 
 been of great coinniercial importance to the (ierinan 1 ‘hnpire. 'These arc the largest salt- 
 beds in the world. before they were worked very considerable (|uantities of salt were 
 imported into Germany but now that country stands at the head of all the salt producing 
 portions of the earth. In 1897 .Stassfurt yielded 273,364 tons of rock-salt ($277,026) 
 and 1,946,188 tons of crude jwtash salts ($6,513,553) ; in the same year the production 
 of potassium chloride in the German Empire amounted to 108,000 tons ($5,764,423). 
 
 Carnallite serves as the chief source for the pre|)aration of potassium 
 chloride; water decomposes it into the more sparingly soluble jxjtassium 
 chloride and the readily soluble magnesium chloride. It is interesting to 
 note that three-fourths of the ])Otassium* chloride sejtarate in solid form 
 when carnallite is heated to 176°. The licpiid separated from this yields 
 carnallite again on cooling to 115°, while magnesium chloride remains 
 dissolved. The chloride crystallizes in vitreous cubes, of sj^ecific gravity 
 1.98. It melts at about 800°, and volatilizes at a strong red heat [see 
 Jahrbuch der Chemie v (1895), 66, 67]. 100 parts of water dissolve 29 
 
 parts of the salt at 0°, and 56 parts at 100°. Potassium chloride is used 
 in making nearly all the other potassium salts, hence it is largely ap- 
 plied technically (see Potassium Carbonate, Potassium Nitrate, Potassium 
 Chlorate). 
 
 Potassium Bromide (^K. bromatuni), KBr, is generally obtained 
 by warming a solution of potassium hydroxide with bromine, when the 
 bromate is also produced : 
 
 6KOH 3Br.2 = 5KBr -|- KBrOg y 3H2O. 
 
 The solution is evaporated to dryness, mixed with charcoal, and ignited, 
 which reduces the bromate to bromide : 
 
 KBrOg -)- 3C = KBr -|- 3CO. 
 
 It is readily soluble in water and slightly in alcohol ; forms cubes of spe- 
 cific gravity 2.4, and melts at about 740°. 
 
 Potassium Iodide (A', iodatuni), KI, may be prepared like the 
 })receding. The iodate produced along with the iodide upon adding 
 iodine to caustic potash may be reduced by hydrogen peroxide : 
 
 KlOg + 3IIA = d 3II2O + 3O, (p. 180). 
 
 It is usually obtained according to the following method : Iodine (3 
 parts) and iron filings (i part) are rubbed together under water ; an equal 
 (|uantity of iodine (i part) is again added to the solution of this ferrous 
 iodide, b'el^, in order to form ferrous-ferric iodide, Fe.Jg, which is then 
 boiled with the recpiired (juantity of pota.ssium carbonate; this will pre- 
 
POTASSIUM CHLORATE. 
 
 277 
 
 cipitate ferrous-ferric oxide ; carbon dioxide escapes, and potassium 
 iodide will be found in the solution ; 
 
 Fcglg -f 4K2CO3 = FegO^ + SKI + 4CO2. 
 
 It forms large white crystals, fuses at about 720°, and is tolerably volatile. 
 Its specific gravity equals 3.0. At medium temperatures it dissolves in 
 0.7 part of water and 2.5 parts of absolute alcohol. The aqueous solution 
 dissolves iodine in large quantity. Many metallic, insoluble iodides dis- 
 solve in it without difficulty, forming double iodides, e. g., Hgl2.2KI. 
 The iodide is employed in medicine and in photography. 
 
 Potassium Fluoride, KFl, is obtained by dissolving the carbonate in aqu^us hydro- 
 fluoric acid. It crystallizes in cubes at ordinary temperatures, with two molecules of 
 water, but above 35° does not contain water of crystallization. It is very .soluble in 
 water. The aqueous solution attacks glass. It is greatly inclined to combine with other 
 fluorides: KFl.HFl; BFI3.KFI. The commercial salt is frequently rich in arsenic. 
 On adding hydrofluosilicic acid to the solution of potassium .salts, a gelatinous precipitate 
 of potassium silico-fluoride, K2SiFlg, is thrown down, wdiich dissolves with difficulty in 
 water. 
 
 Potassium Cyanide, KCN. This salt can be produced by saturating potassium 
 hydroxide with hydrocyanic acid, and by heating yellow prussiate of potash (see Iron). 
 It forms a white, easily fu.sible mass, which deliquesces in the air. The solution may be 
 easily decomposed. The salt crystallizes in cubes, has an alkaline reaction, and smells 
 like prussic acid. As the result of hydrolysis free prussic acid — which is very slightly 
 dissociated — is pre.sent with free alkali. The introduction of carbon dioxide completes the 
 decomposition. By fusion potassium cyanide reduces many metallic oxides, and hence is 
 employed in reduction proces.ses. It is just as poisonous as prussic acid. It is applied 
 in many ways, especially in photography and for galvanic silvering and gilding. Lately 
 it has met with extended application in the extraction of gold from low-grade ores and 
 from sand. Generally it is a mixture of potassium and sodium cyanides which is used for 
 this purpose ; this can be obtained by fusing potassium ferrocyanide with sodium : 
 
 K,Fe(CN)6 -f 2Na = 2NaCN + 4KCN -f Fe. 
 
 Potassium Chlorate (A", chloricuni) KCIO3, is produced when a 
 slight excess of chlorine is introduced into caustic potash, and the hypo- 
 chlorite, formed at first, oxidized thereby to chlorate. This change pro- 
 ceeds most rapidly at 80-90° (see p. 178). The sparingly soluble chlorate 
 separates when the solution cools. 
 
 Technically, a solution of calcium chlorate, produced by slightly super- 
 saturating lime-water with chlorine at 40°, is mixed with a sufficient quan- 
 tity of ])otassium chloride, when potassium chlorate and calcium chloride 
 result : 
 
 Ca(C103)2 + 2KCI = CaCb + 2KCIO3. 
 
 Magnesia can be advantageously substituted for lime-water. 
 
 At present this old method of Liebig is being more and more supplanted 
 by the electrolytic method, particularly since Oettel found that potassium 
 chlorate is formed electrolytically in alkaline solutions of ])otassium chlo- 
 ride without the use of a diaphragm, which se])arated the anode from the 
 kathode liquor. Since the electrolysis of potassium chloride yields caustic 
 (together with hydrogen) at the kathode and chlorine at the anode, the 
 conditions for the production of chlorate are evident. See also Elbs, 
 Chem. Zeit. 1897, 996. 
 
INORGANIC CHEMISTRY. 
 
 278 
 
 A French company in Switzerland first made cldorate electrolytically, while the caustic 
 alkalies were first produced in this way in (Jermany; England followed later. 'I’he ledi 
 nical manufacture of electrolytic chlorate has been in operation since 1891 in Switzerland. 
 The chemical factory at (Iriesheiin began to make caustic potash and soda electrolytically 
 in 1890, Since 1894 the electric factory at Hitterfeld has jmxluced chiefly caustic and 
 bleaching li(juors in the electrolytic way. Other companies have followed those men- 
 tioned, and by some of them soda is now being prepared from electrolytic caustic soda. 
 
 Potassium chlorate crystallizes from the hot solution in shinitig tables 
 of the monoclinic system, which dissolve with difficulty in cold water(Too 
 parts at the ordinary temperature dissolve 6 ])arts of the salt). Its taste 
 is cooling and astringent. When heated it melts at 360° and at higher 
 temperatures gives up a ])ortion of its oxygen, and changes to the chloride 
 and Perchlorate, KCK)^, which on further heating decomposes into 
 oxygen and potassium chloride (see \). 178). With hydrochloric acid it 
 liberates chlorine : 
 
 KCIO3 + 6 IIC 1 = KCl + 3H2O + 3CI2. 
 
 Mixed with sulphur, or certain sulphides, it explodes on heating and 
 when struck a sharp blow. The igniting material upon the so-called 
 Swedish (parlor) matches consists of antimony sulphide and potassium 
 chlorate; when this is rubbed upon the friction surface coated with red 
 phosphorus it ignites. 
 
 Potassium Hypochlorite, KCIO, is formed when chlorine is con- 
 ducted into a solution of potassium hydroxide : 
 
 2KOH + CI2 = KCl + KCIO -f H2O. 
 
 It only exists in aqueous solution; when the latter is evaporated the 
 salt is decomposed into chloride and chlorate : 
 
 3KCIO = 2KCI 4- KCIO3. 
 
 In the presence of an excess of chlorine chlorate is rapidly produced. 
 
 The solution has an odor resembling that of chlorine, and bleaches 
 strongly, especially upon the addition of acids (p. 175). The bleaching 
 solutions occurring in trade (Eau de Javelle) are prepared by the action of 
 chlorine upon solutions of sodium (Eau de Labarraque) and potassium 
 (Eau de Javelle) carbonates; but recently they have been made by the 
 electrolysis of the corresponding chlorides. They also contain free' hypo- 
 chlorous acid. 
 
 The oxy-salts of bromine and iodine are perfectly analogous to those of chlorine. 
 Potassium hromnte, KBrO,, and Potassium iodate, KIO3, are ])repared by the action of 
 bromine or iodine u])on hot potassium hydroxide ; the second is also produced by the 
 action of iodine upon ])otassium chlorate, when the chlorine is directly replaced (p. iSol. 
 'The exj)erimcnts of Klinger and of Bassett apparently prove that a direct replacement of 
 chlorine by iodine does not take place here ; it is rather the oxidation of the latter by the 
 anion (.'Kqof the chlorate. I^)tassiuln iodate can also be made by oxidizing potassium 
 iodide (i |)art) with ])otassium ])ennanganate (2 parts) in aqueous solution. If chlorine 
 be passed through a hot solution of jiotassium iodate or iodide in jK)tassium hydroxide, 
 the p(;rioda((‘ of potassium, KIO,, arises; it is difficultly soluble and when heated de- 
 composes into oxygen and potassium iodate, which then breaks down into potassium 
 iodide and oxygen. 
 
POTASSIUM NITRATE. 
 
 279 
 
 Besides the normal periodates, KIO^, NalO^, other salts exist which are derived from 
 the highest hydroxyl compound, and its anhydro- derivatives (p. 18 1). These 
 
 salts are very numerous, and are in part monoperiodates, 10(011)5 and 102(011)3, and 
 partly polyperiodates, produced by the condensation of several molecules of the highest 
 hydroxides with the exit of water, e. g.^ l203(0H)g, 1205(011 )4, and l20g(0Il)2. 
 
 Potassium Sulphate, K2SO4, is formed in the action of concentrated 
 sulphuric acid upon potassium chloride, and as a by-product in many 
 technical operations. It is also obtained by transposing schonite, 
 MgSO^. K2SO4 -f 6H2O, and other Stassfurt salts containing sulphates 
 with potassium chloride: 
 
 MgSO,.K2SO, -f 2KCI = 2K2SO, -f MgCb- 
 
 It crystallizes without water, in small rhombic prisms, having a bitter, 
 salty taste, and dissolves in 10 parts of water at the ordinary temperature. 
 It melts at about 1070°. It is employed principally for the preparation 
 of potassium carbonate, according to the method of Le Blanc (p. 281). 
 
 The acid or primary salt, KHSO^, crystallizes in large rhombic 
 tables, and is very readily soluble in water. It fuses at about 200°, loses 
 water, and is converted into potassium pyrosulphate, K2S20^, which at 
 600° yields K2SO4 and SO3 (p. 194). 
 
 Potassium Sulphite. — The salts of sulphurous acid— the primary^ KHSO3, and 
 the secondary sulphites, K2SO3 — are produced when sulphuric dioxide comes in contact 
 with a potassium carbonate solution ; they are very soluble and crystallize with difficulty. 
 The first salt shows an acid, the second an alkaline reaction. If sulphur dioxide be 
 passed into a solution of potassium carbonate until effervescence ceases and then cooled, 
 the pyrosulphite, K2S.2O5, corresponding to the pyrosulphate, will crystallize out. 
 
 Potassium Persulphate, K2S20g, results on electrolyzing a saturated solution of 
 acid potassium sulphate, when it separates at the anode as a white crystalline precipitate. 
 It can be crystallized from hot water ; on rapid cooling it separates in small prisms. Its 
 solution oxidizes, has a cooling, salt-like taste and does not yield a precipitate with solu- 
 tions of other metals (the salts of silver, manganese and cobalt excepted). The dry salt 
 commences to decompose at 100° into oxygen and pyrosulphate (see p. 188). 
 
 Potassium Nitrate, Saltpeter {K. nitrictmi), KNO3, occurs in the 
 largest amounts in East India, especially in Ceylon, at Bengal and at 
 Gutscharat (Bombay) in distinct layers; in cavities (in Ceylon) once 
 the lairs of animals and of men’s dwelling-places, which are even now 
 inhabited by hosts of field mice. It is found in abundance in other 
 torrid regions in the soil, upon which it effloresces during the dry sea- 
 son (hence sal, salt; iz^rpa, rock), e. g., in Peru, in Bolivia, in South 
 Africa, also in Egypt, etc. It is produced wherever nitrogenous organic 
 substances decay in the presence of potassium carbonate (aided by micro- 
 organisms), conditions which are present in almost every soil. The 
 intentional introduction of these is the basis of the artificial niter pro- 
 duction in the so-called saltpeter plantations, which were formerly cul- 
 tivated actively in Spain, Hungary, Sweden and Switzerland. Manures 
 and various animal offals are mixed with wood ashes (i)otassium car- 
 bonate) and lime, arranged in porous layers, and submitted to the action 
 of the air (protected from rain) for two or three years, when nitrates are 
 produced from the slow oxidation of the nitrogen. The heaps are then 
 
28 o 
 
 INORGANIC CUKMISTRY. 
 
 treated with water and potassium carbonate added to the solution, which 
 contains potassium, calcium and magnesium nitrates, to convert the last 
 two salts into potassium nitrate : 
 
 Ca(N03)2 + 1^2^03 = CaCOa + 2KNO3. 
 
 The precipitate of calcium and magnesium carbonates is filtered off and 
 the solution evaporated. 'Fhe soils of J^^ast India containing niter are 
 similarly worked. Until the Crimean War (1852-1855) the demand 
 for poiassium nitrate in manufacturing gunjiowder was met almost exclu- 
 sively by h^ast India. The numerous and constant inquiries for the salt 
 led (jerman chemists to transiiose the Chili saljieter by means of Stass- 
 furt potassium chloride into the j)>;tassium salt (conversion saltpeter): 
 
 NaN()3 + KCl = NaCl | KNf)3. 
 
 Warm saturated solutions of e(iui valent quantities of sodium nitrate 
 and potassium chloride are mixed and boiled, when sodium chloride, 
 being less soluble in hot water, will sejiarate. On cooling the solution 
 potassium nitrate, being less soluble in cold water, crystallizes out; 
 sodium chloride is about equally soluble in hot and cold water, for which 
 reason the portion not separated by boiling remains in solution (i). 266). 
 
 Potassium saltpeter crystallizes without water of crystallization in large, 
 six-sided rhombic prisms. It is far more soluble in hot than in cold 
 water; 100 parts of w'ater dissolve 247 parts at 100°, but at 0° only 13 
 ])arts. It possesses a cooling taste, fuses at 340°, and decomposes, when 
 further heated, into oxygen and potassium nitrite, KNO2. Heated with 
 carbon it yields potassium carbonate : 
 
 4KNO3 -|- 5C = 2K2CO3 -(- 3CO2 T 2N2. 
 
 Its'principal use is in the manufacture of gunpowder. This is a granular mixture of 
 potassium nitrate, sulphur, and charcoal. The relative quantities of these constituents 
 are somewhat different in the various kinds of powder (sporting, blasting, and powder 
 free from sulphur). The first consists of 4KNO3 -|- 2C -f- 2S = 2 K^S 0 ^ -[- 2CO2 + 2N2 ; 
 the second : 4KNO3 6C -)- 4S — 2K2S2 -f- 6CO2 -( 2N2» ^he third : 4KNO3 -j- 5C 
 = 2K2CO3 -f- 3Ct)2 + 2N2. These three varieties, mixed in different but simple propor- 
 tions, constitute the powders in general use. Each variety possesses peculiarities as to 
 ignition, combustibility, energy content, heat and gas content. The mixing is made in 
 accordance with the demand for any one or more of these properties. The effectiveness 
 of the powder, therefore, depends upon the disengagement of carbon dioxide and nitrogen 
 gas, the volume of which is almost 1000 times as great as that of the decomposed powder. 
 
 Potassium Nitrite (A". nUrosuvi), KNO2, is obtained by fusing 
 saltpeter with lead (2 parts) which withdraws one atom of oxygen from 
 the former (p. 205). It forms a white, crystalline, fusible mass, which 
 deliquesces in the air. 
 
 Potassium Phosphates. — The potassium salts of phosphoric acid : 
 K^PO^, K2lIP()^, and KH2P(),j, meet with no ju'actical ai)plication, they 
 are readily soluble in water and crystallize ])oorly ; therefore, the sodium 
 salts are generally used. The borates, KIIO2 and K2B^O. -|- 5H2O (see 
 borax), crystallize with difficulty. 
 
 Potassium Carbonate, K2(X).,, ordinarily known as potashes, is 
 the principal ingredient of the ashes of land plants. 
 
POTASSIUM CARBONATE. 
 
 281 
 
 As 1000 parts of wood yield 3.5-28 parts of ashes and the potassium 
 carbonate in the latter is 0.45-4 parts it is obvious that only countries 
 like Russia, Canada, the United States, Hungary and Galicia where 
 there is an abundance of woodland can produce potashes in large 
 amounts from the wood ashes. The latter are extracted with water, the 
 clear filtrate is then evaporated until it solidifies on cooling, when the 
 residue is dehydrated and decolorized by calcination in ovens. The prod- 
 uct is crude potashes. By repeating the preceding treatment purified 
 potashes or pearl ash is obtained. In countries like Belgium, Germany, 
 France and Switzerland where sugar beets are cultivated large quantities 
 of potassium carbonate are separated from the ash of the beets. Beets 
 withdraw large quantities of potash salts from tl.e soil which if the latter 
 is to remain fertile must be returned to it by potash fertilizers. In this 
 direction the Stassfurt potash salts are most valuable. The alkali salts of 
 the beet are in its juice and they remain, when the sugar is extracted 
 from the latter, in the molasses. When the latter is allowed to ferment, 
 then evaporated and subjected to dry distillation, alcohol, ammonia, 
 trimethylamine and other valuable substances are obtained ; furthermore, 
 the residual coke is rich in potash. Its aqueous extract is worked for 
 potassium carbonate. Sheep’s suint contains potassium salts of organic 
 acids which on incineration yield potassium carbonate. The same salt 
 is obtained as a by-product in the manufacture of iodine and bromine 
 from sea-algse and sea-weeds. 
 
 In Germany all these methods for the preparation of potassium carbon- 
 ate give place to those in which the Stassfurt potassium chloride is util- 
 ized, from which potassium carbonate is prepared by two methods : 
 
 1. Method of Le Blanc. — This will be discussed in connection with 
 soda. H. Griineberg was the first to employ it in making potashes. 
 
 2. Method of Ch. R. Engel in Montpellier. — Brecht introduced this 
 method into German industries. Magnesium carbonate is mixed with a 
 solution of potassium chloride, and the liquid, while being stirred, is 
 saturated with carbon dioxide. A double salt of magnesium carbonate 
 and acid potassium carbonate separates while the solution contains mag- 
 nesium chloride : 
 
 SMgCOj + 2KCI + CO2 + 9H2O = 2[MgC03 . KHCO3. 4H2O] + MgCh- 
 
 The double salt is freed from any adherent liquor by washing it with a 
 solution of magnesium bicarbonate, after which it is decomposed under 
 pressure with water at 140°. Basic magnesium carbonate separates in a 
 compact form ; carbon dioxide escapes, and the liquid containing pure 
 potashes is evai)orated and the residue calcined. Potassium carbonate 
 free from sodium salts can be obtained in this way, because sodium chlo- 
 ride does not act upon magnesium carbonate. 
 
 The commercial carbonate is a white, granular, deliquescent powder, 
 melting at 890°, and vaporizing at a red heat. It crystallizes from con- 
 centrated aqueous solutions with molecules of water, in monoclinic 
 prisms; at 100° it loses ^ molecule of water. The solution has a caustic 
 taste and shows an alkaline reaction. When carbon dioxide is con- 
 24 
 
282 
 
 INORGANIC CHEMISTRY. 
 
 ducted through the liquid it is absorl)cd and primary potassium carbonate 
 is produced : 
 
 KjCO, -f 1 1 /) I CO, = 2KIICO,. 
 
 This salt, ordinarily called bi-carbonate, crystallizes in inonoclinic 
 l)risms, free from water. It dissolves in 3-4 ])arts of water and exhibits a 
 neutral reaction. When heated, it decomi)oses into i)otassiuin carbonate, 
 carbon dioxide, and water, 'bhe decornj)osition of the dry salt does not 
 begin until at about 110°, while the aqueous solution decomposes even on 
 evaporation. Potassium carbonate is used almost entirely in the produc- 
 tion of Bohemian or crystal glass. 
 
 Constam and v. Hansen electrolyzed solutions of potassium carbonate, cooled to 
 — 15°, and obtained at the anode Potassium Percarbonate, K2C20g, — a slightly blue- 
 colored, deliquescent powder. Jt is formed, like the persulpliates, by the union of the 
 ions KCO.J, into which (together with potassium ions) the carbonate is decomposed by 
 the current. It resembles the persuljdiates in properties, in so far that when heated to 
 200-300° it rapidly decomposes into carbonate and oxygen. Oxygen escapes from its 
 aqueous solution at 45°. It is a powerful oxidant ; many dyes are bleached by it. 
 Dilute acids evolve hydrogen peroxide from it: K2C2f)fi 4 2iICl 2CO2 + 2KCI -|- 
 H2O2 ; this also occurs with caustic potash : K2C20g | 2KOII = 2K2CO^ -j- H/),. Its 
 solution rapidly reduces manganese peroxide, lead peroxide and silver oxide with energetic 
 liberation of oxygen : Ag20 f K2C20g = Ag, -f- K2CO3 4 - CO, -(- O,. See p. 102. Its 
 chemical structure is very likely analogous to that of the persulphates (pp. i88, 189, 279) : 
 
 O — SO3K O — CO3K 
 
 6 — SO3K 6 — CO3K. 
 
 Potassium Silicate, water-glass, does not possess a constant com- 
 position and cannot be obtained crystallized. It is produced by solution 
 of silicic acid or amorphous silicon dioxide in potassium hydroxide, or by 
 the fusion of silica with potassium hydroxide or carbonate. The concen- 
 trated solution dries in the air to a glassy, afterward opaque, mass, 
 which, when reduced to a powder, will dissolve in boiling water. Potas- 
 sium (and also sodium) water-glass has an extended application, especially 
 in cotton printing, for the fixing of colors (stereochromy), in rendering 
 combustible material fireproof, in soap boiling, etc. 
 
 SULPHUR COMPOUNDS OF POTASSIUM. 
 
 Potassium Hydrosulphide, KSH, is obtained when potassium 
 hydroxide is saturated with hydrogen sulphide : 
 
 KOH + II,S = KSH -b H2O. 
 
 Evaporated in vacuo it crystallizes in colorless rhombohedra, of the for- 
 mula 2KSri -j- H, 0 , which deliquesce in the air. At 200°, it loses its 
 water of crystallization, and at a higher temperature fuses to a yellowish 
 li(]uid, which solidifies to a reddish mass. Like the hydroxide, it has an 
 alkaline reaction. On adding an ecpiivalent quantity of potassium hydrox- 
 ide to the sulphydrate solution, we get potassium sulphide: 
 
 KSH + KOH = K,S -f H,, 0 . 
 
SULPHUR COMPOUNDS OF POTASSIUM. 
 
 283 
 
 Potassium Sulphide, KgS, is usually obtained as a porous mass by 
 gently heating a mixture of potassium sulphate and carbon in well-closed 
 crucibles : 
 
 KjSO^ + 2 C = KjS + 2CO2. 
 
 When fused, it solidifies to a red crystalline mass. It crystallizes from 
 concentrated aqueous solutions with five molecules of water, in colorless 
 prisms, which deliquesce in the air. The solution absorbs oxygen from 
 the air, and is decomposed into potassium hyposulphite and caustic 
 potash : 
 
 2KjS -h HjO + 2O2 = K2S2O3 + 2KOH. 
 
 Potassium hydrosulphide and sulphide precipitate insoluble sulphides 
 from the solutions of many metallic salts. They are decomposed by acids 
 with liberation of hydrogen sulphide. 
 
 When the aqueous solution of the sulphide is boiled with sulphur the 
 poly sulphides, K2S3, K^S^, and K2S-, are formed. The aqueous solutions 
 of the polysulphides are decomposed by acids, with disengagement of 
 hydrogen sulphide and separation of sulphur (milk of sulphur). The so- 
 called liver of sulphur (^Hepar sulphiiris, K. sulphuratumf a liver-brown 
 mass, used in medicine, is obtained by the fusion of potassium carbonate 
 with sulphur, and consists of a mixture of potassium trisulphide with 
 potassium sulphate and hyposulphite. 
 
 The aqueous .solution of the potassium, as well as that of the sodium 
 sulphide, dissolves some metallic sulphides, and forms sulpho-salts with 
 them (pp. 223, 225). 
 
 When dry ammonia is conducted over heated potassium, potassamide 
 (NHjK) results. This is a dark-blue liquid which solidifies to a white, 
 crystalline mass. It sublimes at about 400°, and above that temperature 
 breaks down into its elements. Water decomposes it into potassium 
 hydroxide and ammonia. 
 
 Recognition of the Potassium Compounds. — In all of its com- 
 pounds potassium is present as a positive univalent element. Almost all of 
 the potassium compounds are easily soluble in water. The few exceptions 
 serve for the characterization and separation of potassium. Tartaric acid 
 added to the solution of a potassium salt gives a crystalline precipitate 
 of acid potassium tartrate. Platinic chloride (PtClJ produces in potas- 
 sium solutions a yellow, crystalline precipitate of PtCl^. 2KCI (p. 271). 
 Potassium silicofluoride,K2vSiFlp, is also sparingly soluble and can be used 
 in detecting and estimating potassium. Potassium compounds introduced 
 into the flame of an alcohol or a gas lamp impart to the same a violet col- 
 oration. The spectrum of the flame is characterized by two bright lines, 
 one red and one violet (see Spectrum Analysis). 
 
284 
 
 INORGANIC CHEMISTRY. 
 
 2. RUBIDIUM AND Ci^:SIUM. 
 
 Rl) 85.4. Cs 133. 
 
 Rubidium and caesium are tlie ])erfcct analo^rues of potassium (p. 273). d'hey were 
 discovered by means of the spectroscope, by lUmsen and Kircldioff, in i860. Although 
 only occurring in small quantities, they are yet very widely distributed, and fre(juently 
 accompany potassium in mineral sj^rings, salt, and plant ashes. 'I'he mineral lepidolite 
 contains 0.5 percent, of rubidium; upward of 30 i)er cent, of caesium oxide is present 
 in the very rare pollucite, a silicate of aluminium and cnesium. Stassfurt carnallite also 
 contains rubiilium. The spectrum of rubidium is marked by two red and two violet 
 lines; caesium by two distinct blue lines; hence the names of the.se elements {riibiilns, 
 dark red ; itcsiiis, sky-blue). 
 
 Rubidium and caesium form double chlorides (PtCl^. 2RbCl, PtCl^.CsCl) with plati- 
 num chloride, and they are more insoluble than the double platinum salt of potassium, 
 hence may answer for the separation of these elements from potassium, d'his is also 
 true of the comi)ounds 2RbCl.SnCl^, 2CsCl.SnCl^ and 2SbCl3. 3CSCI, which are rather 
 sparingly .soluble. 
 
 Rubidium and caesium may be obtained free by decomposing the fused chloride with the 
 electric current. Erdmann and Kothner [Ann. Ch. 294 (1897), 58] obtained large yields of 
 rubidium by heating its hydroxide with magnesium in a seamless, knee-shaped iron tube : 
 
 2RbOII -(- 2jMg = 2Rb 2MgO II.^. 
 
 Dry hydrogen is conducted through the tube while heating. The rubidium, which dis- 
 tils over, is collected under liquid paraffin. Metallic rubidium has a silver-white color, 
 with a .somewhat yellowish tinge ; its vapor is greeni.sh-blue. Its .specific gravity equals 
 1.52; its melting point is 38.5°. Oxygen at the ordinary temperature converts it into 
 rubidium dioxide, Rb02, consisting of dark-brown crystals, which are transposed by 
 water at a gentle heat into oxygen and rubidium hydroxide : 
 
 2 Rb 02 + 2H2 = 2RbOII -f H2O + O. 
 
 The dioxide dissolves in water with hissing and the tumultuous evolution of oxygen ; 
 hydrogen peroxide is produced at the same time. The oxide RbO is not known. The 
 iodide at present is quite frequently substituted in medicine for potassium iodide. 
 
 Ccesium like rubidium is isolated by heating coesium hydroxide with magnesium powder 
 in an atmosphere of hydrogen or by the electrolysis of a mixture of caesium and barium 
 cyanides. Electrodes of aluminium are employed for this purpose. Caesium is a silver- 
 white metal, of specific gravity 1.85. It oxidizes quite readily and inflames in the air. 
 It melts at 26.5° and boils at 270°. 
 
 Since 1892 Wells and Wheeler, and also Erdmann, have prepared an interesting 
 series of caesium and rubidium halides. The metals in these appear to be trivalent and 
 also quinquivalent, e. g., RbClBr.^, RbBi^, RbCl^I, CsBrg, CsBr3l2, Cslj. Rubidium 
 iodine tetrachloride, RblCl^, results in conducting chlorine into a rubidium iodide solu- 
 tion. It crystallizes in monoclinic yellow leaflets. It dissolves with difficulty in water. 
 Its solution acts as a powerful oxidant ; it dissolves gold and platinum. It is still undeter- 
 mined whether these are atomic or only double compounds {e.g., RbCl.IClg). The 
 rubidium and caesium compounds are distingui.shed from those of potassium by entering 
 into union more readily and in greater proportion with other halides to form double salts, 
 e. g., with A5CI3, AsBr3 and ASI3, with which potassium double halides have not been 
 prepared. 
 
 3. SODIUM. 
 
 Na = 23.05. 
 
 Sodium is widely distributed in nature, especially as chloride in sea- 
 water and as rock-salt; and isalso found in silicates. Its nitrate is Cliili 
 saltjteter, and its fluoride in union with aluminium fluoride constitutes 
 
SODIUM. 
 
 285 
 
 the cryolite of Greenland. The metal was obtained in 1807, by Davy, 
 by the action of a strong electric current upon fused sodium hydroxide. 
 It was formerly (1855), like potassium, obtained upon a large scale by 
 igniting a mixture of sodium carbonate, finely divided anthracite and 
 limestone in an iron retort : 
 
 Na^COj -{-20 = 2Na -[- 3CO. 
 
 A great advance was made in its mansfacture when Castner (1866), at Oldbury, Bir- 
 mingham, reduced the hydrate instead of the carbonate with carbon impregnated with 
 finely divided, spongy iron. Gay-Lussac and Thenard (1808) had reduced the hydroxide 
 with metallic iron at a white heat : 
 
 2NaOH -j- 2Fe = 2Na -j- 2FeO -)- H2. 
 
 Netto (1898, Wallsend, Newcastle-on-Tyne) invented a most satisfactory process, 
 abandoning it, however, later. It consisted in allowing molten caustic soda to run over 
 wood charcoal placed in vertical iron retorts. Sodium vapors escaped constantly from 
 an exit tube at the top, while fused carbonate ran out from a tube at the bottom : 
 
 3NaOH + C Na2C03 + Na 
 
 The sodium vapors were condensed in flat iron receivers and the liquid metal was 
 collected under rock oil. 
 
 Castner and Kellner electrolyzed a salt solution, using mercury as anode ; the amalgam 
 was then distilled when mercury was expelled and sodium remained. 
 
 Grabau electrolyzed fused .salt after reducing its melting point by adding potassium 
 and strontium chlorides. The sodium then escaped in vapor form. 
 
 Sodium can be made on a small scale by heating the peroxide with freshly ignited 
 wood charcoal : 
 
 3Na20.2 -f 2C = Na.2 + 2Na2C03. 
 
 Calcium carbide may be substituted for the charcoal : 
 
 7Na202 + 2CaC2 = 2CaO -f 4Na2C03 -|- 3Na2. 
 
 The action in both cases is very energetic. 
 
 Sodium in external properties is very similar to potassium. It melts at 
 95.6°, boils at 742°, and is converted into a colorless vapor, which 
 burns with a bright yellow flame in the air. It oxidizes readily on ex- 
 posure, and decomposes water even in the cold, although less energet- 
 ically than potassium. A piece of sodium thrown upon water swims 
 about upon the surface with a rotary movement, the disengaged hydrogen, 
 however, not igniting. If we prevent the motion, by confining the 
 metal to one place, the heat liberated by the reaction attains the ignition 
 temperature of hydrogen, and a flame follows, as was also observed with 
 p jtassium (p. 40). 
 
 Sodium Oxide, NajO, is not definitely known. 
 
 Sodium Peroxide, Na202, has recently been introduced into commerce as a 
 bleaching agent. It is made by heating sodium in a stream of dry air, using vessels of 
 aluminium and a temj)erature below 300°. It is a yellow-white ]X)wder. It melts with 
 greater difficulty than caustic .soda, and at elevated temperatures gives off oxygen. Water 
 decomposes it with caustic soda and oxygen which escapes uj)()n boiling. With ice-water 
 it forms a solution containing both caustic .soda and hydrogen j)eroxide. On careful evapora- 
 tion such solutions yield crystalline hydrates of .sodium superoxide. Anhydrous acids and 
 alcohol appear to decompose it with the formation of a peculiar hydrate : 
 
 NajO, i IlCl = NaCl + Na02H, 
 
286 
 
 INORGANIC CHEMISTRY. 
 
 wliich probably lias the formula Na-()-0-H. Sodium superoxide acts on many organic 
 compounds with the production of flame and the separation of carbon. At a red heat it 
 is superior to all other oxidants in its powerful action (i jiart of the material, 2 parts of 
 soda and 4 parts of .sodium superoxide). 
 
 Sodium Hydroxide, sodium hydrate, or caustic soda, NaOH, like 
 ])Otassium hydroxide, is formed by boiling a solution of soditim car- 
 bonate with calcium hydroxide: 
 
 Na.^COg + Ca(OII)2 = CaCOj + 2NaOII. 
 
 At present it is directly protluced in the soda manufacture by adding a 
 little more carbon to the fusion (see Soda); or by the electrolysis of 
 sodium chloride (jtp. 290, 291). 
 
 'bhe sodium hydroxide which solidifies after fusion is a white, radiating, 
 crystalline mass, and re.sembles caustic ]iotash very much. It attracts 
 water from the air, becomes moist, and coats itself by carbon dioxide 
 absorption with a white layer of sodium carbonate (caustic potash del- 
 i(piesces perfectly, because the resulting carbonate is al.so deliquescent), 
 d'he aqueous solution, called sodium hydroxide, resembles that of potas- 
 sium. Crystals of NaOH -j- separate at 0° from the concentrated 
 
 solution ; they melt at 6°. 
 
 Sodium Chloride, NaCl, is abundant in nature. It is found almost 
 everywhere in the earth and in natural waters; in sea-water it averages 
 2. 7-3. 2 per cent. As rock-salt it forms large deposits in many districts, 
 especially at Stassfurt and Wieliczka in Galicia (p. 276). 
 
 In warm climates, on the coasts of the Mediterranean Sea, sodium chloride is gotten 
 from the sea according to the following procedure : At high tide sea- water is allowed 
 to flow into wide, flat basins (salt gardens), in which it evaporates under the sun’s heat ; 
 the working is limited, therefore, to summer time. After sufficient concentration, pure 
 sodium chloride first separates, and this is collected by itself. Later, there crystallizes a 
 mixture of sodium chloride and magnesium sulphate ; finally potassium chloride, magne- 
 sium chloride and some other salts appear (among them potassium iodide and bromide), 
 the .separation of which constitutes a special industrial branch in some regions. In 
 cold climates, as in Norway and at the White Sea, the cold of winter is employed for the 
 production of salt. In the freezing of sea-water, as well as of other solutions, almost 
 pure ice separates at first ; the enriched sodium chloride solution is then concentrated in 
 the usual way. 
 
 Rock-salt is either mined in shafts, or, where the strata are not so large and are 
 admixed with other varieties of rock, alixiviation process is employed. Borings are made 
 in the earth and wmter runs into them, or into any openings already formed. When the 
 water has saturated itself with sodium chloride, it is pumped to the surface and the brine 
 then further worked. In many regions, especially in Reichenhall, in Bavaria, more or 
 less saturated natural salt or brine springs flow from the earth. The concentration of the 
 non-saturated brine occurs at first in the so-called “graduation” houses. These are long 
 wooden frames filled with fagots, and on letting the salt water run upon them it will be 
 distributed and evaporated by the fall ; the concentrated brine collects in the basin below, 
 and is then evaporated over a free fire. 
 
 Sodium chloride crystallizes from water in transparent cubes, of specific 
 gravity 2.13, which arrange themselves by slow cooling into hollow, four- 
 sided pyramids. It melts at 8 15° and volatilizes at a white heat. It is not 
 much more soluble in hot than in cold water ; 100 parts at 0° dissolve 36 
 l)arts of salt ; at 100”, 39 ])arts. The saturated solution, therefore, contains 
 
SODIUM SULPHATE. 
 
 287 
 
 about 26 per cent, of sodium chloride. If the saturated solution be cooled 
 below — 10, large monoclinic tables (NaCl -f- 2 H,^ 0 ) separate ; these lose 
 water at 0° and become cubes. 
 
 The ordinary sodium chloride usually contains a slight admixture of 
 magnesium salts, in consequence of which it gradually deliquesces in the 
 air ; the perfectly pure salt is not hygroscopic. When heated the crystals 
 crackle, because of the escape of the mechanically enclosed water. 
 
 Sodium Bromide and Iodide crystallize at ordinary temperatures with two molecules 
 of water, which they lose again at 30° ; above 30° they separate in anhydrous cubes. 
 Sodium bromide fuses at 760° and the iodide at 690° ; the former is difficultly soluble in 
 alcohol and the latter is very soluble. 
 
 Sodium Chlorate (NaC103) and Perchlorate (NaClO^) are considerably more solu- 
 ble in water than the corresponding potassium salts. 
 
 Sodium lodate, NalCXj, is obtained in the same manner as the potassium salt, and 
 at ordinary temperatures crystallizes with three molecules of water in silky needles. It 
 is pre.sent in Chili saltpeter. If chlorine gas be conducted through the warmed solution 
 
 of sodium iodate in sodium hydroxide, the periodate 10 <! P- crystal- 
 
 lizes out on cooling. This becomes the normal salt (NalO^ -|- 3H2O) when dissolved in 
 nitric acid. 
 
 Sodium Sulphate (A^atrmm sulphuriciiiJi), Na2SO^, crystallizes at 
 ordinary temperatures with ten molecules of water of crystallization, and 
 is then known as Glauber’s ’idXx. {Sal mirabile Glauberi'). It occurs in 
 many mineral waters, and in large deposits, with or without water of 
 crystallization, in Spain. It is a by-product in the manufacture of 
 sodium chloride from sea-water and brine. It is produced in large quan- 
 tities by heating salt with sulphuric acid : 
 
 2 NaCl + H2SO, = Na^SO, -f 2HCI, 
 
 and is used in making soda (sodium carbonate). Or it may be prepared 
 by the method of Hargreaves, by conducting sulphur dioxide, air and 
 steam over strongly ignited sodium chloride: 
 
 2 NaCl + SO, -f O T H ,0 Na-^SO, + 2HCI. 
 
 More recently the sulphate has been obtained by a transposition of 
 sodium chloride with magnesium sulphate at a winter temperature — a 
 procedure which is prosecuted chiefly in Stassfurt, where immense quan- 
 tities of magnesium sulphate (kieserite) exist : 
 
 2 NaCl + MgSO, = MgCl, + NaaSO,. 
 
 Sodium sulphate crystallizes at ordinary temperatures with ten molecules 
 of water, in large, colorless, monoclinic prisms, which crumble in the air 
 and fall into a white powder. The salt Na2SO^ -j- 10H2O is only stable at 
 temperatures below 34°. It melts at 34°, separating into a saturated 
 solution and the anhydrous salt Na2SO^. The latter melts at about 886°. 
 
 This explains the remarkable alteration which the .solubility of Glauber’s salt sustains 
 in water at 34°. As a rule the solubility of a substance in water does not change sud- 
 
288 
 
 TNOUGANIC CHRMISTRY. 
 
 flenly with the temperature, l)iit f,Mnflually. A siulch'ii cliaufre in .sohihilify is conuccled 
 witli a chanj^e in coTidition of the substance passinjr into soiutif)n. 'I’liis apj)lies also to 
 (dauber’s salt. lOO jiarts of water dissolve, at o°, 12 parts; at 18°, 48 parts ; at 25°, 
 too parts; at 30°, 200 parts; at 34°, 354 ])arts of the hydrous .salt. At the last tem- 
 perature the solubility is f^reatest ; by further increase of heat it ^rradually diminishes; at 
 50°, ICX) parts of water (lissolve only 263 parts ; at lOO°, 238 parts of the salt. 'I'his is 
 explained l)y the fact tliat up to 34° the solubility relates to the salt Na.^SO, ioII.^(). 
 whicli cannot exist at higher temperatures. From this point upwards the solubility 
 relates to the anhydrous salt wliich separates as a deposit in small rhombic pyramids, if 
 the solution of the hydrous salt, saturated at 34°, be healed higher. 
 
 When the .solution, saturated at 34°, is allowed to cool to the ordinary tem])erature, 
 and even lower, not the slightest sei)aration of crystals occurs, although the salt is vastly 
 more in.soluble at lower temperatures than at 34°. Many other .salts form snpersaluraled 
 solutions, although they are less striking than that of (dauber’s salt, d'he supersaturated 
 solution of the latter may be agitated and twirled about without cry.stallization setting in. 
 If, however, a glass rod, or some other .solid body, be introduced into the .solution, it will 
 solidify suddenly to a crystalline mass, d'he particles of dust floating about in the air 
 will have a like effect ; therefore, to preserve the supersaturated solution, the ve.ssel con- 
 taining it should be kept well corked. Ily accurately made investigations, it has been 
 determined that the crystallization of the supersaturated (dauber’s .salt solution is only 
 induced by contact with already formed crystals. These must then be pre.sent every- 
 where in the atmosphere, because only solids that have been expo.sed to the air, and 
 that have not been carefully cleansed afterward, bring about the crystallization. 
 
 In the crystallization of a supersaturated (dauber’s salt .solution considerable heat is 
 disengaged, and the mass increases in tem{)erature. This is becau.se the latent heat of all 
 substances in the liquid condition is greater than in the solid. At io°, occasionally, and 
 of their own accord, tran.sparent crystals, Na2SO^ 71^2^^ .separate from the supersatu- 
 rated solutions. These crystals change readily to anhydrous .sodium sulphate and 
 Glauber’s salt. 
 
 This salt is employed in medicine as a purgative, and finds extended 
 application in the manufacture of glass and in the preparation of soda. 
 
 The primary or acid sodium sulphate, NaHSO^, is obtained by 
 the action of sulphuric acid upon the neutral salt or upon sodium chloride : 
 
 NaCl + H2SO, = NaHSO, -f fICl. 
 
 At ordinary temperatures it crystallizes with one molecule of water, 
 and is perfectly analogous to the potassium salt. 
 
 The sodium salts of sulphurous acid are obtained by conducting sulphur dioxide into 
 solutions of sodium hydroxide or carbonate. The secondary sulphite, Na2S03, crystallizes 
 with seven molecules of water at ordinary temperatures ; in the presence of sodium 
 hydroxide, or by warming the solution, it separates in the anhydrous state. The prunary 
 sulphite, NallSOa, consisting of small, easily .soluble crystaks, gives up sulphur dioxide in 
 the air, and is oxidized to sodium .sulphate. According to Schwicker, when it is neutralized 
 with potassium carbonate it forms yellow, indistinct crystals of sodium potassium sul- 
 phite, NaKSO.j -(- II-iG. Primary potassium sulphite on the other hand, when neutral- 
 ized with .soda, yields the salt NaKSO^ -p 2H2O, consisting of hard, yellowish crystals. 
 'I'his forms entirely different compounds with organic iodides from those which the first 
 .salt (with which it is i.someric, if the water of cry.stallization be di.sregarded) affords. Two 
 isomeric .sodium potassium hyposulphites, NaKS2()3, are obtained when these salts take 
 up sulphur. 'I'he following formulas express their structure : 
 
 K_S( )2-( )Na 
 
 KS-SeVONa 
 
 Na-SC)2-OK 
 
 NaS-S02-OK. 
 
 [Her. 22 (1889), 1728; .see also pp. 186, 198.] 
 
SODIUIM CARBONATE. 
 
 289 
 
 Sodium Hyposulphite, Na2S203, Sodium Thiosulphate, is pre- 
 pared by boiling the aqueous solution of neutral sodium sulphite with 
 flowers of sulphur : 
 
 Na 2 S 03 -j- S = Na 2 S 203 . 
 
 It is obtained as a by-product in the recovery of sulphur from the soda 
 residues. It crystallizes with five molecules of water, in large monoclinic 
 prisms, dissolves very readily in water, and is somewhat deliquescent in 
 the air. At 56° it melts in its water of crystallization ; loses all water 
 at 100°, and decomposes by further heating into sodium sulphate and 
 sodium pentasulphide, Na2S.. When the dry salt is heated in the air, the 
 sulphur of the polysulphide burns with a blue flame. Acids decompose 
 the aqueous solution with separation of sulphur and evolution of sulphur 
 dioxide (p. 197) : 
 
 Na2S203 -f 2HCI = 2NaCl -f SO2 + S -f H2O. 
 
 Like the sulphate, it readily affords supersaturated solutions. The 
 hyposulphite is used as a reducing agent ; chlorine, bromine and iodine 
 are converted by it into the corresponding halogen salts. An iodine 
 solution is instantaneously decolorized by sodium hyposulphite with the 
 production of sodium tetrathionate (p. 199). Chlorine behaves differ- 
 ently; sulphuric acid and sodium chloride are produced. Upon this 
 reaction rests the application of sodium hyposulphite as an mitichlor in 
 chlorine bleaching, to remove the excess of the chlorine, which has a 
 destructive action upon the fiber. In consequence of its property of dis- 
 solving the halogen silver derivatives, it is employed in photography. 
 
 Sodium Carbonate (Soda), Na2C03. This, technically, very im- 
 portant salt is obtained (i) as natural soda, (2) from the ashes of plants, 
 and (3) by chemical methods. 
 
 1. Natural soda disintegrates from the soil in certain districts, as in 
 Hungary, Asia, and Africa, during the hot seasons. It is a constituent 
 of many mineral springs, e. g., Carlsbad, and of the soda seas of Egypt, 
 of Central Africa, and of the coasts of the Caspian and Black Seas, of 
 California, and of German East Africa. In these seas or lakes a salt 
 deposits having the formula 
 
 Na2C03 -h NallCO, + 2H2O. 
 
 In Egy])t it is called iro-tia and in Colombia urao. Its occurrence at 
 Owen’s Lake, Inyo County, California, is the only one of scientific 
 importance. The waters of this lake upon evaporation yield very large 
 quantities of quite pure soda. 
 
 2. Soda from the Ashes of Plants. — Before Le Blanc’s process came 
 into use the greater ])ortion of soda was made from the ashes of sea and 
 coast \)\diY\X.?> (^Chefiopodium, Salsola, Atriplex, Salicornia, etc.); these 
 assimilate the sodium salts of the earth, while the land plants absorb the 
 potassium salts, and for this reason contain potashes in their ash. In 
 Southern France and in Spain beach ])lants were and are yet specially 
 cultivated for this purpose. The Spanish soda was particularly rich in 
 
 25 
 
290 
 
 INORGANIC CHEMISTRY. 
 
 carbonate and controlled the markets for a long period. At present soda 
 from the ashes of plants possesses only a local interest, 
 
 3. In the chemical way soda is pre[)ared almost exclusively from salt 
 (NaCl). Tlie following methods are in use: (i) That discovered in 
 1 794 by Le Blanc; (2) the ammonia-soda i)rocess introduced in 1866, and 
 
 (3) the electrolytic method of recent date. To these may be added the 
 
 (4) cryolite soda process, 
 
 I. In the Le Blanc method the sodium chloride is converted into 
 sodium sulphate by warming with sulphuric acid (pj). 58, 287). When 
 the latter is dry, it is mixed with charcoal and chalk (calcium carbonate) 
 and ignited in a reverberatory furnace, d'wo principal ])hases may be 
 distinguished in this reaction. First, the carbon reduces the sodium 
 sulphate to sodium sulphide: 
 
 Na,SO, + 2C Na,S f- 2 CO,. 
 
 The sodium sulphide then acts upon the calcium carbonate to form 
 calcium sulphide and sodium carbonate: 
 
 Na.^S -f- CaCOj = CaS -(- Na^COg. 
 
 At the same time the high temperature converts a portion of the cal- 
 cium carbonate into calcium oxide and carbon dioxide, which is reduced 
 by the ignited carbon to the monoxide: 
 
 CaC 03 = CaO + CO, ; CO, + C = 2 CO. 
 
 The appearance of the latter, which burns with a bluish flame, indicates 
 the end of the action. The chief products in the soda fusion are, then, 
 sodium carbonate and calcium sulphide, mixed with varying amounts of 
 calcium carbonate, calcium oxide, and foreign substances. This fused 
 mass is called C7'iide soda. It is lixiviated in specially constructed appa- 
 ratus with cold water; the sodium carbonate dissolves, and there remain 
 behind calcium sulphide, calcium carbonate, and a portion of the foreign 
 substance — the soda residue. 
 
 During the lixiviation the caustic lime present in the crude soda acts 
 upon the sodium carbonate with the assistance of water, and there result 
 calcium carbonate and sodium hydrate. The latter passes into solution 
 with the soda : 
 
 CaO T + Na 2 C 03 = CaCOg + 2NaOH. 
 
 It is possible by the Le Blanc process to obtain a preponderance of 
 scKlium hydroxide if the soda fusion be mixed at the beginning with more 
 carbon, heated intensely and the crude soda be then extracted with hot 
 water. 
 
 When the solution is evaporated the soda will separate from the hot 
 lirpiid as a crystalline ])owder, Na^COg -j- H.^O. It is removed from the 
 li(|uid, and new licpiors are introduced, etc. The mother liquor, the 
 so-called red liipior, contains finally caustic soda and sodium suli)hide 
 almost exclusively. The soda flour is freed from the mother licpior by a 
 centrifugal, dried and calcined-— soda. For further purification 
 
SODIUM CARBONATE. 
 
 291 
 
 it is recrystallized from water, when it separates in large, transparent 
 crystals of the formula Na.^COg -j- 10H2O — crystallized soda. 
 
 The by-products and the refuse in the manufacture of soda — hydrochloric acid and 
 soda residues — must be utilized as fully as possible, because of the enormous competition 
 encountered by the manufacturers. To this end the hydrochloric acid is converted by the 
 process described on page 50, into chlorine and bleaching lime and the soda residues are 
 worked in various ways to get their sulphur content into an available form. Of late 
 years the Chance-Claus method has been adopted. It consists in decomposing the 
 residues with carbonic acid: CaS -(- H.^O -f- CO.^ = CaCO.^ I l.,S, and burning the 
 liberated hydrogen sulphide either with an insufficiency of air, when sulphur will sepa- 
 rate : H.^S -|- O = H.^O + S, or with an excess of air when sulphur dioxide is formed : 
 H^S -r 3O = 11.2^^ + SO.^. The last product is then conducted into lead chambers 
 (p. 189). About 70,000 tons of sulphur are recovered annually (regenerated sulphur). 
 
 2. The ammonia-soda process is based upon the transposition of 
 sodium chloride with })rimary ammonium carbonate to ammonium 
 chloride and primary sodium carbonate : 
 
 NaCl ^ NHdICOg ^NallCOg T NH^Cl. 
 
 This change takes place at the ordinary temperature. The acid sodium 
 carbonate being sparingly soluble in cold water, is converted into sodium 
 carbonate upon ignition : 
 
 2 NaHC 03 ^ Na^COg + CO, + H, 0 . 
 
 The ammonium chloride remains in solution. Ammonia is recovered 
 from it by means of lime. In actual practice carbon dioxide is con- 
 ducted under pressure into a concentrated salt solution saturated with 
 ammonia : 
 
 NaCl -f NHg + CO, -f H^O = NH^Cl -f NaHCO^. 
 
 The temperature must not exceed 40°. The carbonic acid is obtained 
 by burning lime, and when the bicarbonate is heated half of it is 
 recovered. The lime from the calcite is used to generate ammonia from 
 ammonium chloride. The raw material in this process consists therefore 
 of salt and limestone which, with the assistance of ammonium salts, are 
 converted into soda and calcium chloride (p. 50).* 
 
 This process is exceedingly simple from the chemical standpoint. 
 Difficulties arose in building the neces.sary ajiparatus on a technical scale 
 and they militated against the general adoption of the method. Fortu- 
 nately they have lieen completely overcome, and this is due in a large 
 measure to E. Solvay, so that at present the ])roduction of soda by the 
 Le Blanc yirocess is becoming less frequent. It is only in England that 
 the old method holds sway. 
 
 3. The Le Blanc method with it.s improvement.s was in operation for a century before 
 it was displaced by the ammonia-soda process, and now the latter is being .seriously 
 threatened by a dangerous rival — the electrolytic production of caustic alkalies, alkaline car- 
 bonates, chlorine and potassium chlorate ( p. 275). When an af|ueous salt solution is elec- 
 trolyzed chlorine .separates at the anode and sodium at the kathode. 'I'he latter acts 
 
 * Magnesite and magnesia can be substituted for limestone, and lime. 
 
292 
 
 INORGANIC CMKMISTRY. 
 
 immediately on tlie water, producing hydrogen and sodium hydroxide, d'lie free chlorine 
 would produce sodium chloride and hypochlorite or chlorate (pp. 175, 176) if the katluxle 
 and anode licpiors were not separated by a porous diaj)hragm. 1 tifficulties, in the construct- 
 ing of the diaidiragms, confronted the technical utilization of the process. 'There is nosul)- 
 stance which does not gradually disintegrate when used as anode. However, all these 
 objectionable features have been, in a measure, overcome. 'The sodium hydroxide is 
 obtained as such or it is separated in the form of the sparingly soluble bicarbonate ui)on 
 conducting carbonic acid through its solution. 'The chlorine is brought into trade in the 
 liquid form or it is changed to bleaching lime. 
 
 4. Considerable cjuantities of soda are obtained at present from cryolite, a compound 
 of aluminium fluoride and sodium fluoride (AlT'l.,. 3NaFl), which occurs in great deposits 
 in Greenland. The pulverized mineral is ignited with burned lime or chalk ; insoluble 
 calcium fluoride and a very soluble compound of aluminium oxide with sodium oxide, 
 called sodium aluminate (see Aluminium), are produced: 
 
 2(A1F13. 3NaFl) -f 6CaO =r bCaFl^ -f Al.Og. 3Na20. 
 
 The mass is treated with water and carbon dioxide, obtained by burning lime, con- 
 ducted into the solution, which causes the precipitation of aluminium oxide, and sodium 
 carbonate dissolves : 
 
 A\, 0 , . 3 Na ,0 + 3H2O + 3CO, = Al, ( 011)6 + 3 Na,C 03 . 
 
 This method is no longer in use in Europe. It continues of value in North America 
 and in Denmark, because the.se countries confrol large deposits of cryolite. Calcium 
 fluoride is largely employed by them in the manufacture of glass and porcelain ; the 
 aluminium oxide is used for making alum, aluminium sulphate or metal. 
 
 At ordinary temperatures sodium carbonate crystallizes with ten mole- 
 cules of water (Na2C03 -f- 10H2O) in large monoclinic prisms, which 
 crumble upon exposure and become a white powder. It melts at 50° in 
 its water of crystallization, and upon heating a pulverulent hydrate, 
 Na2C03 -j- 2H2O, separates, which in dry air has one molecule of water, 
 and at 100° loses all of this. At 30-50° rhombic prisms of the compo- 
 sition Na2C03 7H2O, crystallize from the aqueous solution. The anhy- 
 drous salt absorbs water from the air but does not deliquesce. It melts 
 at 850° and volatilizes somewhat at a very high temperature. 100 parts 
 of water dissolve 7 parts at 0°, and at 38°, 52 parts of the anhydrous 
 salt. At more elevated temperatures the solubility is less, as in the case 
 of the sulphate. Sodium carbonate has a strong alkaline reaction ; acids 
 liberate carbon dioxide from it. 
 
 Primary Sodium Carbonate, ordinary bicarbonate of soda (Nn- 
 triimi bicarbonicuni), NaHC03, produced by the action of carbon 
 dioxide upon the hydrous secondary carbonate : 
 
 Na2C03 + CO3 + II2O 2NaIlC03. 
 
 It crystallizes without water, in small monoclinic tables; it dissolves, 
 however, at ordinary temperatures in lo-ii parts of water, and possesses 
 feeble alkaline reaction. By heating and by boiling the solution it ])asses 
 into the secondary carbonate with disengagement of carbon dioxide, 
 'i'he dry salt decomposes rafiidly even below 100°. By raj)id evaporation 
 small moiioclinic prisms of the so-called sodium sesquicarbonate, 
 Na2(J03 2NallC03 -j- 2II2O, separate. The salt which deposits in the 
 
SODIUM PHOSPHATES. 
 
 293 
 
 sodium seas of Hungary and E'^vpt, has the composition Na.,CO,. -1- 
 NaHCO, + 2Hfi (p. 289). 
 
 Sodium Nitrate, NaNOg, Chili saltpeter, is found in immense 
 deposits in Peru. The saltpeter earth contains of sodium nitrate from 15 
 to 65 parts out of 100 jiarts, the rest being sodium chloride, a little potas- 
 sium nitrate, potassium perchlorate and sodium iodate (pp. 55, i 78). The 
 soil is extracted with boiling water ; on cooling, crude saltpeter sejiarates ; 
 it is purified by recrystallization. It crystallizes in rhombohedra very 
 similar to cubes, hence designated cubic saltpeter, to distinguish it from 
 the prismatic" potassium saltpeter. It fuses at about 318°. In water 
 it is somewhat more easily soluble than potassium saltpeter. In the air it 
 attracts moisture, hence it is not adapted for the manufacture of gun- 
 powder. In other respects it is perfectly similar to potassium nitrate. It 
 is largely used in the manufacture of nitric acid, and especially in pre- 
 paring potassium saltpeter (p. 279). 
 
 Sodium Nitrite, NaNOg, is prepared like potassium nitrite (p. 280), 
 by heating sodium nitrate with lead, iron, or graphite. It crystallizes 
 more readily than potassium nitrite, and does not deliquesce in the air. 
 It occurs in trade in small colorless crystals, containing from 93 to 98 per 
 cent, of the pure salt. It is largely used in the dye industry for the 
 preparation of the azo- compounds. 
 
 Sodium Phosphates.— The sodium salts of phosphoric acid are less 
 soluble and crystallize better than those of potassium. The trisodiwji 
 phosphate, NagPO^, is made by saturating one molecule of phosphoric acid 
 with three molecules of sodium hydroxide, and crystallizes in six-sided 
 prisms with twelve molecules of water. It has a strong alkaline reaction, 
 absorbs carbon dioxide from the air, and is converted into the secondary 
 salt. 
 
 Disodium phosphate, Na2HPO^, is the most stable of the sodium phos- 
 phates, and hence is generally employed in laboratories (JSfatriuin phos- 
 phoricwti). It may be obtained by saturating phosphoric acid with 
 sodium hydroxide to feeble alkaline reaction, or may be prepared on a 
 large scale by decomposing bone ashes (tri calcium phosphate) with an 
 equivalent amount of sulphuric acid and precipitating the calcium as 
 dicalcium phos])hate with soda. It crystallizes at ordinary temperatures 
 v/ith twelve molecules of water in large monoclinic prisms which effloresce 
 rapidly in the air. It separates from solutions with a temperature above 
 30° in non-efflorescing crystals containing seven molecules of water. It 
 is soluble in 4-5 parts of water, and shows a feeble alkaline reaction. 
 Wlien heated the salt loses water, melts at about 300° and becomes 
 sodium pyrophosphate, Na^P20.j, which crystallizes with ten molecules 
 of water, and upon boiling with nitric acid passes into primary sodium 
 phosphate. 
 
 The primary or ?nonosodium phosphate, NaH2P04, crystallizes with one 
 molecule of water, and exhibits a faintly acid reaction. At 100° it loses 
 its water of crystallization, and at 200° becomes disodium pyn'ophosphate, 
 Na2H2P.^O^, which at 240° forms sodium metaphosphate, NaPOg : 
 
 Na2lbP,0- = 2NaP()g 4 IlgC). 
 
294 
 
 INORGANIC CHEMISTRY. 
 
 \Vc get various modifications of the inelapliospliate, according to the 
 conditions of fusing and cooling; they are probably polymerides, cor- 
 res})onding to the torinulas Na.^P./)^^, NaJ’.,()y, etc. Uiion heating sodium 
 metaphosphate with metallic oxides the latter dissolve, and salts of ortho- 
 phosphoric acid are formed, c. .* 
 
 NaP03 -f CuO = CuNaPO,. 
 
 In this manner, characteristic colored glasses (phosyihorus beads) are 
 obtained with various metals. In blowi)ii)e analysis this behavior serves 
 for the detection of the respective metals. 
 
 The salts of arsenic acid i)erfectly analogous to those of phosphoric acid. Of the 
 antiinoniates may be mentioncHl the disodium pyroautimoniate, Na.dl^Sb./)^ -j hi 1^0, 
 Avhich is insoluble in cold water, and is therefore precipitated from the soluble sodium 
 salts on the addition of dipotassium- pyroautimoniate. 
 
 The phosphates show more plainly than the sulphates that the hydrogen atoms of a 
 polyhydric acid, replaceable by metals, are not of equal importance for the “strength ” 
 of the acid, 'fhe hydrogen first replaced acts like the hydrogen of a strong acid, while 
 the second and third follow it successively. This is also seen in the reactions of the 
 aqueous solutions and may be ex|)lained by the theory of electrolytic dissociation as 
 follows: 'Phe dissociation of phosphoric acid into the ions and 11' is that of 
 
 an acid of medium strength (p. 269). 'Phe second hydrogen atom is dissociated like a 
 feeble acid and the third is not at all dissociated in aqueous solution. Consefjuently the 
 solution of sodium triphosphate does not contain the trivalent anion PO^'" together with 
 the sodium ions. As the third hydrogen atom of the acid shows less tendency to dis.so- 
 ciate than water the followdng transposition immediately takes place : PO^'" -j- 
 
 H' HO' P()^H" -|- HO', so that the solution contains the ions PO^H" — 2Na' 
 
 and Na' — OH', which are the cause of its alkaline reaction. 
 
 Sodium Borate. — The normal salts of boric acid, B(0H)3, and 
 metaboric acid, BO . OH (see p. 242), are not very stable. The ordi- 
 nary alkaline borates are derived from tetraboric acid (HgB^O^), which 
 results from the condensation of four molecules of the normal boric acid : 
 
 4B(0H)3-5H,0 = H,BA. 
 
 The most important of the salts is borax, which crystallizes at ordinary 
 temperatures with ten molecules of water in large monoclinic prisms, 
 Na^B^O^ -f- 10H2O. Borax occurs naturally in some lakes of Thibet, 
 whence it was formerly imported under the name of tinkaL At present, 
 it is prepared artificially by boiling or fusing boric acid with sodium car- 
 bonate. At ordinary temj)eratures, the crystals dissolve in 14 parts of 
 water, at 100° in one-half ])art; the solution has a feeble alkaline reac- 
 tion. When heated to 70° rhombohedra crystallize from.tbe concen- 
 trated solution, and have the composition NaJl^O^ -|- sH^O, formerly 
 known as octahedral borax. Both salts puff up when heated, lose water, 
 and yield a white, porous mass {burned borax^, which fuses at 880° to a 
 transparent vitreous mass (Na.^B^O^). In fusion this dissolves many metal- 
 lic oxides, forming transjiarent glasses {borax beads), which frequently 
 jjossess characteristic colors; thus co])per salts give a blue and chromic 
 oxide gives a green glass. Therefore, borax may be em])loyed in blow- 
 lu'jje tests for the detection of certain metals. Upon this property of 
 
LITHIUM. 295 
 
 dissolving metallic oxides depends the apidication of borax for the fusion 
 and soldering of metals. 
 
 Sodium Silicate (sodium water-glass) is analogous to the i)otassium 
 salt, and is most readily obtained by fusing quartz with sodium sulphate 
 and charcoal. 
 
 The sulphur compounds of sodium are also analogous to those of potas- 
 sium. 
 
 Sodium Nitride, NaNg (sodium azoimide), is the sodium salt of hydrazoic acid, 
 Ngil. It results from derivatives of azoimide, by neutralizing the free acid or by decom- 
 posing the ammonium salt with caustic soda. W. Wislicenus prepared it by heating 
 sodamide (p. 283) to 150-250° in a current of nitrous oxide : 
 
 NaNIIg + N2O = NaNg -f H,0. 
 
 The water produced here decomposes a portion of sodamide into sodium hydroxide and 
 ammonia : 
 
 NaNHg + H 2 O = NaOH -f NHg 
 
 [Ber. 25 (1892), 2084; see also Z. f. anorg. Ch. 6 (1894), 38]. It can be recrystallized 
 from water or may be precipitated by alcohol from water. Its solution reacts alkaline 
 and has a very salty taste. It is not exploded by a blow, but this occurs when it is 
 heated. [Curtius, Ber. 24 (1891), 3346; see also p. 133.] 
 
 Recognition of Sodium Compounds. — Almost all the sodium 
 salts are easily soluble in water, sodium pyroantimoniate, Na2H2Sb20y, 
 excepted; this is precipitated from solutions of sodium salts by potassium 
 pyroantimoniate, and can serve for the detection of sodium. Sodium 
 compounds, exposed in a colorless flame, impart to the latter an intense 
 yellow. The spectrum of the sodium flame is characterized by a very 
 bright yellow line, which, when more strongly magnified, splits into two 
 lines. (See Spectrum Analysis.) 
 
 4. LITHIUM. 
 
 Li = 7.03. 
 
 Lithium occurs in nature only in small quantities, but is tolerably 
 widely disseminated, and is found in some mineral springs and in the 
 ashes of many plants, notably in that of tobacco and the beet. As a com- 
 pound silicate, it occurs in lepidolite or lithia mica; as phosphate (with 
 iron and manganese) in triphylite, and (with aluminium, sodium, and 
 fluorine) in amblygonite. 
 
 The metal is separated from the chloride, or, better, from the more 
 easily fusible mixture of equal ])arts of lithium chloride and potassium 
 chloride, by means of the electric current, and is silver-white in color, 
 decomposing water at ordinary temperatures. Its specific gravity is 0.59. 
 It is the lightest of all the metals, and swims upon naphtha. It melts at 
 180°, and burns with an intense light. It burns energetically in hydro- 
 gen at a red heat to lithium hydride, Lill, which is a comparatively stable 
 
296 
 
 INORGANIC CHEMISI'RY. 
 
 wliile powder. The lilhiuiii salts arc at present j)re|)ared almost entirely 
 from amblygonite ; they are very similar to the salts of sodium, but closely 
 approach those of magnesium. 
 
 Lithium Chloride, LiCd, crystallizes, at ordinary temperatures, in 
 anhydrous, regular octahedra ; below 10°, however, it has two molecules 
 of water, and delicpiesces in the air. 
 
 Lithium Phosphate, and Lithium Carbonate, 
 
 Li^C(),j, are difficultly soluble in water; therefore they are precipitated 
 from solutions of lithium salts by sodium j)hosi)hate or carbonate. l>y 
 strong ignition the carbonate loses carbon dioxide. So far as the.se two 
 salts and also lithium fluoride, which is soluble in very little water, are 
 concerned, lithium aiijiroaches the metals of the second group (|). 273). 
 Its compounds color the flame a beautiful red; the spectrum shows an 
 intense red line together with a faint yellow line. 
 
 AMMONIUM COMPOUNDS. 
 
 Upon p. 128 we observed that ammonia combines directly with the 
 acids to form salt-like compounds, which are analogous to the metallic 
 salts, especially those of potassium with which they are isomorphous. The 
 univalent group, NH^, playing the role of metal in these derivatives, is 
 called ammofiium, and the derivatives of ammonia, ammonium coin- 
 pounds. The metallic character of the group NH^ is confirmed by the 
 existence of ammonium amalgam, which, as regards its external appear- 
 ance, is very similar to the sodium and potassium amalgams. Ammonium 
 amalgam may be prepared by letting the electric current act upon ammo- 
 nium chloride, NH^Cl, viz., by immersing the negative platinum elec- 
 trode into a depression in the ammonium chloride, which is filled with 
 mercury and stands upon the positive electrode. Then, as in the case of 
 the decomposition of potassium or sodium chloride, the metallic ion — 
 ammonium — separates on the negative pole, and combines to an amalgam 
 with mercury. The amalgam may also be obtained if sodium amalgam 
 be covered with a concentrated solution of ammonium chloride: 
 
 (Hg 4- Na) and NH.d yield (Hg + NH,) and NaCl. 
 
 Sodium amalgam. Ammotiium amalgam. 
 
 Ammonium amalgam forms a very voluminous mass with a metallic 
 appearance. It is very unstable, and decomposes rapidly into mercury, 
 ammonia and hydrogen. 
 
 The ac|ueous solution of ammonia reacts strongly alkaline, and from its 
 entire behavior we must assume the existence of ammonium hydroxide 
 (NH/>M) in the solution. This is justified because there are many 
 organic derivatives of ammonium hydroxide, in which the hydrogen of 
 the ammonium is re|)laced bv hvdrocarbon residues; e. g., tetramethyl 
 ammonium hydroxide, N(CIi,,), OH. 'These are thick liquids, of strong 
 basic reaction and, in all respects, are very similar to potassium and 
 sodium hydroxides. 
 
AMMONIUM COMPOUNDS. 
 
 297 
 
 Ammonium Chloride, NH^Cl {Sal ammoniac wyi), is sometimes found 
 in volcanic districts, and was formerly obtained by the dry distillation of 
 camel’s dung (p. 1 26). At present it is prepared almost exclusively from the 
 ammonia water of gas works. This water contains ammonium sesquicar- 
 bonate in addition to theotherammonium salts. It is distilled with lime 
 and the escaping ammonia caught in hydrochloric or sulphuric acid. The 
 ammonium chloride is carefully heated and recrystallized or sublimed. By 
 sublimation it is obtained as a compact, fibrous mass. It dissolves in 2.7 
 parts of cold and in one part of boiling water, and crystallizes from the 
 solution in small, feather-like, grouped octahedra or cubes, of sharp, 
 salty taste. When heated, ammonium chloride sublimes without melting; 
 at the same time a dissociation into ammonia and hydrochloric acid is 
 sustained, but these products recombine to ammonium chloride on cool- 
 ing. The dissociation begins at 280° and is complete at 350°, and the vapor 
 density corresponds to that of a mixture of similar molecules, of NHjand 
 HCl, i. €., J 7-07 f 3646 _ _ 26.76 (02 = 32). A like decomposition is sus- 
 tained by the ammonium chloride when its solution is boiled ; ammonia 
 esca])es and the solution contains some free hydrochloric acid. 
 
 Ammonium Sulphate, (NHO2SO4, crystallizes without water in 
 rhombic prisms, and is soluble in two i)arts of cold and one part of hot 
 water. It fuses at 140°, and by further heating decomj^ioses. Most of it 
 is used as a fertilizer. 
 
 Ammonium Persulphate (pp. 188, 279), (NH4)2S20,^, obtained by the electrolysis 
 of ammonium sulphate, is at present made on a large scale and is used as an oxidant. It 
 consists of very soluble monoclinic crystals. It decomposes when its aqueous solution is 
 evaporated, yielding ammonium sulphate, free sulphuric acid and oxygen. This salt is 
 applied in making the other persulphates. 
 
 Ammonium Nitrate, NH^NOg, is isomorphous with potassium 
 nitrate and deliquesces in the air. It melts at 159° ; at 170° decomposi- 
 tion into nitrous oxide and water commences and is tumultuous at 240° 
 (p. 211). It has been recently applied in the manufacture of blasting 
 material. 
 
 Ammonium Nitrite, NH^N02, is present in minute quantities in the 
 air, and results from the action of the electric spark upon the latter when 
 moist. It may be obtained by the saturation of aqueous ammonia with 
 nitrous acid [Z. f. anorg. Ch. 7 (1894), 34], and in a perfectly pure 
 condition by the decomposition of silver or lead nitrite by ammonium 
 chloride. Heat decomposes it, especially when in concentrated solution, 
 into nitrogen and water (p. 115). 
 
 The decomposition of ammonium nitrite into water and nitrogen and ammonium 
 nitrate into nitrous oxide and water are both exothermic reactions, occurring with the 
 disengagement of heat and are independent of the j^ressure of the disengaged gas ; the 
 components do not reunite to form their original compounds. This is not a case of dis- 
 sociation fp. 274). 
 
 Ammonium Hyponitrite, NIl4-0-N=N-0-NIl4, has been pre]:)ared recently 
 by conducting ammonia into an ethereal solution of hy])onitrous acid. It forms white 
 crystals. It melts at 64° with violent decomposition. It gradually breaks down at the 
 ordinary temperature into ammonia, water and nitrous oxide [see p. 212 and Hantzsch 
 and Kaufmann, Ann. Chem. 292 (1896), 317]. 
 
298 
 
 INORGANIC CHF.MISTKY. 
 
 Ammonium Carbonate.— 'riu; ncufralnx secoiUaty (MI, )./;(),„ 
 separates as a crystalline powder, when aininonia gas is condncied ihrcnigli 
 a concentrated solution of commercial ammonium carbonate. It ])arts 
 with ammonia in the air and becomes the primary or acid salt, NI I,HC( >>3, 
 which, when heated to 58°, dissociates into carbon dioxide, ammonia, 
 and water. 
 
 d'he common commercial, so-called sesquicarbonate of ammonium, is 
 generally a mixture of ])rimary ammonium carbonate with ammonium 
 carbamate (NHjIiCC)., NH.^CCA^. NH,. d'he latter may be obtained 
 by the direct union of carbon dioxide with ammonia; water immedi- 
 ately converts it into the neutral salt (p. 234). lOxjjerience has demon- 
 strated that the commercial salt rpiite often contains carbon dioxide and 
 ammonia in the same proportion as the acid salt, (Nlt jHCO,. It arises 
 in the decay of many nitrogenous carbon comi)Oiinds, e. g., the urine, 
 and was formerly prepared by the dry distillation of bones, horn, and 
 other animal substances. At ])resent it is obtained by heating a mixture 
 of ammonium chloride, or sulphate, with calcium carbonate. It then 
 sublimes as a white, tr;ins])arent, hard mass. 
 
 Primary Ammonium Carbonate, NH^HCO.,, is obtained by satu- 
 rating ammonium hydroxide with carbon dioxide. It is a white, odor- 
 less powder, rather insoluble in water. In aqueous solution it gradually 
 loses carbon dioxide and is changed to the secondary carbonate. 
 
 Ammonium Phosphates. — 'Fhe most important of these is the 
 secondary ammo 7 iium-sodium phosphate, NH^NaHFO^ -f- qll^O, ordi- 
 narily termed salt of phosphorus (Sal ?nicrocosmicum'). It is found in 
 guano and in decaying urine. It can be obtained by the crystallization 
 of a mixture of disodium phosphate and ammonium chloride : 
 
 Na,HPO, -f NH,C 1 = NI^NallPO, + NaCl. 
 
 It consists of large, transparent, monoclinic crystals. When heated it 
 fuses, giving up water and ammonia and forms a transparent glass of 
 sodium metaphosphate, NaPOg (p. 294). It serves in blowpipe tests for 
 the detection of various metals. 
 
 The tertiary ammonium phosphate, (NHJ3P0^, separates, in crystalline 
 form, upon mixing concentrated solutions of phosphoric acid and ammo- 
 nia. Upon drying, it loses ammonia and passes into the secondary salt, 
 (NHJ.^HPO^, which changes to the prhnary salt, (NH^)H2PO^, when its 
 solution is boiled. This is in harmony with what was said on p. 294 rela- 
 tive to the behavior of j)hosi)horic acid. 
 
 Ammonium Nitride, N./NIip, the ammonium salt of hydrazoic acid, obtained from 
 an organic body, diazohi])])uramide, or by saturating hydrazoic acid vvitli ammonia, is pre- 
 ci])itated from its alcoliolic solution byetlier in the form of a snow-white, crystalline pow- 
 der. It separates from alcoliol in com|)act, colorless leaflets, consisting of step-like 
 groups of crystals. It resembles ammonium chloride in this res])ect. It crystallizes from 
 water in large, (rans|\arent ])risms, which soon become opaque. It has a slight alkaline 
 reaction ; it is not hygroscopic ; it dissolves readily in water and in alcohol. It is 
 exceedingly volatile ; it graduallv disappears in the air and it is also carried off by 
 af)ueous and alcoholic vapors. It sublimes, when gently heated, in small, shining 
 prisms; on rapid healing it explodes (Clurtius, [). 132). 
 
METALS OF THE SECOND GROUP. 
 
 299 
 
 Ammonium Sulphide, (NHJ,^S, results ii[)on mixing i volume of 
 hydrogen sulphide with 2 volumes of ammonia at — 18°. It is a white 
 crystalline mass, dissociating, at ordinary temperatures, into NH^SH and 
 NHg. In aqueous solution it also seems to dissociate into its constituents. 
 At 45° it completely dissociates into ammonia and hydrogen sulphide ; 
 
 (NHJ2S = 2NII3 + H^S. 
 
 2 vols. I vol. 
 
 Ammonium Hydrosulphide, NH^SH, is produced upon conduct- 
 ing hydrogen sulphide into an alcoholic ammonia solution. It is com- 
 pletely dissociated at 45° : 
 
 NH^SH = NHg -f HgS. 
 
 I vol. 1 vol. 
 
 It is obtained in aqueous solution by saturating aqueous ammonia with 
 hydrogen sulphide. At first the solution is colorless, but on standing in 
 contact with the air becomes yellow, owing to the formation of ammonium 
 polysulphides, (NH^}2Sn. The so-called yellow ammonium sulphide is 
 more easily obtained by the solution of sulphur in the colorless hydrosul- 
 phide. Both solutions are often employed in laboratories for analytical 
 purposes. 
 
 Recognition of Ammonium Compounds. — All ammonium salts 
 are volatile or decompose upon heating. The alkalies and other bases 
 liberate ammonia from them, which is recognized by its odor and the 
 blue color it imparts to red litmus-paper. Platinum chloride produces 
 a yellow crystalline precipitate of ainmonio-platinum chloride, PtCl^.- 
 2NH^C1, in solutions of ammonium chloride. An excess of tartaric acid 
 precipitates primary ammonium tartrate. 
 
 METALS OF THE SECOND GROUP. 
 
 Be 9. 1 
 
 Mg 24.36 
 
 Ca 40 Sr 87.6 Ba 137.4 
 
 Zn 65.4 Cd 112 
 
 Ilg 200.3. 
 
 The second group of the periodic system (see Table, p 246) comprises 
 chiefly the bivalent metals, which form compounds of the type MeXg, and 
 in their entire deportment exhibit many analogies. Their special rela- 
 tions and analogies are more closely regulated by their position in the 
 periodic system. Beryllium and magnesium belong to the two small 
 periods whose members are similar but do not show complete analogy. 
 Beryllium exhibits many variations from magnesium, and in manv prop- 
 erties approaches aluminium; just as lithium attaches itself to magnesium 
 (P; 273)- 
 
 The metals calcium, strontium, and barium constitute the second 
 members of the three great periods, are very similar to one another (p 
 244), and in accord with their strong basic character, attach themselves 
 to the alkali metals — potassium, rubidium and ctesium. Zinc, cadmium. 
 
300 
 
 INORGANIC CHEMISTRY. 
 
 and mercury, wliicli correspond to them and constitute the second sub- 
 group, really belong to the right, negative sides of the three great jjeriods. 
 'They fall in with the heavy metals, are much less basic, and resemble the 
 alkaline earth metals only in their combination types. In conseipience 
 of the double periodicity of the three great periods both sub-groui)S (Ca, 
 Sr, Ba and Zn, Cd, Hg) exhibit many analogies to magnesium and 
 beryllium. 
 
 I. GROUP OF THE ALKALINE EARTHS. 
 
 Calcium, . . Ca = 40 Strontium, . . Sr = 87.6 IJarium, , . P>a = 137.4. 
 
 The metals of this group are termed alkaline earth metals, because 
 their oxides attach themselves in their i)roperties, on the one side to the 
 oxides of the alkalies, upon the other to the real earths (alumina, etc.). 
 'I'heir atomic weights bear almost the same ratio to one another as those 
 of the alkali metals, hence the alkaline earth metals show the same grada- 
 tion in properties as the elements of the ])otassium group. With in- 
 crease in atomic weight and atomic volume, their chemical energy and 
 basicity become greater. Barium decomposes water energetically, and 
 oxidizes more readily than strontium and calcium. In accord with this, 
 we find barium hydroxide a stronger l)ase ; it dissolves rather readily in 
 water, does not decompose upon ignition, and absorbs carbon dioxide 
 rapidly from the air. Barium carbonate is also very stable, fuses at a 
 white heat, and only disengages a little carbon dioxide. Calcium hy- 
 droxide, on the other hand, dissolves with more difficulty in water, and 
 when ignited, breaks down into water and calcium oxide; the carbonate 
 also yields carbon dioxide when similarly treated. In its entire charac- 
 ter, strontium stands between barium and calcium. All these affinity re- 
 lations find full expression in the heat of formation of the corresponding 
 compounds (p. 327). 
 
 While the alkaline earth metals are similar to the alkalies in their 
 free condition and in their hydroxides, they differ essentially from them 
 by the insolubility of their carbonates and phosphates, and still more of 
 their sulphates. Barium sulphate is only soluble to the slightest degree 
 in water and acids, while strontium sulphate is about 40 times and cal- 
 cium sulphate about 800 times as soluble (p. 309). 
 
 The atomic weights of the.se metals have been determined in part from their specific 
 lieats, but mainly from their isomorphism with the metals of the magnesium group 
 (pp. 253, 255). 
 
 1. CALCIUM. 
 
 Ca = 40. 
 
 Calcium belongs to the class of elements most widely distributed upon 
 the earth’s surface. As calcium carbonate (limestone, marble, chalk) 
 and the sulphate (gyi)sum, alabaster), it represents immense deposits in 
 
CALCIUM. 
 
 301 
 
 all stratified formations. As phosphate, it constitutes phosphorite, as 
 fluorite, both of which are abundant. As silicate, it is found in 
 most of the oldest crystalline rocks. 
 
 The metal is obtained, according to Moissan, by heating calcium 
 iodide with an excess of sodium to a dark-red heat. The liberated 
 metal dissolves in the sodium, and when cold the latter is removed by 
 anhydrous alcohol. It dissolves with the evolution of hydrogen to 
 sodium alcoholate. The calcium remains as a brilliant white crystalline 
 powder. It can also be prepared by the electrolysis of the fused chloride 
 or iodide. Although the affinity of calcium for oxygen is less than that 
 of the alkalies, yet the oxide (also barium oxide and strontium oxide) 
 cannot be reduced to metal by ignition with carbon, iron, or sodium — 
 due, probably to the non-fusibility of the oxide. These oxides, how- 
 ever, are reduced in the electric furnace; the liberated metals com- 
 bine at once with carbon to yield metallic carbides (pp. 253, 307). 
 
 Calcium is a yellow, shining metal, of specific gravity 1.55-1.6. In 
 dry air it is tolerably stable, in moist it is covered with a layer of 
 hydroxide. It decomposes water with considerable energy. It fuses at 
 a red heat, and in the air burns with a brilliant yellow light. If the 
 metal be heated in nitrogen to a red heat calcium nitride, Ca3N2, results. 
 This is a brown mass which water decomposes with the formation of 
 calcium hydrate and ammonia. Lithium, strontium, barium and mag- 
 nesium behave similarly (p. 316). Calcium slowly unites with hydrogen 
 at the ordinary temperature, but rapidly at a red heat, forming calcium 
 hydride, CaHj, an earthy, gray powder, which decomposes water more 
 energetically than calcium itself. 
 
 Calcium Oxide, CaO (lime), may be obtained ])ure by igniting the 
 nitrate or carbonate. It is prepared on a large scale by burning the 
 ordinary limestone or marble (CaCOg) in lime-kilns. It is a grayish- 
 white amorphous mass, which becomes crystalline at 2500° and liquid 
 like water at 3000°. It vaporizes at higher temperatures (Moissan). 
 The oxyhydrogen flame thrown upon a piece of lime causes it to emit an 
 extremely intense white light (Drummond’s lime light). In the air 
 lime attracts moisture and carbon dioxide, becoming calcium carbonate ; 
 burned lime unites with water with evolution of much heat, breaking 
 down into a white voluminous powder of calcium hydroxide, Ca(OH)2 — 
 slaked Ihjie. 
 
 When limestone contains large quantities of alumina, magnesium carbonate, or other 
 constituents, the lime from it slakes with difficulty, and is known as poor lime, to dis- 
 tinguish it from pure fat or rich lime, which is readily converted into a powder of 
 hydroxide with water. 
 
 Calcium Hydroxide, Ca(OH)2 (slaked lime), is a white, porous 
 powder which dissolves with difficulty in cold water (i part in 760 parts), 
 but with still more difficulty in warm water; the solution saturated in 
 the cold (lime-water) becomes cloudy uj)on warming. It has a strong 
 alkaline reaction. Milk of Imie is slaked lime mixed with water. In the 
 air it attracts carbon dioxide and forms calcium carbonate. At a red 
 heat it decomposes into oxide and water. 
 
302 
 
 INORGANIC CHEMISTRY. 
 
 Slaked lime is employed in the j)re}jaration of ordinary mortar, a mix- 
 ture of calcium hydroxide, water and sand. The liardening of the mor- 
 tar in the air depends mainly upon the fact that the calcium hydroxide 
 combines with the carbon dioxide of the air to form the carbonate, and 
 at the same time acts upon the silicic acid of the sand forming a calcium 
 silicate, which, in time, imi)aits durability to the mortar. 
 
 Hydraulic mortar, or cement, is produced by gently igniting a mixture 
 of limestone or chalk with aluminium silicate (clay) and (piartz powder. 
 On stirring the jiowdered, burnt mass with water it soon hardens, and is 
 then not affected by water. Some naturally occurring limestones, con- 
 taining upwards of 20 jier cent, of clay, yield hydraulic cements, without 
 any admixtures after burning. Their comi)Osition varies, and also the 
 ])rocess of their hardening ; the latter, however, depends princij)al]y upon 
 the formation of calcium and aluminium silicates. 
 
 Calcium Peroxide, Ca02, is precipitated as a hydrate in crystalline 
 leaflets, if lime-water be added to a solution of hydrogen jjeroxide. It 
 contains eight molecules of water, which it gradually loses in dry air. 
 
 The halogen derivatives of calcium, like those of other metals, are 
 })repared by the solution of the oxide or carbonate in the aqueous haloid 
 acids. They are formed also by the direct union of calcium with the 
 halogens; this occurs with evolution of flame. Technically, calcium 
 chloride is often obtained as a by-product, e. g., in the preparation of 
 ammonia (see Soda). 
 
 Calcium Chloride, CaCl2, crystallizes from aqueous solution with six 
 molecules of water, in large, transparent, six-sided prisms, which deli- 
 quesce in the air. In vacuo it loses four molecules of water. When 
 heated, it melts in its water of crystallization, loses water, but it is only 
 after it has been exposed above 200° that it becomes anhydrous; then it 
 is a white, porous mass. The dry salt melts at 806°, and solidifies on 
 cooling to a crystalline mass, which attracts water energetically, and may 
 be employed in the drying of gases and liquids. The dry calcium chlo- 
 ride also absorbs ammonia, forming the compound CaCl2 . 8NH3. The 
 crystallized hydrous salt dissolves in water with reduction of temperature; 
 by mixing with snow or ice the temperature is lowered to ^ — ^_48°. Upon 
 fusing the dry chloride in moist air it will partly decompose into the 
 oxide and hydrogen chloride. 
 
 Calcium bromide and iodide are very similar to the chloride. 
 
 Calcium Fluoride, CaFl.^, occurs in nature as fluorite, in large cubes 
 or octahedra, or even in compact masses. It is often discolored by im- 
 ])urities. It is found, in sparing quantities, in the ashes of idants, bones, 
 and the enamel of the teeth. A soluble fluoride added to the solution of 
 calcium chloride throws down insoluble calcium fluoride as a white vol- 
 uminous ])recii)itate. If, however, the calcium cldoride solution be poured 
 into a boiling dilute solution of potassium fluoride, the calcium fluoride 
 is obtained in little transparent crystals. 
 
 'I'he fluoride is jierfectly insoluble in water, and is only decomposed by 
 
CALCIUM HYPOCHLORITE. 
 
 303 
 
 Strong acids. It fuses easily at a red heat, serving, therefore, as a flux in 
 the smelting of ores. When heated it phosphoresces. 
 
 Calcium Hypochlorite, Ca(C10)2, is not known in a pure condi- 
 tion. The so-called bleaching lime or chloride of lime, obtained by con- 
 ducting chlorine, at ordinary temperatures, over slaked lime, contains 
 calcium hypochlorite as the active principle. 
 
 According to analogy to the action of chlorine upon potassium or 
 sodium hydroxide (p. 278), the reaction in the case of calcium hydroxide 
 may be expressed by the following equation : 
 
 2 Ca(OH )2 + 2 CI 2 = Ca(C10)2 -f- CaCq -f 2 H 2 O. 
 
 This would incline us to regard chloride of lime as a mixture of cal- 
 cium hypochlorite, calcium chloride, and water. In accordance with 
 the equation, the completely chlorinated chloride of lime must contain 
 48.9 per cent, of chlorine, which is never the case, because a portion of 
 the calcium hydroxide invariably remains unaltered. Calcium chloride 
 does not exist free in bleaching lime, because it is not withdrawn from 
 the latter by alcohol, and nearly all the chlorine of the bleaching lime 
 can be expelled by carbon dioxide. It is therefore probable that the 
 Cl 
 
 compound, Ca<Q^I = CaCl20 (Odlingand Lunge), is present in bleach- 
 ing lime; its formation, then, from calcium hydroxide would correspond 
 to that of the chloride and hypochlorite from the caustic alkalies; 
 
 NaOH , 
 NaOH “t- '-h 
 
 Chloride of lime is a white, porous powder with an odor resembling 
 that of chlorine. The aqueous solution has a strong alkaline reaction, 
 and bleaches. It decomposes in the air as the carbon dioxide of the 
 latter liberates hypochlorous acid. Even in closed vessels it gradually 
 decomposes with the evolution of oxygen ; the decomposition is hastened 
 I)y sunlight and heat, and may occur with explosion. Hence chloride 
 of lime should be preserved in loosely closed vessels, in a cool, dark 
 place. 
 
 Dilute hydrochloric or sulphuric acid will expel chlorine from chloride 
 of lime ; the quantity liberated is just twice that which the hypochlorite 
 eventually found in it contains (p. 175) : 
 
 CaCl(Ori) -I- 2 IICI = CaCq + 11,0 -f Cl,. 
 
 By the action of sulphuric acid all the chlorine of the bleaching 
 powder is liberated, whereby calcium chloride first yields hydrochloric 
 acid : 
 
 CaCl(OCl) + = CaSO^ + 11,0 f Cl,. 
 
304 
 
 INORGANIC CHEMISTRY. 
 
 The api)lication of chloride of lime for the production of chlorine in 
 chlorine bleaching and disinfection is based on this deiiorlment. 
 
 The (juantity of chlorine set free by acids from the chloride of lime 
 rei)resents its quantity of so-called active chlorine ; good chloride of lime 
 should contain at least 25 per cent. 
 
 When a small (iuantity of coballic oxide or a cobalt salt is added to the solution of 
 bleaching lime, and heat aj)[)lied, a regular stream of oxygen is disengaged ; this is an 
 advantageous method of preparing oxygen. Other oxides, like those of manganese, 
 copper, and iron, behave similarly. In this reaction there occurs, apparently, a cata- 
 lytic action of the oxides, d’he reaction is explained, doubtless, in the same way as the 
 action of hydrogen peroxide upon certain oxides (sec pp. 8 S, 102). 
 
 On warming bleaching lime with ammonia, the following decomposition occurs : 
 
 3CaCl(OCl) I 2NII3 = 3CaCb + 3H,0 + N^. 
 
 This reaction can be used for the preparation of nitrogen. 
 
 Calcium Sulphate, CaSO^, is very abundant in nature. In an 
 anhydrous condition it forms the mineral anhydrite, crystallizing in forms 
 of the rhombic system. With two molectiles of water it occurs as 
 gypsum, in large monoclinic crystals or in granular, crystalline masses 
 (alabaster, etc.). It also separates as a fine crystalline ])owder, CaSO^ d- 
 2H2O, vv’hen soluble calcium salts are precipitated with sulphuric acid. 
 Calcium sulphate dissolves with difficulty in water; i ])art at average 
 temperatures is soluble in 400 parts of water. When heated to 110° 
 gypsum loses all its water, and becomes burjtt g^\)’s,\\m ; when this is pul- 
 verized and mixed with water, it forms a paste which hardens to a solid 
 mass in a short time. The hardening is dependent ui)on the reunion of 
 anhydrous calcium sulphate with two molecules of water. On this de- 
 pends the use of burned gypsum for the production of casts, figures, 
 etc. In case gypsum is heated above 160° (dead-burnt gypsum), it will 
 no longer harden with water ; the naturally occurring anhydrite behaves 
 in the same manner. 
 
 Calcium Nitrate, Ca(N03)2, is produced by the decay of the nitrog- 
 enous organic substances in the presence of lime, therefore, it is fre- 
 quently found as an efflorescence upon walls (in cattle stables). It 
 crystallizes from water in monoclinic prisms, containing four molecules 
 of water; the anhydrous salt deliquesces in the air and is easily soluble 
 in alcohol. By the action of potassium carbonate or chloride, calcium 
 nitrate can be transposed into ])otassium nitrate (p. 280). 
 
 Calcium Phosphates. — The tertiary phosphate, Ca3(PO^)2, is found 
 in slight quantities in most of the mountain rocks. In combination with 
 calcium chloride and fluoride, it crystallizes as apatite. As phosphorite, 
 it forms compact masses, more or less intimately mixed with other con- 
 stituents, and occurs in immense deposits in Spain, France, Germany, 
 Russia and Florida. When these minerals disintegrate the calcium phos- 
 ])hate ])asses into the soil and is absorbed by the plants. In the latter, 
 it accumulates chiefly in the seeds and grains. In the animal kingdom, 
 it is principally found in the bones, the ashes of which contain upwards 
 of 82 per cent, of tricalcium ])hosj)hate. T'ertiary calcium phosphate 
 is nearly insoluble in water. If disodium phosphate be added to the 
 
CALCIUM CARBONATE. 
 
 305 
 
 aqueous solution of a calcium salt, and then ammonium hydroxide, it 
 will separate as a gelatinous precipitate, which, after drying, becomes a 
 white amorphous powder. It is very readily soluble in acids, even acetic. 
 
 The secondary calcium phosphate, CaHPO^ -f 2H2O, is sometimes 
 present in guano, in the form of small, shining prisms; it separates as 
 an.amorphous precipitate, if disodium phosphate be added to a solution 
 of calcium chloride mixed with some acetic acid. When ignited, it 
 passes into pyrophosphate, Ca.^P.^O.^ (p. 216). 
 
 The primary phosphate, Ca(H.,P04).„ is produced by the action of sul- 
 phuric or hydrochloric acid upon the first two phosphates. It is readily 
 soluble in water. It changes to metaphosphate when it is heated : 
 
 Ca(ITPOj2 - Ca(P 03 b + 211,0. 
 
 Calcium phosphate is present in all plants. Its presence in the soil is, therefore, an 
 indispensable condition for its fertility. When there is a scarcity of phosphoric acid it 
 must be added. To this end, bone meal and pulverized phosphorite were formerly em- 
 ployed. Since, however, the phosphoric acid is contained in these substances as tri- 
 calcium phosphate, which is not easily absorbed by the plants, the primary phosphate is 
 extensively employed at present as a fertilizer, or, better, the mixture resulting from the 
 action of sulphuric acid upon the tertiary salt, which also contains calcium sulphate. 
 Superphosphate is the name applied to the resulting mass. 
 
 The Thomas slag, obtained in the dephosphorization of iron oi'es by the Gilchrist- 
 Thomas process, constitutes a very valuable and important crude product from which to 
 prepare calcium phosphate. It also contains a crystallized phosphate of the formula 
 Ca^F.^O;,, which may be regarded as a derivative of the normal hydrate, r( 0 H )5 : 
 
 2P(0H)5— H.,0 = HgP^Og. 
 
 Calcium Carbonate, CaCOg, is very widely distributed in nature. 
 It crystallizes in two crystallographic systems, hence is dimorphous. In 
 rhombic crystals, with the specific gravity 3.0, it forms aragonite. In 
 hexagonal rhombohedra, with the specific gravity 2. 7, it occurs as calcite. 
 Iceland spar, employed for optical purposes, is perfectly pure, transparent 
 calcite. The common calcite, which constitutes immense mountain 
 chains, is an amorphous or indistinctly crystalline stratum, and is usually 
 mixed with other constituents, as clay. When the limestone is granular 
 and crystalline, it is termed marble. Dolomite also forms large layers, 
 and is a compound or amorphous mixture of calcium and magnesium 
 carbonates, with generally an excess of the former. Chalk is very pure 
 amorphous calcium carbonate, consisting of the shells of microscopic sea 
 animals. Calcium carbonate is, further, a regular constituent of all plants 
 and animals ; the shells of eggs, of mussels, even corals and pearls, consist 
 chiefly of it. 
 
 A soluble carbonate, added to the aqueous solution of a calcium salt, 
 precipitates calcium carbonate as a white, amorphous powder, which soon 
 becomes crystalline. In the cold, it assumes the form of calcite ; upon 
 boiling the liquid, it generally changes into aragonite crystals. Magnesite 
 (MgCOg), siderite (P'eCO.j), smithsonite (ZnCO.^) and rhodochrosite 
 (MnCOg) are isomorphous with calcite, while witherite (BaCO,,), stron- 
 tianite (SrCOg) and cerinssite (PbC03)are isomorphous with aragonite. 
 
 Calcium carbonate is almost insoluble in pure water; but dissolves 
 somewhat in water containing carbon dioxide, as it very probably is 
 26 
 
3o6 
 
 INORGANIC CHEMISTRY. 
 
 changed lo ihe primary carljonate, Ca(HCC).,)2. One tliousand parts of 
 water at 15° and incdiuin barometric j)res.siire dissolve about 0.385 part 
 of ])rimary calcium carbonate [Z. f. anorg. Ch. 17 (1898) 170]. Id)r 
 this reason, we find calcium bicarbonate dissolved in all natural waters. 
 When the solution stands ex|)Osed, or if it be heated, carbon dioxide 
 escai)es, and the secondary carbonate again separates out. The formation 
 of lime scales, thermal tufts, stalactites, boiler scales and similar deposits 
 are due to this. Calcium carbonate, like all carbonates, is decomj)osed 
 by acids, with evolution of carbon dioxide. 
 
 At a red heat, it dissociates into calcium oxide and carbon dioxide. Tlie change 
 begins at 600°. Tlie tension of dissociation (]). 274) is 85 mm. at 860°, and at 1040° 
 510-520111111. For this reason calcium carbonate is not decomposed wlien heated in a 
 sealed tube. 
 
 Calcium Silicate, CaSiO,, occurs as white, crystalline wollastonite. 
 It is also a constituent of most natural silicates and of the artificial sili- 
 cate fusion — glass. 
 
 Glass. — The silicates of pota.ssinm and sodium are readily fusible and soluble in water. 
 The silicates of calcium and the other alkaline earths are insoluble, very difficultly fusible, 
 and generally crystallize when they cool. If, however, the two silicates be fused together, 
 ail amorphous, transparent mass, of average fusibility, results ; it is only slightly attacked 
 by water and acids — it is glass. To prejiare the latter, a mixture of sand, lime, and 
 soda, or potash, is heated to fusion in a muffle furnace. Instead of the carbonates of 
 potassium and sodium a mixture of the sulphates with charcoal can be employed ; the carbon 
 reduces the sulphates to sulphides, which form silicates when fused with silicon dioxide. 
 
 Glass, depending on its constituents, arranges itself into two clas.ses : 
 
 1. Lime glass (free from lead). 
 
 2. Lead glass. 
 
 1. The lime glass eonsists either of sodium-lime (soda glass) or potassium-lime (potash 
 glass). 
 
 Soda g/ass — a mixture of sodium and calcium silicates — buses readily, and is employed 
 for window-panes and ordinary glass vessels. Poiash or BoJnviian glass, also called 
 crown glass, consists of calcium and potassium silicates, is not very fusible, is harder 
 and withstands the action of water and acids better than soda glass ; it is, therefore, 
 employed in the manufacture of chemical glassware. 
 
 The best lime glass, normal glass, is a trisilicate corresponding to the formula- 
 
 KgO (or Na20) CaO -f- bSiOg. 
 
 2. Lead glass is constituted like lime glass, eontaining, however, lead oxide for 
 calcium oxide. 
 
 Glass crystal or flint glass is composed of potassium and lead silicates. It is not as 
 hard, fuses with tolerable readiness, refracts light strongly, and when polished, acquires 
 a clear luster. On this account it is employed for optical purposes (for lenses, prisms) 
 and is u.sed in ornamental glassware. Strass, a lead glass containing boron trioxide, is 
 used to imitate ])recious stones. The opaque varieties of enamel consist of lead glass 
 and in the fused glass are insoluble admixtures, as tin dioxide and calcium phosphate. 
 
 Ordinary winclow or bottle glass is obtained by the fusion of rather impure materials ; 
 in consc(]uence of Ihe ])resenceof ferrous oxide it is ordinarily colored green. To remove 
 this coloration, manganese peroxide, MnO,^, is added to the fusion. It oxidizes a por- 
 tion of the ferrous to ferric oxide, the silicate of which is colored slightly yellow, while 
 manganese foi ins a violet silicate. 'I'liese colors, violet and green, almost neutralize each 
 other as complementaries. d’he colored glasses contain silicates of colored metallic 
 oxides: chromic and cn|)ric oxides color green; cobaltic oxide, blue ; cupn)us oxide, a 
 ruby red, etc. 
 
 'The sulphur compounds of calcium are very much like those of the alkalies. Calcium 
 
STRONTIUM. 
 
 307 
 
 Sulphide, CaS, is most readily obtained by beating the sulphate with carbon, and is a 
 whitish-yellow mass. It melts in the electric furnace and crystallizes on cooling. When 
 it is dissolved in water we get Calcium Hydrosulphide, Ca(SlI)2 (in addition to cal- 
 cium hydroxide), which decomposes on boiling the aqueous solution. When calcium 
 oxide is ignited with sulphur in a closed crucible a yellowish-gray mass is obtained, which 
 consists of calcium polysulphides and sulphate. Milk of lime boiled with sulphur yields 
 a deep-yellow solution of calcium polysulphides and calcium hyposulphite. When these 
 solutions are acted upon by acids, finely divided sulphur — milk of sulphur — is precipi- 
 tated and hydrogen sulphide set free. If the reverse occur, viz., the addition of a .solu- 
 tion of polysulphides to an excess of dilute acids, hydrogen persulphide will separate 
 (p. no). 
 
 Calcium Carbide, CaCj. Its importance in illumination has been dwelt upon on 
 pp. 154, 155. Fr. Wohler (1862) first described it. He obtained it by heating intensely 
 an alloy of zinc and calcium with carbon : 
 
 ZnCa -(- Cj — CaCg -j- Zn. 
 
 W Borchers was the first person to prepare it in the electric furnace. In 1891 he emphasized 
 the fact, after years of experience, that all oxides could be reduced by carbon heated 
 electrically. Wilson made carbide accidentally on a large scale in 1892. Moissan fol- 
 lowed shortly after. Both chemists obtained the carbide in crystals. It maybe procured 
 pure in shining gold-like, non-transparent crystals, if burnt marble is heated electrically 
 with sugar carbon : 
 
 CaO + 3C = CaCj + CO. 
 
 Its specific gravity is 2.2. It resists most of the strongly reacting substances, yet water 
 and dilute acids decompose it energetically with liberation of acetylene (p. 154). At 
 present it is made in many factories from unslaked lime and coal or coke (see the publi- 
 cation of Liebetanz, p. 155). Strontium and barium carbides are similar to calcium 
 carbide. 
 
 2. STRONTIUM. 
 
 Sr = 87.6. 
 
 This element is rather rare in nature, and is principally found in stron- 
 tianite (strontium carbonate) and celestite (strontium sulphate). Its 
 compounds are very similar to those of calcium. 
 
 The metal is obtained by the electrolysis of fused strontium chloride. 
 It has a brass-yellow color, and a specific gravity of 2.5. It oxidizes in 
 the air and burns with a bright light when heated. It decomposes water 
 at the ordinary temperature. 
 
 Of the comj)ounds of strontium we may mention the following: 
 
 Strontium Oxide, SrO, is most readily obtained by igniting the 
 nitrate. It unites with water, with strong evolution of heat, forming 
 
 Strontium Hydroxide, Sr(OH)2, which is more readily soluble in 
 water than calcium hydroxide. It crystallizes from aqueous solution 
 with eight or nine molecules of water. When ignited it decomposes into 
 strontium oxide and water, but with more difficulty than calcium hy- 
 droxide. 
 
 Strontium Chloride, SrCl2 -f- 6H2O, crystallizes from water in hex- 
 agonal tables, which deliquesce in the air; it is easily soluble in alcohol. 
 
 Strontium Sulphate, SrSO^, is much more difficultly soluble in water 
 than calcium sulphate, but is not as insoluble as barium sulphate [one 
 liter dissolves 107 mg. at 
 
3o8 
 
 INORGANIC CHEMIS^IRY. 
 
 Strontium Nitrate, Sr(N()J^, is obtained l)y dissolving tlic carl)on- 
 ate in nitric acid, and is rcatlily soluble in water, but insoluble in alcohol. 
 It crystallizes from warm solutions in anhydrous octahedra, but from cold, 
 with four molecules of water, in monoclinic i)risms. Mixed with com- 
 bustible substances it colors the flame a beautiful carmine-red, and for 
 this reason is employed in ji^ rotechny. 
 
 Strontium Carbonate, SrCO.,, is [irecipitated from aqueous solutions 
 ol strontium salts by soluble carbonates, as an amorj)hous, insoluble pow- 
 der. When ignited (iioo°) it breaks down into strontium oxide and 
 carbon dioxide, 'bhis decom])osition does not, however, occur as easily 
 as with calcium carbonate. 
 
 8. BARIUM. 
 
 Ba = 137.4. 
 
 Barium occurs in nature in large masses, as heavy spar (or barium sul- 
 phate), and as witherite (barium carbonate). All its compounds are dis- 
 tinguished by their high specific gravity, hence the name barium, from 
 [■iafthq, heavy. In accordance with its general character, barium is a 
 stronger liasic metal than either strontium or calcium (p. 299). 
 
 The barium salts are either prepared from the natural witherite, by dis- 
 solving it in acids, or from heavy spar. The technical procedure con- 
 sists in heating barium sulphate with carbon and calcium chloride when 
 barium chloride and potassium sulphide (in addition to carbon monoxide) 
 are formed : 
 
 BaSO, + 4C + CaCb = Bad, + CaS + 4CO. 
 
 Metallic barium, like strontium and calcium, was first obtained (Hum- 
 phry Davy, t8o8) by electrolysis. Bunsen irsed the fused chloride for this 
 purpose. The following method is more convenient : Sodium amalgam 
 is added to a hot saturated barium chloride solution ; the sodium dis- 
 ])laces the barium, which forms an amalgam with the mercury. The 
 resulting liquid barium amalgam is kneaded with water, to remove all the 
 sodium, and then heated in a stream of hydrogen, to volatilize the 
 mercury. 
 
 Barium is a bright yellow metal, of specific gravity 3.6. It fuses at a 
 red heat, but vaporizes with difficulty. It is rapidly oxidized in the air; 
 like sodium it decomposes water very energetically, even at the ordinary 
 temjierature. 
 
 Barium Oxide, BaO, is obtained by the ignition of barium nitrate. 
 It is a gray, amorphous mass, of specific gravity 5.5, and is fusible in the 
 oxyhydrogen flame. With water it yields the hydroxide, with evolution 
 of much heat. 
 
 Barium Hydroxide, Ba(OH)2, is ])recipated from concentrated 
 solutions of barium salts by potassium or sodium hydroxide, not, however, 
 by ammonium hydroxide. It is olitained on a large scale by igniting 
 heavy spar with carbon, when barium sulphide is formed, and its aqueous 
 solution desulj)hurized with zinc oxide: 
 
 BaS -|- ZnO | H,() = ZiiS + Ba(()H), ; 
 
BARIUM. 309 
 
 also Upon heating barium carbonate together with carbon in a current 
 of steam : 
 
 BaCOa + C + lip = Ba(OH)2 + 2CO. 
 
 At the ordinary temperature it dissolves in 20 parts of water; upon 
 boiling, in 3 parts. From aqueous solution it crystallizes with eight mole- 
 cules of water in four-sided prisms or leaflets. The solution — called 
 baryta water — is strongly alkaline and is very similar to the alkalies. 
 When exposed to the air it absorbs carbon dioxide and becomes turbid, 
 with separation of barium carbonate. At a red heat it fuses, like the 
 caustic alkalies, without decomposition, and on cooling solidifies to a 
 crystalline mass. 
 
 Barium Peroxide, BaO.^, is produced when barium oxide (ignited 
 barium nitrate) is heated in a current of dry air (free from carbonic acid) 
 or oxygen ; it always contains barium oxide. To purify it, the commercial 
 peroxide is rubbed together with water and added to cold, very dilute 
 hydrochloric acid, until the latter is almost saturated. An excess of 
 baryta water is added to the solution, containing barium chloride and 
 hydrogen peroxide, when hydrated barium peroxide, BaO.^ + 8H2O, sepa- 
 rates in shining scales, which, upon warming, readily lose water and 
 break down into a white powder consisting of barium peroxide. The 
 latter is a compact gray mass when obtained directly from the oxide. 
 
 The peroxide dissolves in dilute acids, with production of hydrogen 
 peroxide. Concentrated sulphuric acid sets free ozonized oxygen from 
 it. When strongly ignited (at about 700°) it decomposes into barium 
 oxide and oxygen. See p. 81 for its use in the technical production of 
 oxygen. It is also employed for bleaching purposes and for the prepara- 
 tion of hydrogen peroxide; see pp. 99, 100. 
 
 Barium Chloride, BaCl2, crystallizes from aqueous solution, with 
 two molecules of water, in large, rhombic tables, which are stable in the 
 air. It dissolves readily in water, and is poisonous, like all soluble 
 barium salts. 
 
 Barium Nitrate, Ba(N03)2, crystallizes in anhydrous, shining octa- 
 hedra, of the regular system, soluble in 12 parts of cold and 3 parts of 
 hot water. Technically it is made by mixing concentrated solutions of 
 barium chloride with sodium nitrate, when barium nitrate is precipitated. 
 It is employed as a green fire in pyrotechny. 
 
 Barium Sulphate, BaSO^, is found in nature as heavy spar, in 
 rhombic j)risms, with a specific gravity of 4.6. It is obtained artificially 
 by the precipitation of barium salts with sulphuric acid, in the form of a 
 white, amorphous powder, almost insoluble in water and acids. Under 
 the name of permanent white, it is used as a paint, as a substitute for poi- 
 sonous white lead, from which it is also distinguished by its unaltera- 
 bility. 
 
 Barium Persulphate, BaS.^O^ fpp. 188, 198), obtained by rubbing the ammonium 
 salt together with barium hydroxide, crystallizes from its solution {in vacuo) in prisms, 
 containing four molecules of water of crystallization. It soon decomposes, yielding 
 barium sulphate, suh)huric acid, and oxygen : 
 
 BaSPg -f II2O = BaSO^ i 11 , 80 ^ -p O. 
 
310 
 
 INORGANIC CHEMISTRY. 
 
 Barium Carbonate, IkiCO.,, as witherite, occurs iu sliiuing, rlioml)ic 
 crystals, and is i)recipilatc(l from barium solutions by soluble caibonatcs, 
 as a white, amorphous i)()wder. It fuses at a white heat, and loses carbon 
 dioxide at 1400-1500°. 
 
 Barium Sulphide, llaS, is obtained l)y igniting the sul|>hate with 
 carbon. It dissolves in water, with decomjiosition into hydroxide and 
 hydrosulphide. 
 
 RECOGNITION OF THE COMPOUNDS OF THE ALKALINE 
 
 EARTHS. 
 
 The carbonates and phosphates of this group are insoluble in water; 
 hence are precipitated from the aqueous solutions of their salts upon the 
 addition of soluble carbonates and phosphates (of the alkalies). The 
 sulphates are also sparingly soluble in acids; for this reason they are 
 thrown down from acid solutions l^y soluble sulphates or free sulf)huric 
 acid; the precipitation is complete, even with calcium, if alcohol be 
 added to the solution, d'he hydroxides of the alkaline earths, which are 
 more or less soluble in water, are only precipitated by sodium or potas- 
 sium hydroxide from concentrated solutions. In solutions of barium 
 salts hydrofluosilicic acid produces a crystalline y^reciju'tate of barium 
 silicofluoride, BaSiFlg, and potassium chromate one of barium chromate, 
 BaCrO,. 
 
 Koblrausch and Rose determined, by means of the electric conductivity of solutions, 
 the solubility of many of the salts which have been mentioned — the so-called insoluble 
 salts. They found that a liter of water at 15° dissolved o. i mg. of silver iodide ; 0.4 mg. 
 of silver bromide ; 0.5 mg. of mercuric iodide ; 1.7 mg. of silver chloride ; 3.1 mg. of mer- 
 curous chloride ; 2.6 mg. of barium sulphate ; 107 mg. of strontium sulphate ; 2070 mg. 
 of calcium sulphate ; 13 mg. of calcium carbonate ; 24 mg. of barium carbonate ; 3.8 mg. 
 of barfum chromate; 0.2 mg. of lead chromate; 9 mg. of magnesium hydroxide. See 
 Z. f. anorg. Ch. 5 (1894), 237. 
 
 The flame colorations produced by the volatile compounds are very 
 characteristic; calcium salts impart a reddish-yellow color; strontium, 
 an intense crimson ; barium, a yellowish-green. The spectrum of calcium 
 exhibits several yellow and orange lines, and in addition, a green and 
 a violet line (see the Spectrum Table); that of strontium contains, 
 besides several red lines, an orange and a blue, which are less distinct 
 but very characteristic. Finally, the barium spectrum consists of several 
 orange, yellow, and green lines, of which a bright green is particular! \ 
 prominent. 
 
 DIAMMONIUM COMPOUNDS. 
 
 Tlie .same relation which ammonia sustains to the univalent alkali metals (p. 296), 
 hvtirazine or dianiide, N^II^ (pp. 125, 131 ), discovered by Curtins, bears to the bivalent 
 alkaline earth metals. Diamide, however, does not always act as a di-acid base, which 
 might be inferred from what has been said, for in its most stable derivatives it is as mon- 
 acid as ammonia. It may be assumed that the di-acid compounds contain the bivalent 
 II 
 
 radical diainnioniiiin, 1^2* counterpart of the radical ammonium, and that they 
 
DIAMMONIUM COMPOUNDS. 
 
 3II 
 
 accordingly arrange themselves with the salts of barium, strontium and calcium. The 
 similarity of diammonium to the alkaline earth metals is shown also in the difficult solu- 
 bility of the sulphate and by its inability to form alums. 
 
 Again, diamide manifests properties allying it with the alkali metals. Its hydrate, 
 N.^H^ . 11.^0 = N.^IIgOIl, generally behaves like a mon-acid base. Its chloride, N2HgCl,^, 
 breaks down below 100° into hydrochloric acid and the chloride N.^HgCl, which does not 
 yield any more hydrochloric acid unless destroyed by heating. The hydrate N2H4 . - 
 2H2O = N2Hg(OH)2 is only stable in aqueous solution and upon evaporation changes to 
 N2ll4. H2O = N2II5OM, boiling without decomposition. It forms but one nitrate of the 
 composition N2H4.HN0.4, but one sulphocyanide N2II4.HCNS, and but one diam- 
 monium nitride, N2H4 . N3H. 
 
 It would therefore appear that both the bivalent radical N2Hg and the univalent radical 
 NjHg are present in the compounds of diamide, and further that those containing the 
 radical N2H5 are the more stable derivatives. 
 
 
 f OH 
 
 Dihydrate. 
 
 N2H5OH 
 
 Mono-hydrate. 
 
 Unstable Diammonium Compounds. 
 
 Dichloride. Di-iodide. 
 
 Stable Diammonium Compounds. 
 
 N2H5CI N2H5I 
 
 Mono-chloride. Mono-iodide. 
 
 N.HgNOg. 
 
 Nitrate. 
 
 N 2 HgS 04 . 
 
 Sulphate. 
 
 (N 2 Hg) 2 S 04 
 
 Semi-sulphate. 
 
 The hydrazine .salts generally crystallize well. Their great reducing power is one 
 of their characteristics. They precipitate metallic silver from ammoniacal silver solutions ; 
 cuprous oxide and metallic copper from alkaline copper solutions (Fehling’s solution), and 
 gold from acid solutions of gold chloride (distinction from hydroxylamine). The hydra- 
 zine salts are very poisonous. With care nitrous acid will liberate hydrazoic acid from 
 them : 
 
 N2H4 + HNO2 = N3H -f 2H2O, 
 
 otherwise nitrogen. The addition of copper sulphate to the solutions of hydrazine salts 
 causes the separation of a sparingly soluble double salt CUSO4 . (N2H-).2S04 ; this can be 
 used to separate hydrazine from its solutions and from mixtures. The corresponding 
 double salt.s, with the salts of nickel, cobalt, iron, manganese, cadmium and zinc, are also 
 soluble with difficulty and are anhydrous. P'or the detection by means of benzaldehyde 
 (benzalazine, m. p. 93°) see Organic Chemistry. Consult Curtius and Schrader, Jr. 
 prakt. Ch. 50 (1894), 31 1. 
 
 II 
 
 Diammonium Sulphate, N2II4.II2SO4 = (N2Hg)S04, can be made by treating 
 solutions of all the other hydrazine salts with sulphuric acid. It consist of vitreous plates, 
 which are soluble in about 33 parts of water at 20°. They deflagrate when quickly heated. 
 
 II 
 
 Diammonium Chloride, N2II4. 2IICI = (N2Hg)Cl2, obtained by transposing the 
 sulphate with barium chloride, forms large, vitreous octahedra. It is very soluble in 
 water. It melts with decomposition at about 200°. At 100° it passes gradually into 
 monohydrate. 
 
 I 
 
 Hydrazine Hydrate, (diammonium monohydrate), N2II4. = (N2ll5)OII, can 
 be obtained by distilling hydrazine sulphate with caustic potash. It is a strongly refract- 
 ing, not very mobile liquid, which fumes in the air and boils at 118.5° under 739.5 mm. 
 pres.sure. It has a faint odor, not reminding one of ammonia, has a caustic taste, is 
 hygro-scojjic and attracts carbon dioxide from the air. Below — 40° it solidifies to a 
 leafy crystalline mass. Its specific gravity equals 1.03 at 21°. It sinks in water and 
 only mixes with it after some time. It can be preserved in a pure condition and in con- 
 centrated solution. Dilute aqueous .solutions decompose completely in (he course of 
 time. Its .solution and its vapors color red litmus blue. It breaks dowm (from its vapor 
 density) at 171° and 760 mm. jire.ssure into hydrazine and water, but on the other hand 
 it vaporizes in a vacuum at 100° without decomposition. It has a powerful reducing 
 
312 
 
 INORGANIC CHKMISTKY. 
 
 action upon many oxides ; with cliromic acid and mercuric oxide it causes explosion. 
 Nothing delinite is known as to its chemical constitution. [See Curtins and .Schul/, Jr. 
 j)rakt. Ch. 42 (1890), 521 ; Curtins and Schrader, ibid. 50 (1894), 318J. Upon distill- 
 ing the hydrate with barium oxide under reduced pressure the diamide boiling only at 
 lower, is set free. It fumes strongly in the air and is combustible. 
 
 I 
 
 Diammonium Monochloride, NjH^. HCl _ (N.^Il5)Cl, results when the dichloride 
 is heated. It crystallizes in long white needles, melting at 89°. It is very soluble in 
 water. 
 
 I 
 
 Diammonium Semisulphate, ( N^I I^).^ . I I.^SD^ — (N^n.),^SO,, is formed ui)on 
 neutralizing hydrazine hydrate with sulphuric acid. It crystallizes in large, delicjuesccnt 
 plates, melting at 85°. 
 
 Another reTnarkable body is the Diammonium Nitride, N5II5 — N,^lI^.N3ll = 
 
 (N2ll5)N.j, prepared by Curtins from ammonium nitride and hydrazine hydrate. It is the 
 
 II 
 
 ammonium salt of hydrazoic acid. 'I'he other salt, Ngl 1^, = K^If^ . fN^II).^ = (N^I I,.)(N.j)2, 
 corresponding to the dichloride described above, could not be obtained. Diaminonium 
 nitride (diaminonium monazide, if the salts of hydrazoic acid are designated azides) con- 
 sists of large, vitreous prisms, melting at about 50°, deliijuescing in the air and volatiliz- 
 ing at the ordinary temperature, but more readily with acpieous and alcoholic vapors, 
 d'he salt dissolves with difficulty in alcohol. When ignited it burns cpiietly with a yellow 
 flame, without smoke and without residue. If heated rapidly in the air, by contact with 
 a wire raised to a white heat by ignition with detonating metallic azides or fulminating 
 mercury, a powerful explosion will take place even if the salt has deliquesced [Ber. 24 
 (1891), 3348 ; see also Ammonium Nitride, p. 298.] 
 
 2. METALS OF THE MAGNESIUM GROUP. 
 
 Id' this group are usually included beryllium, magnesium, zinc, and 
 cadmium. However, these metals do not exhibit complete analogy, as is 
 clearly seen from their position in the periodic system (p. 300). Beryl- 
 lium shows the greatest variations. It approaches aluminium, while mag- 
 nesium resembles not only zinc and cadmium, but also the alkaline earth 
 metals, calcium, strontium, and barium. Its similarity to the latter is 
 expressed by the basic nature of its oxide, whereas it resembles zinc and 
 cadmium mainly in isomorphism of compounds. 
 
 Beryllium and magnesium bear the same relations to calcium, stron- 
 tium, and barium, as lithium and sodium bear to the metals of the potas- 
 sium grou]:). 
 
 The alkaline character of the alkaline earth metals gradually diminishes 
 from barium to calcium, and becomes almost nothing in magnesium and 
 beryllium, which possess the lowest atomic weights (see p. 300). Mag- 
 nesium and beryllium are scarcely capable of decomposing water, even at 
 l)oiling temperatures. Their oxides and hydroxides are almost insoluble 
 in it ; the hydroxides decompose, on gentle ignition, into oxides and 
 water. 'I'lieir carbonates are very unstable; their chlorides, too, suffer a 
 partial decom])osition into oxide and liN’drogen chloride, even on drying, 
 d'lie solubility of the sulphates of magnesium and beryllium further dis- 
 tinguishes them from the metals of the alkaline earth group. 
 
MAGNESIUM. 
 
 3^3 
 
 The specific properties of beryllium and magnesium are maintained in 
 zinc and cadmium, which constitute a natural group with the former. 
 Zinc and cadmium do not decompose water at a boiling heat; their 
 hydroxides are insoluble in it, and are not very stable; their carbonates 
 and chlorides easily undergo decomposition; their sulphates are readily 
 soluble in water. The similarity is further expressed by the isomorphism 
 of most of their compounds. 'I'hus, magnesium and zinc sulphates'crys- 
 tallize with seven molecules of water, in perfectly similar forms. If the 
 solution of a mixture of both salts be allowed to crystallize, we get crystals 
 with variable quantities of zinc and magnesium : the formation of such 
 tsomorphous mixtures in ad libitu 7 n proportions is a characteristic indica- 
 tion of the isomorphism of compounds chemically similar (p. 255). 
 
 The difference between beryllium and magnesium upon the one side, 
 and zinc and cadmium on the other, is shown distinctly in their specific 
 gravities. While the first two elements possess a low specific gravity 
 (Be 1.8, Mg 1.74), zinc and cadmium (with specific gravities of 7.1 
 and 8.6) belong to the so-called heavy metals (see p. 252). 
 
 The difference in specific gravity determines, also, many differences in 
 chemical character. The light metals (especially the alkalies and alkaline 
 earths) form rather unstable sulphides, readily soluble in water, while the 
 sulphides of zinc and cadmium, like those of all heavy metals, are insolu- 
 ble in water, and usually in dilute acids; hence magnesium and beryl- 
 lium are not, while zinc and cadmium are, precipitated by hydrogen sul- 
 phide or alkaline sulphides as sulphides from solutions of their salts. 
 Further, the oxides of the light metals are very stable, and are only 
 reduced by ignition with carbon if they are readily fusible (like potassium 
 and sodium oxides); the heavy metals, on the other hand, are easily separ- 
 ated from their oxides by carbon. Zinc and cadmium oxides are reduced 
 by carbon, while those of magnesium and beryllium are not altered. 
 
 1. MAGNESIUM. 
 
 Mg = 24.36. 
 
 Magnesium is very abundant in nature, and almost always accompanies 
 calcium in its compounds. As carbonate, it occurs in compact masses, 
 as magnesite, etc. Dolomite, which forms entire mountains, is an iso- 
 morphous mixture of calcium and magnesium carbonates. Magnesium 
 is also present in most of the natural silicates; its soluble salts are con- 
 tained in almost all natural waters. In conjunction with the alkali salts 
 they constitute the most important salts of the Stassfurt beds (see p. 
 276). 
 
 Metallic magnesium may be obtained by the electrolysis of the chloride 
 (Bunsen, 1852) or by heating the same with sodium (Bussy, 1830). It is 
 more easily prepared by heating the double chloride of magnesium and 
 sodium with metallic sodium and fluorspar, the latter serving merely as a 
 flux : 
 
 MgCb.XaCl -f- 2Na = 3NaCl 4 Mg. 
 
 27 
 
3^4 
 
 INORGANIC CHKMlh'I'RV. 
 
 At present magnesium is ol)tained in large (luantities by the electrol- 
 ysis of fused carnallite [see Dammer, Chemische 'Fechnologie, Ikl. ii, 
 
 (1895)]- 
 
 Magnesium is a brilliant, almost silver-white metal, of specific gravity 
 1.74. It is tenacious and ductile, and when heated may be converted 
 into wire and rolled out into thin ribbons. It melts at about 700°, and 
 distils at a bright-red heat. At ordinary temperatures, it scarcely oxidizes 
 in the air; it burns, wlien heated, with an extremely intense white light, 
 owing to the glowing non-volatile magnesium oxide. Magnesium light 
 is rich in chemically active rays, and, for this reason, it is employed in 
 photograi)hy for artificial illumination. Its alloy with zinc is generally 
 em})loycd as a substitute for pure magnesium, as it burns with an 
 equally bright light. Its intense light has led to its use in pyrotechny. 
 Boiling water is very slowly decom])osed by magnesium. It dissolves 
 easily in dilute acids, forming salts; the alkalies do not attack it. Many 
 metallic oxides and acids when heated with magnesium powder lose their 
 oxygen and are reduced by it (see p. 237). 
 
 Magnesium Oxide, MgO, or ;/n 7 ^i^;/esia, formed by the combustion 
 of magnesium, is ordinarily obtained by the ignition of the hydroxide or 
 the carbonate (^Magnesia iista). It is a white, very voluminous, amor- 
 phous powder, which finds apjdication in medicine, d'he feebly ignited 
 magnesia combines with water, with slight generation of heat, to produce 
 magnesium hydroxide. 
 
 In Stassfurt the magnesium chloride liciuors, produced in the prepara- 
 tion of potassium chloride, are utilized in j^roducing a very good mag- 
 nesia, which is used in the manufacture of fire-brick or fire-clay (p. 276). 
 
 At high temperatures magnesia conducts electricity and it can be made to glow and 
 become luminous by the current. If, therefore, little thin cylinders of magnesia are 
 heated until they reach the point of conductivity they will answer for the production of 
 the electric light (Nernst). 
 
 Magnesium Hydroxide, Mg(OH)2, is precipitated from solutions 
 of magnesium salts by potassium or sodium hydroxide as a gelatinous 
 mass. Dried at 100° it is a white amorphous powder. It is almost in- 
 soluble in water and alkalies; moist red litmus-paper is, however, colored 
 blue. Ammonium salts dissolve it quite easily, forming soluble double 
 salts. Magnesium hydroxide attracts carbon dioxide from the air and 
 forms magnesium carbonate. It yields the oxide and water when gently 
 ignited. 
 
 Magnesium Chloride, MgCl2, is present in traces in many mineral 
 s])rings. It may be prepared by the solution of the carbonate or oxide 
 in hydrochloric acid ; in large quantities it is obtained as a by-])roduct 
 in the technical ])roduction of potassium chloride. When its solution 
 is evaporated the salt crystallizes out with six molecules of water in deli- 
 quescent crystals, isomori)hous with calcium chloride. When these are 
 heated they give up water, and there occurs at the same time a partial 
 decomj)osition of the chloride into oxide and hydrogen chloride: 
 
 MgC:i2 1 11,0 MgO 4- 2IIC1. 
 
MAGNESIUM SULPHATE. 
 
 315 
 
 As magnesium chloride is produced in large quantities in various technical processes, 
 repeated efforts liave been made to utilize the above reaction for the preparation of hydro- 
 chloric acid, by conducting steam over heated magnesium chloride. 
 
 Lately the liquors have been evaporated and the magnesium chloride dehydrated, 
 whereby it breaks down in part into hydrochloric acid and magnesium oxide. When the 
 residual magnesium oxychloride is heated in an air current, magnesia and chlorine are 
 produced (Weldon-Pechiney ; see p. 50). 
 
 To get anhydrous magnesium chloride, ammonium chloride is added to 
 the solution of the former. The double salt, MgCl2.NH^Cl -|- 6H2O, 
 is formed. When this is heated it first loses water, and at 460° throws 
 off ammonium chloride, leaving anhydrous magnesium chloride. The 
 latter can also be obtained by heating the hydrated salt in a current of 
 hydrochloric acid. This is a leafy, crystalline mass, which fuses at 708°, 
 and distils undecomposed at a red heat; it is very deliquescent in the air. 
 
 Double salts, similar to the above, are also formed from potassium and 
 calcium chloride. The potassium double salt, MgCl2.KCl -f- 6H2O, 
 occurs in considerable deposits as carnallite at Stassfurt (p. 276). 
 
 Magnesium Sulphate, MgSO^, is found in sea-water and in many 
 mineral springs. With more or less water it is kieserite^ which abounds 
 extensively at Stassfurt. At ordinary temperatures it crystallizes with 
 seven molecules of water, MgSO^ -j- 7H2O, in four-sided rhombic prisms, 
 readily soluble in water (at 0° in 2 parts of w’ater). It has a bitter, salt-like 
 taste, and serves as an aperient. It crystallizes with six molecules of water 
 'from solutions heated to 70°; at 0°, however, it has twelve molecules. 
 When heated to 150° these hydrates lose all their water of crystallization, 
 excepting one molecule, which escapes above 200°. One molecule of 
 water, in magnesium sulphate, is, therefore, more closely combined than 
 the rest. Many other salts containing water deport themselves similarly. 
 The more intimately combined water is termed water of constitution. 
 
 Magnesium sulphate forms double salts with potassium and ammonium 
 sulphates, which crystallize with six molecules of water in monoclinic 
 prisms, e. g : 
 
 MgSO,.K2SO, 4- 6H2O. 
 
 The sulphates of zinc and several other metals, e. g., iron, cobalt, and nickel, in their 
 bivalent forms, are very similar to magnesium sulphate. Their sulphates crystallize with 
 seven molecules of water, and are isomorphous. They form double salts with potassium 
 and ammonium sulphates ; these crystallize with six molecules of water, and are iso- 
 morphous, e. g. : 
 
 ZnSO, + 7H2O ZnSO, . + 6H2O. 
 
 FeSO^ -b 7H2O FeSO^. K2SCh -f 6 Uf. 
 
 The constitution of these double salts may be viewed in the same way as that of potas- 
 sium-sodium sulphate, or of mixed salts of polybasic acids. We may suppose that in the 
 given instance the bivalent metal unites two molecules of sulphuric acid : 
 
 II 
 
 ^04/Mg -f 6 H 20 . 
 
 Magnesium Phosphates. — The tertiary phosphate Mg,(POJ2> 
 accompanies the tertiary calcium phosphate in small quantities in bones 
 
INORGANIC CHEMISTRY. 
 
 316 
 
 and in plant ashes. The secondary phosphate, MgHPO^ -|- is 
 
 precipitated from the soluble magnesium salts, by disodium phosphate 
 (NajHPO^) as a salt dissolving with difficulty in water, in the pres- 
 ence of ammonia and ammonium salts, magnesium-ammonium phosphate , 
 MgNH^PO^ -f- ^H/), is precipitated as a crystalline powder insoluble in 
 water. This double salt is found in guano, forms in the decay of urine, 
 and is sometimes the cause of the formation of calculi. The primary 
 salt, MgHj(PO^)2, has not been obtained. 
 
 The magnesium salts of arsenic acid, H^AsO^, are very similar to those 
 of phosphoric acid. Magnesium-ammonium arseniate, MgNH^AsO^ -f- 
 6H2O, is likewise almost insoluble in water. 
 
 Magnesium Carbonate, MgCO,, occurs in nature as magnesite, 
 crystallized in rhombohedra ( isomorphous with calcite), or more often 
 in compact masses. Combined with calcium carbonate, it forms dolo- 
 mite, to which, when pure, is ascribed the formula, CaCO^.MgCOg; 
 however, it usually contains an excess of calcium carbonate. On adding 
 sodium or potassium carbonate to the acpieous solution of a magnesium 
 salt, some carbon dioxide escapes, and a white precipitate forms, which 
 consists of a mixture of magnesium carbonate and hydroxide. If the 
 precipitate be dried at low temperature, we obtain a white, voluminous 
 powder, whose composition generally corresponds to the formula 
 Mg(OH)2. 3MgC03 -f- 4H,^0. 'Phis salt is em[)loyed under the name 
 Magnesia alba in medicine. If it be suspended in water, and carbon 
 dioxide passed through it, the salt will dissolve, and upon standing 
 exposed to the air, crystals of neutral cai'bonate, MgCOg 3^2^? sepa- 
 rate. When these are boiled with water they give up carbon dioxide 
 and are again converted into the basic carbonate. The naturally occur- 
 ring magnesite sustains no change when boiled with water, and it is 
 only when it is heated above 300° that it decomposes into magnesium 
 oxide and carbon dioxide. 
 
 Freshly prepared magnesium carbonate dissolves in acid alkaline car- 
 bonates and crystallizes from them as 4MgC03 -{- 15H2O. 
 
 Magnesium carbonate yields isomorphous double salts, with potassium 
 carbonate and ammonium carbonate ; e. g., MgC03.K2C03 -f- 4H2O. 
 
 Of the silicates of magnesium, we may mention olivine (Mg2SiO^), 
 serpentine (Mg3Si207 -f" 2H2O), talc (Mg^Si^O^j -f- H2O), sepiolite or 
 meerschaum (Mg2Si30g -f- 2H2O). The mixed silicates of magnesium 
 and calcium are very numerous; to these belong asbestos, the augites and 
 horn-blendes. 
 
 Magnesium Nitride, Mg.^Nj (pp. 116, 124, 126), is a light, porous, yellow-colored 
 mass. liriegleb and (ieuther found that it was produced when magnesium was heated 
 in a current of ammonia : 
 
 Mg3 t 2NIl3 = Mg3N, -f 3lb. 
 
 It is also obtained pure when magnesium is heated in nitrogen. When introduced into 
 water, the latter is made to boil and large volumes of ammonia escape [see Pasch- 
 kowezky, Jr. prakt. Ch, 47 (1893), 89]. 
 
BERYLLIUM. 
 
 317 
 
 Recognition of Magnesium Compounds. — The fixed alkaline hy- 
 droxides precipitate magnesium hydroxide from soluble magnesium salts; 
 the carbonates throw down basic magnesium carbonate. The precipitates 
 are insoluble in pure water and the alkalies, but dissolve readily in solu- 
 tions of ammonium salts. In the presence of the latter, neither the alka- 
 line hydroxides nor carbonates cause precipitation. In the presence of 
 ammonia and ammonium chloride, disodium phosphate precipitates mag- 
 nesium-ammonium phosphate, MgNH^PO^ insoluble in water. 
 
 2. BERYLLIUM. 
 
 Be = 9.1. 
 
 Among the metals of the second group beryllium occupies a position similar to that of 
 lithium in the first group (pp. 272, 299) ; in both elements, which have the lowest atomic 
 weight in their group, the specific group character is considerably diminished, or does not 
 find expression. As lithium attaches itself in many respects to magnesium, so does beryl- 
 lium approach aluminium. Like the latter, it is scarcely at all attacked by nitric acid, but 
 dissolves easily in sodium or potassium hydroxide, with elimination of hydrogen. Like 
 aluminium oxide, that of beryllium dissolves in the alkalies, and is almost invariably 
 accompanied by the former in its natural compounds. However, beryllium, in most of 
 its compounds, stands nearer to magnesium than to aluminium. The determination of 
 the vapor density of beryllium chloride (see below) and of certain organic derivatives has 
 finally established the atomic weight and the valence of this element. 
 
 Beryllium is not very abundant in nature and is found principally in beryl, a double 
 silicate of aluminium and beryllium, Al2Be3(Si03)(,. Emerald has the same composi- 
 tion, and is only colored green by a slight amount of chromium oxide. 
 
 It was in these minerals that V auquelin discovered beryllium in 1797. It is also present 
 in leucophane (a silicate), together with aluminium, fluorine and sodium ; and in gado- 
 linite (a silicate) with ferrous oxide, yttrium, cerium, lanthanum, and other rare earths. 
 
 Metallic beryllium is obtained by the ignition of the chloride, or better, potassium beryl- 
 lium chloride (or potassium beryllium fluoride, BeFl2.2KFl), with sodium, and by the 
 electrolysis of BeFlg. KFl. It is a white, ductile metal, of specific gravity 1.8. Its specific 
 heat at the ordinary temperature equals 0.408 ; the atomic heat is, therefore, 3.7 
 (p. 254). It does not decompose water, even upon boiling. It does not oxidize in the 
 air at ordinary temperatures. When finely divided and heated it will burn in the air 
 with a very bright light. It is readily dissolved by dilute hydrochloric and sulphuric 
 acids ; also by potassium and sodium hydroxides. 
 
 Beryllium Chloride, BeCl2, is obtained, like aluminium chloride, by the ignition of 
 a mixture of beryllium oxide and carbon in a stream of chlorine. It sublimes in shin- 
 ing needles, which deliquesce in the air. Its vapor density corresponds to the molecular 
 formula BeCl2 == 79.9 (Nilson). Pure beryllium chloride dissolves in water, forming a 
 colorless solution, from which it crystallizes with four molecules of water; upon drying it 
 suffers a decomposition similar to that of magnesium chloride. 
 
 The salts of beryllium have a sweet taste, hence it has been called glticinwn. 
 Ammonium hydroxide precipitates a white, gelatinous beryllium hydroxide, Be(OH)2, 
 from solutions of the soluble salts. This dissolves readily in sodium and potassium 
 hydroxide, and in a mixture of ammonia and ammonium carbonate, but on boiling, sepa- 
 rates again from solution. When heated, the hydroxide breaks down into water and 
 beryllium oxide, Bef), which is a white, amorjflious powder, of specific gravity 2.96. 
 
 Beryllium Sulphate, BeSO^, crystallizes from water at various temperatures, with 
 four or seven molecules of water ; it contains two molecules at 105°, and is anhydrous at 
 
INORGANIC CIIKMIS'I'RY. 
 
 318 
 
 250-260°, It docs not crystallize with inagiicsiuin siilpliatc in an isoinorphous niixlnrc. 
 d'he double salt, HeSO^. K.^St >4 -)- 2ll2<), docs not dissolve readily in water. Compare 
 Kruss and Moraht, Her. 23 (1890), 727, and Ann. Chein. 262 (1891), 38. 
 
 ;j. ZINC. 
 
 Zn 65.4. 
 
 The natural compounds of the heavy metals liave generally a high spe- 
 cific gravity, frequently jiossess metallic luster, usually occur in the older 
 crystalline rocks in veins, and are termed ores. The most important zinc 
 ores are the carbonate, ZnCO,, the silicate, and sphalerite or blende, ZnS. 
 The iirincipal sources of these ores are in vSilesia, hingland, Belgium, 
 Poland, the United ^States, Prance, and Spain. To get the metal the 
 carbonate or sulphide is converted into oxide by roasting it in the air; the 
 product is then mixed with carbon and ignited in cylindrical clay tubes. 
 In this manner the oxide is reduced : 
 
 ZnO -L C — Zn CO, 
 
 and the liberated zinc distilled off. The receivers contain the fused, 
 compact zinc and a gray, pulverulent mass, called zinc-dust, which con- 
 sists of a mixture of zinc oxide with finely divided metal. This material 
 is used in laboratories as a strong reducing agent. See Mylius and Fromm 
 on the purification of zinc, Z. f. anorg. Ch. g (1895), 144. 
 
 Metallic zinc has a bluish-white color, and exhibits rough, crystalline 
 fracture; its specific gravity equals 7-7. i . At ordinary temperatures it 
 is brittle and can be pulverized; at 100-150° it is malleable and can be 
 rolled into thin leaves and drawn out into wire. At 200° it becomes 
 brittle again and may be easily broken. It melts at 420° and distils 
 at about 950°. 
 
 It becomes coated with a thin layer of basic carbonate in moist air. 
 Heated in the air it burns to zinc oxide with a very intense, bluish-white 
 light. Compact zinc will only decompose water at a red heat ; zinc-dust, 
 however, acts at ordinary temperatures. Perfectly pure zinc is only 
 slowly attacked by dilute acids at the ordinary temperature; it dissolves 
 by boiling with potassium or sodium hydroxide, as well as ammonia, 
 with liberation of hydrogen : 
 
 Zn 4- 2NaOII = Zn(ONa)2 + Hj. 
 
 Owing to its slight alteration in the air zinc meets with extensive appli- 
 cation as sheet-zinc for coating statues and in architectural adornment, 
 and in galvanizing sheet-iron. It also forms an important constituent 
 of many valuable alloys, such as brass and argentan (see these). 
 
 Zinc Hydroxide, ZnfOH)^, is preci])itated as a white, amorphous 
 powder, from acpieous solutions of zinc salts, by alkalies, and is soluble in 
 excess of the reagent. When heated it deconq)oses into water and zinc 
 oxide. 
 
 Zinc Oxide, ZnO, is usuallv jirepared by igniting the precipitated 
 basic carbonate, and, as zinc iviiile, is employed as a stable white ])aint. 
 'I'he oxide obtained by burning the metal is a white, voluminous, floccu- 
 
2INC. 3^9 
 
 lent mass, called Flores zi/ici, or Lana philosopJiica. When zinc oxide is 
 heated it acquires a yellow color, which disappears on cooling. 
 
 Zinc oxide occurs in nature as zincite^ colored by impurities. 
 
 Zinc Chloride, ZnCl2 {Zincutn chloratiini), anhydrous, is obtained by 
 heating zinc in a stream of chlorine, by the evaporation of the solution 
 of zinc in hydrochloric acid, and by the distillation of zinc sulphate with 
 calcium chloride. It forms a white, deliquescent mass, which fuses when 
 heated and distils at about 730°. Molten zinc chloride is a transparent, 
 very mobile, strongly refracting liquid, attracting moisture more readily 
 than phosphoric anhydride [Lorenz, Z. f. anorg. Ch. 10 (1895), 82]. It 
 crystallizes from concentrated hydrochloric acid solution with one mole- 
 cule of water in deliquescent octahedra. When the aqueous solution of 
 zinc chloride is evaporated it partly decomposes (like magnesium chlo- 
 ride) into zinc oxide and hydrochloric acid. When the concentrated 
 zinc chloride is mixed with zinc oxide, a plastic mass is obtained, which 
 hardens rapidly; a mixture of magnesium chloride and oxide does the 
 same. In both instances the hardening depends upon the formation of 
 
 Cl . . 
 
 basic oxychlorides, e. g., Zn<Qp^ Zinc chloride forms deliquescent 
 
 double salts with the alkaline chlorides, e. g., ZnCl2.2KCl. With 
 ammonia it yields various compounds, of which ZnCl2.NH3 is character- 
 ized by great stability. 
 
 Zinc Sulphate, ZnSO^, is obtained by dissolving zinc in sulphuric 
 acid. It is prepared upon a large scale by carefully roasting zinc- 
 blende (ZnS) ; the zinc sulphate is extracted by water. It crystallizes 
 at ordinary temperatures from aqueous solutions with seven molecules of 
 water-(zinc or white vitriol, Zincum sulphuricum') in rhombic crystals, re- 
 sembling those of magnesium sulphate. It affords double salts with the 
 alkaline sulphates; these contain six molecules of water (p. 315). 
 
 Zinc Carbonate, ZnCOg, occurs native as smithsonite in hexagonal 
 crystals, isomorphous with those of calcite. Sodium carbonate precipi- 
 tates a mixture of ZnCOg. H2O, which rapidly becomes crystalline, and 
 the basic carbonate 2ZnC03. 3Zn(OH)2. H2O. 
 
 Zinc Sulphide, ZnS, is zinc-blende or sphalerite, usually colored 
 brown by ferric oxide or other admixtures. Ammonium sulphide precipi- 
 tates it as a white compound from zinc solutions. Although fused zinc 
 reacts with difficulty with sulphur, zinc-dust combines with the latter in 
 powdered form quite readily, and if the mixture be heated or struck with 
 a hammer the union is accompanied by an explosion. Zinc sulphide is 
 insoluble in water, but is readily dissolved by dilute acids, excepting acetic 
 acid; therefore it may be precipitated by hydrogen sulphide from zinc 
 acetate solutions. This reaction serves to separate zinc from other metals. 
 The sulphide is also applied as a white pigment ; especially a mixture of 
 it with barium sulphate (lithopone — made from zinc sulphate and barium 
 sulphide). 
 
 Zinc Silicate, Zn2SiO^-(- H2O, occurs in rhombic crystals as calamine. 
 In the analytical way zinc is characterized by the fact that its hydroxide 
 dissolves both in caustic soda and potash as well as in ammonia; hydro- 
 gen sulphide precipitates white zinc sulphide from these solutions. 
 
320 
 
 INORGANIC CHEMISTRY. 
 
 4. CADMIUM. 
 
 Ccl = 1 12. 
 
 Cadmium very often accompanies zinc in its ores. As much as 5 per 
 cent, of this metal is present in the .Silesian zinc ores; it was first dis- 
 covered in them in 181 7 by Hermann ; shortly after Stromeyer recognized 
 it as a new element. Jleing more volatile than zinc, in olitaining the 
 latter it distils off first, and may be easily separated from the first por- 
 tions of the distillate. It is a white, tenacious, and rather soft metal, of 
 specific gravity 8.6. It fuses at 320° and l)oils at 770°. It does not alter 
 much in the air. Heated, it burns with the separation of a brown smoke 
 of cadmium oxide. It dissolves with difficulty in dilute hydrochloric and 
 sulphuric acids, but readily in nitric acid. Zinc throws out the metal 
 from solutions of the soluble cadmium salts. 
 
 St. Claire Deville found the specific gravity of cadmium vapors at 1040° and II. 
 Biltz at about 1700° to approximate closely 112 (0.^;=r32). 'i’herefore, the molecular 
 weight of cadmium is 112. Since the atomic weight of cadmium (determined from its 
 specific heat) is also 1 12, it follows that the gas molecule of cadmium consists of but 
 atom. We know that the molecules of other elements in the gaseous state are composed 
 of two or more atoms (O.^, N2, P4). Cadmium, therefore, forms an exception to this rule. 
 This is also true of mercury, zinc and perhaps, too, of other bivalent metals. The.se 
 relations remind us of the behavior of the hydrocarbon residues (radicals) ; while the 
 bivalent or quadrivalent groups, <?. , ethylene. C.^H^, and acetylene, exist in free 
 
 condition, the univalent groups (as CH, and CNj cannot appear free, but double them- 
 selves when separated from their compounds. 
 
 Cadmium Hydroxide, Cd(0H)2, is precipitated as a white powder, 
 from the soluble cadmium salts, by the alkalies; it is insoluble in sodium 
 and potassium hydroxides, but dissolves readily in ammonium hydroxide. 
 
 Cadmium Oxide, CdO, is prepared by igniting the nitrate. It is a 
 brownish-black ])owder, consisting of microscopic octahedra. When ob- 
 tained by heating the carbonate or hydroxide, it is a brown, amorphous 
 powder. 
 
 Cadmous Hydroxide CdOH, and its oxide, Cd^O, have been prepared. The first 
 is a grayish- white compound, while the second is composed of yellow, translucent crystals. 
 It is obtained on heating the hydroxide to a temperature at which sulphuric acid gives 
 off dense white fumes. The hydroxide is a reducing agent, yielding hydrogen with 
 hydrochloric acid and oxides of nitrogen with nitric acid. The oxide conducts itself 
 similarly. These compounds are of interest, as they foreshadow the tendency toward 
 the formation of lower oxides, so strongly shown by mercury {^Am. Cheni. Jr., 12, 493 . 
 
 Cadmium Chloride, CdCl2, crystallizes from aqueous solution, with 
 two molecules of water, and may be dried without decomposition. The 
 anhydrous salt melts at 540° and sublimes in scales. 
 
 Cadmium Iodide, Cdl2, is obtained by the direct action of iodine 
 upon metallic cadmium in the j)resence of water. It crystallizes from 
 the latter in hexagonal tal)les. It is used in photography. 
 
 Cadmium Sulphate, CdSOj, crystallizes at — 20°, like the sulphates 
 of zinc and magnesium at the ordinary temperature, with seven mole- 
 cules of water, but at the ordinary temj)erature with -I molecules of 
 water ; the crystals effloresce in the air. It, however, forms double salts 
 
CADMIUM. 
 
 321 
 
 with the sulphates of the alkali metals, e. g., CdSO^ . K.^SO^ -f dH/) ; 
 these are perfectly analogous to thv/Se of zinc and magnesium, and iso- 
 morphous with them (p. 319). 
 
 Cadmium Sulphide, CdS, occurs native as greenockite, in yellow 
 hexagonal prisms. Hydrogen sulphide precijutates it from cadmium salt 
 solutions as a yedow powder, insoluble in dilute acids. It is employed 
 as a pigment. 
 
 Almost all the alloys of cadmium have a low melting point. Freshly 
 prepared cadmium amalgam is a white ])lastic mass, which soon becomes 
 hard ; it was formerly used for filling teeth. 
 
 The magnitude of the atomic weight of Mercury would place the lat- 
 ter in the group of zinc and cadmium. The relationship of these three 
 heavy metals occupying a similar position in the three great periods are, 
 from a physical point of view, distinguished among the heterologous mem- 
 bers by their ready fusibility and volatility, which nearly reach a maxi- 
 mum in them. In the homologous series, zinc, cadmium, mercury, these 
 properties increase with rising atomic weight (just as with the metals of 
 the potassium group, p. 272) : 
 
 
 Zn 
 
 Cd 
 
 Hff 
 
 Atomic weight, 
 
 65-4 ' 
 
 I 12 
 
 200.3 
 
 Melting point, 
 
 420° 
 
 320° 
 
 - 39 - 5 ° 
 
 Boiling point, 
 
 950° 
 
 770° 
 
 357 ° 
 
 Specific gravity, 
 
 7.1 
 
 8.6 
 
 13-5 
 
 Atomic volume, 
 
 9.1 
 
 13 
 
 14.7 
 
 The gradation in the heat of formation of their compounds (p. 327) 
 clearly indicates that mercury must be arranged in a group with cad- 
 mium and zinc. 
 
 Like zinc and cadmium, it yields compounds of the form HgX^, in 
 which it a],)pears bivalent. These derivatives are, in many respects, simi- 
 lar to the corresponding compounds of zinc and cadmium. Thus, mer- 
 curic sulphate forms double salts with the alkaline sulphates, which crys- 
 tallize with six molecules of water (HgS04.K2S0^ -f 6H2O), and are iso- 
 morphous with the double sulphates of the magnesium group (p. 315). 
 The similarity, however, limits itself to few compounds. Since the 
 properties of each group sustain a slight change by virtue of the increas- 
 ing atomic weight, we are not surprised to observe this to be very evident 
 in the case of mercury (with the high atomic weight 200.3), especially as 
 the middle ftransition) member of the third great period is not known 
 (p. 245). Mercury differs essentially from zinc and cadmium in that, in 
 
 II 
 
 addition to the compounds of the form HgX2 (mercuric compounds), it 
 
 ' I 
 
 is also capable of yielding those of the form HgX (mercurous compounds^ 
 in which it seems to be univalent. Here we meet an instance, frequently 
 
322 
 
 in()R(;anic chemis'I'rn . 
 
 observed, in wliich one and the same metal (as witli tlic most metalloids) 
 is capable of forming compounds of two or more forms, which are to be 
 referred to a different valence of the metal ; and it often haj)i)ens that 
 the derivatives of a metal, ap[)earing in different forms or types, are fre- 
 quently more essentially distinguished from one another than the com- 
 pounds of different elements having the same tvpe. Thus, the mercuric 
 
 II 
 
 compounds (IlgX.J are similar to those of zinc and cadmium, after the 
 
 I 
 
 same form, while the mercurous com})ounds (HgX) exhibit great resem- 
 I I 
 
 blance to the cuprous (CuX) and silver (AgX) compounds, constituted • 
 according to a similar tyj)e. 
 
 It shows that the similarity of the compounds is influenced not only by 
 the nature of the metals, but frequently, to a marked degree, by the 
 forms or types according to which they are constituted (p, 330). 
 
 Heretofore it has been generally customary to assume the presence of two atoms of 
 mercury, united to a bivalent group, in the molecule of the mercurous derivatives ; hence 
 the appearance of a univalent character : 
 
 llg-Cl Hg-NOs 
 
 I'lg-Cl Hg-NO.,. 
 
 Chloride. Nitrate. 
 
 The most recent researches, however, show that the.se compounds can be represented 
 by the simple formulas HgCl, HgN03 ; see pp. 325, 332. 
 
 5. MERCURY. 
 
 Hg = 200.3. 
 
 Mercury {Hydrargyrum, i. e., water-silver') occurs in nature principally 
 as the sulphide, cinnabar ; more rarely native in the form of little drops 
 scattered through rocks. Its most important localities are Almaden, in 
 Spain; New Almaden, in California; Idria, in Austria; Tuscany, in 
 Italy; Mexico ; Peru; China; Japan. Considerable quantities have been 
 found also in Russia. 
 
 The metallurgical separation of mercury is very simple. Cinnabar is 
 roasted in reverberatory furnaces, whereby the sulphur burns to dioxide, 
 and the mercury vapors are condensed in large chambers with the aid of 
 water. There is a regular and gradual return to the old method of dis- 
 tilling cinnabar with lime or iron from iron retorts. Commercial mer- 
 cury usually contains a slight quantity of other metals dissolved in it. 
 h'or its purification, it is poured in a thin stream into a deep layer of dilute 
 nitric acid or of ferric chloride solution, or it is shaken with a solution 
 of i)otassium dichromate and dilute sulifliuric acid, by which the admixed 
 metals are more easily di.ssolved than the mercury. Finally, the metal is 
 passed through chamois skin and distilled (best in a vacuum). Ber. 12 
 (1879), 204, 576. 
 
 Mercury is the only metal which is liquid at ordinary temperatures. 
 At 0° its s|)ecific gravity equals 13.596 and at 15°, 13.559 ; it solidi- 
 fies at — 39-5°, and crystallizes in regular octahedra ; it evaporates some- 
 
MERCUROUS COMPOUNDS. 
 
 323 
 
 what at medium temperatures, and boils at 357°. Its vapors are very 
 poisonous. The specific gravity of the vai)or of mercury is 200.3, 
 (O2 = 32). Therefore, the molecular weight of the metal is 200.3, 
 its atomic weight is also 200.3, molecule, like that of cadmium and 
 zinc, is composed of only one atom. At ordinary temperatures, mercury 
 is not altered by exposure to the air ; near the boiling point, however, it 
 gradually oxidizes to red mercuric oxide. Hydrochloric and cold sulphuric 
 acids do not act upon mercury ; hot sulphuric acid converts it into mer- 
 curic sulphate, with evolution of sulphur dioxide. Even dilute nitric acid 
 will readily dissolve it. It combines with the halogens and sulphur at 
 ordinary temperatures. 
 
 Mercury dissolves almost all metals (not iron) forming amalgams. It 
 unites with potassium and sodium upon gentle warming, with production 
 of heat and light. When the quantity of potassium and sodium exceeds 
 2 per cent., the alloy is solid and crystalline; by less amount it remains 
 liquid. [Compare Z. f. phys. Ch. 29 (1899), 119.] Tin amalgam is 
 employed for coating mirrors. 
 
 Sodium amalgam may be made by placing rather large pieces of sodium upon the 
 bottom of an iron crucible, covering them with wire gauze, so that they cannot rise to the 
 surface, and pouring mercury upon them. With large quantities the reaction begins cf 
 itself. Finally, heat must be applied while stirring. 
 
 Mercury forms two series of compounds, mercurous and mercuric. The 
 first are analogous to the cuprous, and have the form HgX. In them 
 mercury appears to be univalent. In many respects the mercurous com- 
 pounds are similar to the cuprous and silver derivatives. The halogen 
 compounds are insoluble, and darken on exposure to light. 
 
 In the mercuric derivatives, HgX2, mercury is bivalent, and is very 
 much like zinc and cadmium. Mercuric compounds almost always form, 
 if the substance reacting with the mercury is in excess; when the oppo- 
 site is the case, mercurous salts result. Mercuric derivatives, by the 
 addition of mercury, pass into the mercurous, e. g., 
 
 Hg(NO,y+ Hg = Hg,(NO,),. 
 
 This change may also be effected by reducing agents (sulphurous acid, 
 phosphorous acid and stannous chloride). Oxidizing agents convert the 
 mercurous into mercuric compounds. 
 
 MERCUROUS COMPOUNDS. 
 
 Mercurous Chloride, HgCl or Hg2Cl2, calomel {Hydrargyrum chlor- 
 atum), is an amorphous white precipitate, produced by the addition of 
 hydrochloric acid or soluble chlorides to the solution of mercurous salts. 
 It is generally formed by the sublimation of mercuric chloride with mer- 
 cury ; or a mixture of mercuric sulphate, mercury and sodium chloride 
 is sublimed : 
 
 IlgSO, -f 2NaCl -f Hg = Na 2 SO, + 2HgCl. 
 
324 
 
 INORGANIC CHEMISTRY. 
 
 It then forms a radiating, crystalline mass (cjuadratic prisms) of sjiecific 
 gravity 7.2. Calomel is insoluble in water, in alcohol, and in dilute acids ; 
 it gradually decomi)oses when exi)osed to the light, with separation of 
 mercury. When heated, it sublimes without fusing, foiling hydro- 
 chloric acid and hot concentrated solutions of calcium chloride and 
 alkaline chlorides decompose it into mercuric chloride and free mercury: 
 
 2lIgC] = HgCl, + Ilg. 
 
 When ammonium hydroxide is j^oured over calomel, it blackens (hence 
 the name calomel, from xak(>/xeX(x', beautiful black); it is not surely 
 known what chemical change occurs heie. 
 
 The vapor density of calomel vapors at 400° (first determined by Mitscherlich, and con- 
 firmed by Deville and Troost, Rieth and (Idling) is 235 (< >2 = 32) ; which corresponds 
 to the formula HgCl. As formerly supposed, and as recently demonstrated by V. Meyer 
 and Harris in opposition to Fileti, calomel breaks down completely into mercury and 
 mercuric chloride [Her. 28 (1895), 364]. This mixture has the same vapor density that 
 undecomposed calomel would show : 
 
 iigci + Hgci = ng+ iigcb- 
 
 I vol. I vol. I vol. I vol. 
 
 The question, whether the mercurous compounds contain one or two atoms of mer- 
 cury, whether, for exam[)le, the formula Hg.^CI.^ or llgCl properly belongs to calomel, 
 can, therefore, not be decided by the determination of its vapor density alone. It must 
 also be ascertained of what the vapor consists — and this proof free from objection has not 
 yet been given. From electro-chemical experiments and from the analogy to cuprous 
 and silver chlorides it would seem very probable that mercurous chloride has the simple 
 formula HgCl [Z. f. anorg. C h. 9 (18951, 442]. 
 
 Mercurous Nitride, Ngllg or Hg.2(N3l2, is formed by adding mercurous nitrate to 
 solutions of hydrazoic acid or its alkali salts. It is insoluble in water. It consists of 
 microcrystalline needles, becoming yellow in the light and blackened by ammonia, just 
 like calomel. When heated or struck it explodes with great violence. Its separation 
 into its elements is accompanied with a brilliant blue light (p. 312). 
 
 Mercurous Iodide, Hgl or Hg2l2» is prepared by rubbing together 
 8 parts of mercury with 5 parts of iodine, or by precipitating mercurous 
 nitrate with potassium iodide. It is a yellowish-green powder, insoluble 
 in water and in alcohol. Light changes it to mercuric iotiide and mercury. 
 Aqueous solutions of potassium iodide have the same effect. 
 
 Mercurous Oxide, Hg^O, is a black mass, and is formed by the 
 action of potassium or sodium hydroxide upon mercurous salts. In the 
 light or at 100° it decomposes into mercuric oxide and mercury. 
 
 Mercurous Nitrate, HgNOg or Hg2(N03\, is produced by allowing 
 dilute nitric acid to act upon excess of mercury in the cold. It crystal- 
 lizes with one molecule of water in large monoclinic tables. It dissolves 
 readily in water acidulated with nitric acid ; pure water partly decom- 
 poses it with the separation of a yellow basic salt of the composition 
 
 „ OH 
 
 The nitric acid solution of mercurous nitrate oxidizes when exposed to 
 the air, and gradually becomes mercuric nitrate; this may be jirevented 
 by adding metallic mercury to the solution, whereby the resultant mer- 
 curic salt is again changed to the mercurous state: 
 
 Hg.N03)2 f Ilg = Hg,(N03). 
 
MERCURIC COMPOUNDS. 
 
 325 
 
 Mercurous Sulphate, Hg2(SOJ, results when an excess of mercury 
 is heated gently with concentrated sulphuric acid ; it separates as a crystal- 
 line precipitate, difficultly soluble in water, if sulphuric acid be added to 
 a mercurous nitrate solution. It fuses upon ap})lication of heat, and de- 
 composes into sulphur dioxide, oxygen and mercury. 
 
 Mercurous Sulphide, HggS, is not definitely known. A black com- 
 pound is produced in dilute solutions of mercurous nitrate by hydrogen 
 sulphide or alkaline sulphides. It contains mercury and mercuric sul- 
 phide : 
 
 Hg^S =: Hg + HgS. 
 
 MERCURIC COMPOUNDS. 
 
 Mercuric Chloride, HgClj, corrosive sublimate (^Hydrargyrum bichlo- 
 ratum), is produced when mercuric oxide is dissolved in hydrochloricacid, 
 or metallic mercury in aqua regia. It is obtained on a large scale by the 
 sublimation of a mixture of mercuric sulphate with sodium chloride: 
 
 HgSO, + 2 NaCl = HgCb + Na^SO,. 
 
 It crystallizes from water in rhombic prisms, and dissolves at medium 
 temperatures in 15 parts of water, at 100° in 2 parts; it is still more 
 soluble in alcohol. Its specific gravity is 5.4. It fuses at 265° and 
 boils at 307°, Its critical pressure is about 420 mm. (p. 229). The 
 vai)or density corresponds to the molecular formula HgClj (= 271.2). 
 
 Reducing substances, like sulphur di'<xide and stannous chloride, 
 change it to insoluble mercurous chloride: 
 
 2 HgCl 2 -f SO2 + 2H2O = Hg 2 Cl 2 + H 2 S 0 ^ + 2HCI. 
 
 Stannous chloride first precipitates mercurous chloride: 
 
 2HgC]2 + SnClj = Hg2Cl2 + SnCh, 
 
 which is afterward reduced, by the excess of stannous chloride, to me- 
 tallic mercury : 
 
 iig2a2 + Sncq = 2Hg -f Sncq. 
 
 Mercuric chloride is greatly inclined to form double salts with metallic 
 chlorides, e. g., HgCl2.KCl -j- H2O. These may be regarded as the 
 salts of a hydrochlormercuric acid, H.HgClgfp. 271). The acidity of 
 an aqueous solution of mercuric chloride is neutralized by the addition 
 of sodium chloride, which would indicate the formation of such salts. 
 When ammonium hydroxide is added to its solution, a heavy white pre- 
 cipitate is produced. Its composition varies with the concentration, the 
 temperature, and the quantities of mercuric chloride and ammonia. 
 White precipitate (^Hydrargyrum prcecipitatum alinin') has the formula 
 HgClNHj. This compound may be regarded as mercuric chloride, in 
 
326 
 
 INORGANIC CHKMISTRY. 
 
 which one atom of chlorine is replaced Ijy the amido-gronp, NH, : 
 
 , and it has been called amido-mcrcury chloride. Or it may 
 
 be derived from ammonium chloride by the substitution of a mercury 
 atom for two hydrogen atoms, NHglljCl, mercuramnumium chloride. 
 Similar mercuri- and mercuro-amimmium derivatives are numerous. 
 
 Mercuric Iodide, Hgl2 (^Hydrargyrum hiiodafum), is formed by the 
 direct union of mercury with iodine. When i)otassium iodide is added 
 to a solution of mercuric chloride, mercuric iodide separates as a yellow 
 precipitate, which immediately becomes red. It is readily soluble in 
 mercuric chloride and ])otassium iodide solutions; it crystallizes from 
 alcohol in bright red cpiadratic pyramids. Upon heating dry mercuric 
 iodide to 150°, it suddenly becomes yellow, and fuses at 223° to a red 
 liquid; a i)ortion subliming before this in yellow, shining, rhombic 
 needles. On touching these with some solid, they become red, with 
 sei)aration of heat, and are changed into an aggregate of quadratic 
 pyramids. Mercuric iodide is therefore dimorphous. It crystallizes 
 (depending on the temperature), in either yellow or red forms, from 
 methylene iodide, 01^^12, of which 100 jjarts dissolve 16.6 parts of mer- 
 curic iodide at 100°. 
 
 The great resistance shown by all the mercury halides toward concen- 
 trated sulphuric acid is noteworthy. Even in the heat the halogen is very 
 slowly ex])elled. They are also very stable toward caustic soda and pot- 
 ash. This is in harmony with the tendency of mercury to yield deriva- 
 tives which are only slightly dissociated. 
 
 Mercuric Cyanide, Hg(CN)2, is obtained by dissolving yellow mercuric oxide in 
 aqueous prussic acid or by heating Prussian blue with mercuric oxide and water. It crys- 
 tallizes from water in white quadratic prisms. It is also soluble in alcohol and can be 
 extracted from its aqueous solution by shaking with ether. It breaks down on heating 
 into mercury and cyanogen gas (p. 236). The mercury and cyanogen group seem to be 
 differently combined from other metals and cyanogen, e. g., potassium cyanide. Mercuric 
 cyanide has a very low degree of dissociation. Silver nitrate does not precipitate silver 
 cyanide from its aqueous solution, nor does an alkaline hydroxide throw out mercuric 
 oxide. Hydrogen sulphide and alkaline sulphides decompose it with formation of mer- 
 cury sulphide and hydrocyanic acid or alkaline cyanide. Unlike potassium cyanide it is 
 very stable toward acids. 
 
 Mercuric Oxide, HgO, is obtained by the prolonged heating of 
 metallic mercury near the boiling point in the air, orUy the ignition of 
 mercurous or mercuric nitrate and mercury. It forms a red, crystalline 
 powder (^Hydrargyrum oxydatum rubruui), of specific gravity 1 1.2. When 
 sodium hydroxide is added to a solution of mercuric chloride, mercuric 
 oxide separates as a yellow, amorjihous itrecijn'tate {Hydr. oxyd. flavum via 
 hu?nida paratuni). Both modifications become black when heated, but 
 change to a yellowish-red on cooling. Mercuric oxide decomposes into 
 mercury and oxygen at about 400°. 
 
 Mercuric oxide combines directly with ammonia, to form the compound 
 2HgU. NII3, which explodes with violence when heated. 
 
 Mercuric Nitrate, Ilg^NO^b^- It is difficult to obtain this salt pure, 
 because it is inclined to form basic com})ounds. A solution of it may be 
 
MERCURIC COMPOUNDS. 
 
 327 
 
 made by dissolving mercury or mercuric oxide in an excess of hot nitric 
 acid. On diluting the solution with water the basic salt, Hg(N03)2. 2HgO 
 -j- HjO, separates, and this may be converted into pure mercuric oxide 
 by boiling with water. 
 
 Mercuric Sulphate, HgSO^, is produced by digesting mercury or 
 its oxide with an excess of concentrated sulphuric acid. It forms a white, 
 crystalline, insoluble mass, which becomes yellow on heating. It yields 
 the hydrate HgSO^ -f- H^O with a little water, but much of the latter 
 decomposes it into sulphuric acid and the yellow insoluble basic salt, 
 HgSO^. 2HgO ( Turpcthum mme 7 'ale, Turpeth mineral). 
 
 Mercuric sulphate forms double salts wjth the alkaline sulphates, e. g., 
 HgSO^.K2SO^ 6H2O ; these are isomorphous with the corresponding 
 double salts of the magnesium grouj) (p. 315). 
 
 Mercuric Sulphide, HgS, 0( curs in nature as cinnabar, in radiating 
 crystalline masses, or in hexagonal prisms of dark-red color. It is obtained 
 artificially by rubbing together mercury and flowers of sulphur with water, 
 or it is produced as a black microcrystalline precipitate by the precipita- 
 tion of a solution of a mercuric salt with hydrogen sulphide. If the black 
 sul])hide be heated with exclusion of air it sublimes as a dark-red mass of 
 radiating crystalline structure, and is perfectly similar to natural cinnabar. 
 A similar conversion of the black modification into the red is effected by 
 continued heating of the same to 50° with a solution of potassium or 
 ammoniumsulphide. The red mercury sulphide thus obtained is employed 
 as artificial cinnabar in painting. 
 
 The mercury compounds can be readily recognized by the following 
 reactions. When heated with dry sodium carbonate in a small test- 
 tube, mercury escapes, and condenses u])on the side in metallic drops. 
 Tin, copper and zinc throw out metallic mercury from its solutions. 
 If a piece of sheet copper be dipped into the same, mercury is de- 
 posited as a gray coating, which on being rubbed acquires a metal- 
 lic luster. When a dry mixture of a mercuric salt, lime and potassium 
 iodide is heated yellow (or red) mercuric iodide sublimes. This is 
 the only iodide not decomposed by caustic lime. The mercurous com- 
 pounds are distinguished from the mercuric by their precipitation by 
 hvdrochloric acid. 
 
 The heats of formation of the chlorides of the seven metals of the second group are 
 as follows : 
 
 (Mg,Cl 2 ) = 15 1 
 
 (Ca,Cb) 170 (Zn,Cl2) =-- 97 
 
 (Sr,Cl2) = 185 (Cd,Ch) = 93 
 
 (Ba,Cl2) = 195 (llg,('l,) -54. 
 
 I'he heats of formation in the first sub-class increase with the rise in the atomic weights, 
 while the opposite is observed in the second sub-class. The heats of formation of other 
 derivatives of the two groups manifest similar relations. 
 
328 
 
 INuKGANIC CHEMISTRY. 
 
 Comparing these numbers with the quantity of heat which is disengaged in the forma- 
 tion of aqueous hydrochloric acid : 
 
 II,C 1 = 22; ir,Cl,Aq 39.3, 
 
 we find explained the behavior of the metals toward this acid. All metals liberating a 
 greater (juantity of heat than 39.3 Cal. in the formation of their chlorides (calculated for 
 I equivalent of metal) are in condition to decompose the dilute acid. Most of the metals 
 belong to this class; mercury, copper, silver, gold, lead, thallium, and some others, set 
 free a less amount of heat, and hence are not able to decompose dilute hydrochloric acid. 
 The slight quantity of heat developed in the formation of hydrogen sulphide, 
 
 S,H, = 4.5, 
 
 indicates that the same is readily decomposed by all the metals. In the same way, by 
 adding the heat of solution, 
 
 S,H2,Aq = 9.2, 
 
 we can easily ascertain which metals are precipitated by hydrogen sulphide from their 
 chlorides, etc. 
 
 If in the thermo-chemical equation, 
 
 (Me,Cl„Aq) (S,H2,Aq) = (Me,S) -f 2(II,Cl,Aq), 
 
 the sum of the heat developed upon the right side is greater than that upon the left, the 
 reaction will occur (precipitation of metallic sulphides) ; in the opposite case the sulphide 
 is decomposed by the dilute hydrochloric acid. 
 
 COPPER, SILVER, AND GOLD. 
 
 Considering the magnitude of their atomic weights, copper, silver, and 
 gold, bear the same relation to the alkali group, especially to sodium, as 
 zinc, cadmium, and mercury bear to magnesium : 
 
 Na = 23.015 
 
 Cu= 63.6 
 Ag= 107.93 
 Au = 197.2 
 
 Mg= 24.36 
 Zn = 65.4 
 Cd =112 
 Hg = 200.3. 
 
 They occupy an entirely analogous position in the three great periods of 
 the periodic system of the elements (p. 246), and constitute the transition 
 from the elements of group VIII, especially from cobalt, palladium, and 
 platinum, to the elements of group II — zinc, cadmium, and mercury : 
 
 Co= 59 
 Pd — 106 
 Pt = 194.8 
 
 Cu = 63.6 
 Ag^ 107.93 
 Au = 197.2 
 
 Zn = 65.4 
 Cd =112 
 Ilg = 200.3. 
 
 This intermediate position of the three elements about to be discussed 
 is clearly shown in their entire physical deportment. While the ele- 
 ments of group VIII, with the last members, cobalt, palladium, and 
 ])latinum, fuse with difficulty and do not volatilize, copper, silver, 
 and gold constitute the transition to the readily fusible and volatile 
 elements, zinc, cadmium, and mercury. They take an intermediate 
 I)osition, too, with reference to their coefficients of expansion, their 
 
COPPER, SILVER, AND GOLD. 
 
 329 
 
 atomic volumes, and other physical properties. It is noteworthy that the 
 ability to conduct heat and electricity attains its maximum in copper, sil- 
 ver, and gold. 
 
 Not only are the properties of the free elements determined by the 
 position of the latter in the periodic system, but those of their deriva- 
 tives, and especially such as depend upon the valence of the elements, are 
 influenced to a marked degree by the above relation. In consequence of 
 the double periodicity of the great periods, copper, silver, and gold attach 
 themselves to group I, and especially to sodium, just as the elements im- 
 mediately following, zinc, cadmium and mercury, arrange themselves with 
 group II and magnesium. Hence we find copper, silver, and gold, like 
 
 I 
 
 sodium, yielding compounds of the form MeX, in which they appear univa- 
 lent. Some of these are isomorphous, e.g., silver sulphate, Ag2S04, with 
 sodium sulphate, Na2SO^; sodium chloride, NaCl, cuprous chloride, CuCl, 
 and silver chloride, AgCl, crystallize in forms of the regular system. 
 
 But we may say that the similarity of copper, silver, and gold to 
 sodium is confined to these few external properties. Just as the heavy 
 metals, zinc, cadmium, and mercury differ in many properties from the 
 light metal magnesium (p. 313), so do the metals copper, silver, and 
 gold, possessing a high specific gravity, distinguish themselves in a still 
 higher degree from the light metal sodium. They possess all the prop- 
 erties belonging to the heavy metals, which are mainly characterized by 
 the insolubility of the oxides, sulphides, and many salts, so that they 
 arrange themselves in a less marked degree with the alkali metals. 
 
 I 
 
 In the compounds constituted according to the form MeX, in which 
 copper, silver, and gold appear univalent, they exhibit great similarity. 
 Thus, the chlorides, CuCl, AgCl, and AuCl, are white and insoluble in 
 water; soluble, however, in concentrated hydrochloric acid, ammonia, 
 the alkaline hyposulphites, etc., and form double chlorides, very similar 
 to each other, with other chlorides. While silver only enters compounds 
 
 I 
 
 of the form AgX, copper and gold are capable of yielding another form ; 
 
 I II 
 
 copper forms, besides cuprous, CuX, also cupric, CUX2, derivatives, in 
 which it appears to be bivalent. The latter are much more stable than the 
 former, and embrace the ordinary copper salts. Gold, however, besides 
 
 I III 
 
 furnishing aureus, AuX, compounds, has aurA derivatives, AuXg, in which 
 the metal appears to be trivalent (see the Thallium Group). 
 
 While copper and gold, in their lower forms, are analogous to silver (and 
 in less degree, to sodium), the cuprf<: derivatives show a great resemblance 
 to the compounds of the metals of the magnesium group and other metals 
 in their bivalent combinations. Thus, the sulphates of zinc, magnesium, 
 cupric oxide (CuO), ferrous oxide (FeO), nickelous oxide (NiO), cobalt- 
 ous oxide (CoO), and manganous oxide (MnO), are similarly constituted, 
 resemble each other, are isomorphous, and form entirely analogous double 
 salts (p. 315) with the alkaline sulphates. In the same way the carbonates 
 
 II II 
 
 (MeCOg), the chlorates and bromates (MeCl20g -f 6H2O) and others, are 
 28 
 
330 INORGANIC CHEMISTRY. 
 
 similarly conslituted and isomorphons. In its aur/c derivatives, gold cx- 
 
 III 
 
 hibits some similarity to the aluminium compounds (MX ), to those of 
 
 III 
 
 indium (InX,,) and other metals, in their trivalent combinations. Here 
 we see, as already ob.served with mercury (j). 321), that the similarity of 
 ilie compounds of the metals is influenced by the similarity of forms or types, 
 according to which they are composed, i. e., by the valence of the metals. 
 If a metal form several series of compounds of different types, each series 
 is usually more or less similar to the compounds of other metals of like 
 type. In this manner is shown the resemblance of the compounds of the 
 following types ; 
 
 Na^C 
 
 AggO 
 
 Cu.^C) 
 
 Aiqf) 
 
 TljO 
 
 Sodium oxide. 
 
 Silver oxide. 
 
 Cuprous oxide. 
 
 Aurous oxide. 
 
 Thallous oxide. 
 
 MgO 
 
 Zn(.) 
 
 CuO 
 
 Fet ) 
 
 IlgO 
 
 Magnesium oxide. 
 
 Zinc oxide. 
 
 Cupric oxide. 
 
 Ferrous oxide. 
 
 Mercuric oxide. 
 
 AiPs Au,03 TI,()3. 
 
 Aluminium oxide. I-erric oxide. Auric oxide. Tliallic oxide. 
 
 The character of their derivatives, varying with the degree of combina- 
 tion or valence, becomes quite marked with chromium, manganese and 
 iron, as we shall later see. The heavy metals also exhibit a strongly 
 positive and basic character in their univalent combinations. Thus, silver 
 oxide (Ag.^O) and thallous oxide (T^O) are strong bases, forming neutral- 
 reacting salts with acids, and even cuprous and aurous oxides are more 
 strongly basic than their higher forms of oxidation. The metalloidal 
 character of the metals and the acid nature of their oxides begin to appear 
 in their trivalent combinations. Thus, in the hydroxyl derivatives of 
 aluminium, indium, and gold, A1(0H)3, In(OH)3, Au(0H)3, hydrogen 
 may be replaced by the alkalies just as in boric acid, B(OH)3. Their 
 higher forms of oxidation show, like those of the metalloids, a pronounced 
 acid-like character (as Pb02, PtOg, CrOg, FeOg) which is only lessened by 
 a high atomic weight of the metal (as in PbOg and PtOg). 
 
 The heats of formation of the copper, silver and gold compounds of the MeX type 
 .show the same gradation as those of the zinc group (p. 327) : 
 
 (Na, Cl) = 97.6 (Ag, Cl) = 29.3 
 
 (Cu, Cl) = 32.8 (Au, Cl) = 5.8. 
 
 COPPER. 
 
 Cu = 63 . 6 . 
 
 Native copper is found in large quantities in America (Lake Superior), 
 China, Japan, also in Sweden and in the Urals. It frecpiently occurs 
 crystallized in cubes and octahedra. The most im|)ortant and most widely 
 distributed of its ores are: cuprite (Cu.^O), malachite (CuCOg. Cu(OH).2) 
 
COPPE'K. 
 
 33 ^ 
 
 and aziirite (2CuCO.^. Cu(OH)2') (basic carbonates), chalcocite (Cu.^S), and 
 especially chalco})\ rile (Ci]2S . Fe2S3) and bornite (3CU2S. Fe2S3). 
 
 Metallurgy of Copper. — 'I'he extraction of copper from its oxygen ores is very simple : 
 metallic copper is melted out when the ores are ignited along with charcoal. The sul- 
 phur ores are more difficult to work. i. The divided material is first roasted in the air, 
 whereby the sulphur content is reduced : 
 
 3(Cu 2S. Fe2S3) -f 21O = 6CuO -f- 5FeS -j- FeO -f ySOg. 
 
 2. The mass is afterward fused, when we get the so-called coarse metal, which con- 
 tains iron sulphide in addition to copper sulphide : 
 
 6CuO + sFeS ^ (3CU2S + FeS) + 4FeO + SO2. 
 
 3. Crude copper (black copper) is obtained by further roasting and by the “fusion 
 reaction” of this mass: 
 
 3CU2S + FeS -f 9O := 3CU2O -f FeS -f 3SO2 = 3CU2 + FeO -f 4SO2. 
 
 The ferrous oxide formed here is taken up by the silica to form a slag. In the more 
 recently constructed works, particularly in the United States, the sulphide ores are first 
 roasted and then melted in cupola or reverberatory furnaces, after which the coarse metal 
 is subjected to the Bessemer process and the resulting black copper is electrolyzed. See 
 Chem. Zeit. 1895, 977 ; 1897, 995 ; also Jahrb. d. Chem. v (1895', 287. 
 
 In refining copper electrcilytically, plates cast from black copper are suspended as 
 anodes in the solution of copper sulphate. The kathodes are plates of pure copper. The 
 copper is dissolved from off the anode and deposited on the kathode. Impurities, includ- 
 ing any of the nobler metals, settle at the bottom as mud. 
 
 Copper can also be obtained in the wet way from its ores. The pulverized ore is 
 treated with solutions of ferric chloride or sulphate, when the copper sulphide is trans- 
 posed to soluble copper salts in the sense of the following equations : 
 
 CU2S -f Fe2Clg = 2FeCl2 + CU2CI2 -j- S ; 
 
 CuS -f Fe2Clg = CuCb f- 2FeCl2 -}- S ; 
 
 CU2S 4 - 2Fe2(SOb3 = 2CUSO4 -f 4FeS04 S. 
 
 The copper is then precipitated electrolytically from the solution. 
 
 The following countries are rich in copper : United States (Lake Superior and Mon- 
 tana), Spain and Portugal, also Chili, Japan, and Germany (Prussia, Mansfield). 
 
 The following are excellent reference works on the metallurgy of copper : Ost’s Tech- 
 nologie, Dammer’s Technologie (p. 190), and Durre’s Vorlesungen fiber allgemeine 
 Hfittenkunde (Halle, 1898). 
 
 To obtain chemically pure copper, the pure oxide is heated in a stream 
 of hydrogen, or the solution of copper sulphate is decomposed by 
 electrolysis. 
 
 Metallic copper possesses a characteristic red color, and transmits a 
 green light in thin leaflets. It is rather soft and ductile, and possesses a 
 specific gravity of 8. 5-8. 9. It melts at about 1080°, and vaporizes in 
 the oxyhydrogen flame. It remains unaltered in dry air; in moist, it is 
 gradually coated with a green layer of basic copper carbonate. When 
 heated in the air, it oxidizes to black cupric oxide. 
 
 Copper is not changed by dilute hydrochloric or sulphuric acids; if it 
 be moistened with these, and exj)osed to the air, it absorbs oxygen, and 
 gradually dissolves. It issimilarly dissolved by ammonium hydroxide. Hot 
 concentrated sulphuric acid converts it into copper sulphate, with evolu- 
 
33 ^ 
 
 INORGANIC CHEMISTRY. 
 
 tion of sulphur dioxide. It dissolves in dilute nitric acid in the cold, 
 with evolution of nitric oxide. Zinc, iron and al.so jihosphorus precipi- 
 tate metallic copper from the aqueous solutions of its salts. 
 
 Copper forms two series of compounds, known as cu\nous and cupr/V. 
 In the latter co})})er is bivalent : 
 
 CuO CuGj Cu(OII), CuSO,. 
 
 These are more stable than the cuprous derivatives; the ordinary copper 
 salts belong to them. In many respects they resemble the compounds 
 of other dyad metals, especially those of the magnesium group, and the 
 ous comj)ounds of iron (FeO), manganese (MnO), cobalt and nickel 
 (see p. 329). 
 
 The cuj)rous compounds are, on the other hand, very unstable, absorb 
 oxygen from the air, and pass into cupric derivatives. They show some 
 similarity to the mercurous derivatives (p. 323), and possess an analogous 
 composition : 
 
 CuCl Cul CU2O Cu.^S. 
 
 Oxygen salts of cuprous oxide are not known. 
 
 From the formulas given above, copper, like silver, is univalent in the cuprous com- 
 pounds. This is also experimentally indicated by the boiling and melting points shown 
 by solutions of cuprous salts in pyridine, piperidine, and methyl sulphide. The salts 
 are not dissociated in these solutions ; they are non-conductors. From what has been 
 said on p. 268 some of them probably have the simple formula, especially cuprous 
 bromide, CuBr. 
 
 Other salts are disposed to appear as double molecules ; the cyanide is only known as 
 Cu 2(CN)2'. We can therefore assume, as was done in the case of the mercurous salts, 
 the presence of a bivalent group of two metallic atoms : 
 
 CN-Cu-Cu-CN Cl-Cu-Cu-Cl. 
 
 The vapor density of cuprous chloride even at 1 700° corresponds to the formula CU2CI2. 
 Both formulas should be given equal value. [See pp. 322, 324, 339, and Werner, Z. f. 
 anorg. Ch. 15 (1897) l]. 
 
 CUPROUS COMPOUNDS. 
 
 Cuprous Oxide, CiqO, occurs as cuprite crystallized in regular 
 octahedra. It is obtained artificially by boiling a solution of copper sul- 
 phate and grape sugar with potassium hydroxide, when it separates as a 
 crystalline, bright red powder. It does not change in the air, and is 
 readily soluble in ammonium hydroxide. The solution absorbs oxygen, 
 and while forming cupric oxide acquires a blue color. By the action of 
 sulphuric and other oxygen acids, it forms cupric salts, half of the copper 
 sejjarating as metal : 
 
 CU2O -f II2SO, CuSO, + Cu + H2O. 
 
 d'he hydroxide^ Cu.2(OH)2, is precipitated by the alkalies as a yellow pow- 
 
CUPROUS COMPOUNDS. 333 
 
 der from hydrochloric acid solutions of cuprous chloride, CU2CI2. It 
 oxidizes in the air to cupric hydroxide. 
 
 Cuprous Chloride, CuC.l or CU2CI2, is produced by the combustion 
 of metallic copper in chlorine gas (together with cupric chloride, CuClj), 
 upon conducting hydrochloric acid over copper at alow, red heat; by 
 boiling the solution of cupric chloride with copper: 
 
 CuCb -|- Cu = CU2CI2, 
 
 and by the action of many reducing substances upon cupric chloride. It 
 is most conveniently made by passing sulphur dioxide through a concen- 
 trated solution of copper sulphate and sodium chloride, when it separates 
 as a white, shining powder, consisting of small tetrahedra. It fuses at 
 430°, and distils at about 1000°; its vapor density at 1600-1700° corre- 
 sponds to the formula CujClj. In moist air it rapidly becomes green* 
 owing to oxygen absorption, and the formation of basic cupric chloride. 
 Cl 
 
 Cu<oH- becomes black on exposure to light. Cuprous chloride is 
 
 readily soluble in concentrated hydrochloric acid and in ammonium 
 hydroxide; both solutions jjossess the characteristic property of absorb- 
 ing carbon monoxide. Colorless leaflets separate from solutions saturated 
 with the gas. Their composition seems to be CuaCb -f CO -j- 2H2O. 
 
 Cuprous Iodide, Cul or Ci]2l2, is precipitated from soluble cupric 
 salts by potassium iodide : 
 
 2CuS0^ -f 4KI = CU2I2 + 2X280^ + I2. 
 
 By extracting the co-precipitated iodine by means of ether it is 
 obtained as a gray powder, insoluble in acids. 
 
 Cuprous Sulphide, CU2S, occurs as chalcocite crystallized in forms of 
 the rhombic system. It is produced by burning copper in sulphur vapor, 
 and by heating cupric sulphide in a current of hydrogen ; after fusion it 
 solidifies in crystals of the regular system. Combined with silver sul- 
 
 Cu 'I 
 
 phide it constitutes the mineral stromeyerite, | S or CU2S. Ag2S, iso- 
 morphouswith chalcocite. 
 
 Copper Hydride, CuH or CU2H2, belongs to the derivatives of 
 univalent copper. Jf a solution of copper sulphate be digested with 
 hypophosphorous acid, the hydride separates as a yellow amorphous pre- 
 cipitate which soon acquires a brown color. At 60° it decomposes into 
 copper and hydrogen. 
 
 With hydrochloric acid it forms cuprous chloride : 
 
 CU2H2 + 2HCI = Cu2a2 + 2H2. 
 
 (Compare Mylius and Fromm, Ber. 27 (1894), i, 647.) 
 
334 
 
 INORGANIC CHKMIS'IRV. 
 
 CUPRIC COMPOUNDS. 
 
 'riie cuiiric salts, when liydrous, are generally colored blue or green ; 
 they are colorless in the anhydrous condition. 
 
 Cupric Hydroxide, Cu(()ll),^, separates as a voluminous bluish pre- 
 cipitate when sodium or potassium hydroxide is added to soluble copper 
 salts. \Vhen heated, even under water, and particularly in the jiresence 
 of alkali, it loses water, and is changed to black cupric oxide. In the 
 presence of alkaline chlorides, or of cui)ric chloride it is i)artly reduced 
 to a cuprous salt with the production of an alkaline peroxide in the 
 evolution of oxygen : 
 
 2Cu(0II), + 2KBr = CujBr, + K./)^ f 2ll,0. 
 
 (Spring and Liicion, Z. f. anorg. Ch. 2 (1892), 195). 
 
 Cupric Oxide, CuO, is usually obtained by the ignition of copper 
 turnings in the air, or by heating cupric nitrate. It forms a black 
 amorphous powder, which, at higher temperatures, settles together and 
 acquires a metallic luster. By heating with organic substances their 
 carbon is converted into carbon dioxide, and the hydrogen into water, 
 th-; ciq^ric oxide being reduced to metal ; upon this rests the application 
 of cupric oxide in the analysis of such comj)ounds. 
 
 Copper oxide and hydroxide dissolve in ammonium hydroxide with a 
 dark-blue color. The solution possesses the power of dissolving wood 
 fiber (cotton-wool, linen, filter-paper, etc.) — Schweizer' s reagent. 
 
 Cupric Chloride, CUCI2, is formed by the solution of cupric oxide 
 or carbonate in hydrochloric acid. It crystallizes from aqueous solution, 
 with two molecules of water, in bright green rhombic needles, and is 
 readily soluble in water and alcohol. When heated, it parts with its 
 water, becoming anhydrous chloride, which at a red heat is decomposed 
 into chlorine and cuprous chloride. It yields beautifully crystallized 
 double salts with potassium and ammonium chlorides. Cupric bromide 
 is like the chloride ; the iodide is not known, since in its formation it at 
 once breaks down into cuprous iodide and iodine. 
 
 Copper Sulphate, CuSO^ + cupric suli)hate, copper vitriol — 
 
 may be obtained by the solution of copper in concentrated sulphuric 
 acid. It is produced on a large scale by roasting chalcocite. It forms 
 large blue crystals of the triclinic system, which effloresce somewhat upon 
 exposure. At 100° the salt loses four molecules of water ; the fifth sepa- 
 rates above 200°. The anhydrous sulphate is colorless, absorbs water 
 very energetically, and returns to the blue hydrous compound. 
 
 Although copi)er sulphate only crystallizes with five molecules of water, 
 it is capable, like the sulphates of the magnesium group, of forming 
 double salts with potassium and ammonium sulphates, which crystallize 
 with six molecules of water, and are isomorphous with the double salts 
 of the metals of the magnesium group. 
 
 ('oppcr sulphate is em[)loyed in electro-plating. When its solution is 
 decomposed by the galvanic current copjier separates at the negative 
 ])ole, and deposits in a regular layer upon the conducting objects con- 
 
CUPRIC COMPOUNDS. 
 
 335 
 
 nected wiih the electrode. The positive pole is a copperplate: the 
 ions, SO^, dissolve a quantity of it equal to the copper deposited on the 
 negative pole; hence the solution always has the same concentration. 
 
 Ammonium hydroxide added to a copper sulphate solution in sufficient 
 quantity to dissolve the cupric hydroxide produced at first, changes the 
 color of the liquid to a dark blue. From this solution alcohol precipitates 
 a dark-blue crystalline mass with the composition CuSO^.qNHg -}- H2O. 
 Heated to 150° this compound loses water and two molecules of ammonia, 
 and becomes CUSO4.2NH3. It is supposed that these compounds are 
 ammonium salts in which a part of the hydrogen is replaced by copper; 
 they have been designated cuprammoniimi compounds, e. g. : 
 
 Cuprammonium sulphate. 
 
 The other soluble copper salts yield similar compounds with ammonium 
 hydroxide. 
 
 Cupric Nitrate, Cu(N03)2, crystallizes with three or six molecules of 
 water, has a dark-blue color and is readily soluble in water and alcohol. 
 Heat converts it into cupric oxide. 
 
 Copper Carbonates. — The neutral salt (C11CO3) is not known. 
 When sodium carbonate is added to a warm solution of a copper salt the 
 
 basic carbonate, CuCOg . Cu(OH)2, or separates as a 
 
 green precipitate. It occurs in nature as malachite, which is especially 
 abundant in Siberia. Another basic salt, aCOgCu . Cu(OH ^2, is the beau- 
 tiful blue nzurite. 
 
 Copper Arsenite separates as a beautiful bright green precipitate, 
 upon the addition of sodium arsenite to a copper solution. This preci- 
 ])itate has never been obtained in homogeneous form and of a definite 
 composition. It was formerly employed as a pigment, under the name 
 of Scheele' s green, but at present, owing to its poisonous character, it has 
 been replaced by other green colors (Guignet’s green and aniline green). 
 
 Cupric Sulphide, CuS, is a black compound, precipitated from cop- 
 per solutions by hydrogen sulphide. It is insoluble in dilute acids. 
 When moist, it slowly oxidizes in the air to cupric sulphate. Heated in 
 a stream of hydrogen, it forms cuprous sulphide, CUgS. 
 
 Alloys of Copper. — Pure copper is very ductile, and may be readily 
 rolled, and drawn out into a fine wire. It does not make good cast- 
 ings, because it contracts unequally upon cooling and does not fill out 
 the moulds. For such pur])oses, alloys of copper are em]:)loyed which, 
 in addition, possess other technically valuable properties. The most 
 important copper alloys are: 
 
 Brass, consisting of three parts of copjjer and one part of zinc. It has 
 a yellow color, and is considerably harder than ]jure co])per. Ordinarily, 
 one or two per cent, of lead are added to the brass, which facilitates its 
 
33 ® 
 
 INORGANIC CHEMISTRY. 
 
 working upon the turning-lathe. Tombac contains 15 per cent, of zinc, 
 and has a gold-like color. The alloy of i part of zinc and 5.5 parts of 
 copper answers for the manufacture of spurious gold leaf, d'he alloys of 
 copper with tin are called bronzes. Most of the modern bronzes also 
 contain zinc and lead; those from Japan, gold and silver. The cannon 
 bronze contains 90 per cent, of coi)per and 10 per cent, of tin ; bell-metal 
 has 20-25 cent, of tin. 
 
 Argentan is an alloy of copper, zinc, and nickel (see Nickel). The 
 German cop])er coins consist of 95 per cent, of coj)jjer, 4 [)er cent, of tin, 
 and I per cent, of zinc. 
 
 Of the more recently introduced alloys of copi)er we may mention : 
 
 Phosphor-bi'Ofize. — This consists of 90 parts of cop|)er, 9 parts of tin 
 and 0.5-0. 8 parts of phosphorus. By the last ingredient the bronze is 
 increased in hardness, and its solidity and resistance to oxidation are also 
 increased. It is em})loyed in making axle bearings and various parts of 
 machinery. 
 
 Silicon bronze, containing silicon instead of phosphorus, is charac- 
 terized by great firmness and conductivity. It is used for telephone 
 wires and as conductii'g cables for electric railways. 
 
 Manganese bronze contains 70 per cent, of copi)er and 30 per cent, of 
 manganese (it is the cupro-manganese of Letrange), and may be melted 
 with copper and copper alloys, im]3arting to these solidity and great 
 hardness. It is especially employed in making ship-propellers. 
 
 See page 347 for Aluminium Bronze. 
 
 Recognition of Copper Compounds. — Most copper compounds 
 containing water have a blue or green color. With the exception of 
 copper sulphide they all dissolve in ammonium hydroxide, with a blue 
 color. Potassium ferrocyanide produces a reddish-brown precipitate or 
 color in aqueous solutions of cupric salts. When a clean piece of iron is 
 introduced into a co|)per solution, it becomes covered with a red layer 
 of metallic copper. Volatile copper compounds tinge the flame blue or 
 green. The spectrum of such a flame is characterized by several blue or 
 green lines. 
 
 SILVER. 
 
 Ag = 107.93. 
 
 Silver occurs native. Its most important ores are argentite or silver 
 glance (silver sulphide, Ag^S) and various compounds with sulphur, 
 arsenic, antimony, copper and other metals (light-red silver ore, prous- 
 tite, Ag^AsSj, and dark-red silver ore, pyrargyrite, Ag3SbS3). Of rarer 
 occurrence are combinations with chlorine (horn-silver, AgCl), bromine 
 and iodine. Slight quantities of silver sulphide are ])resent in almost 
 every galenite (PbS). The principal localities for silver ores are the 
 United States, Mexico, Bolivia, Australia, Saxony and Chili. 
 
SILVER. 
 
 337 
 
 Metalliirf:;y of Silver. — The metallurgy of silver ores involves methods dependent upon 
 the silver content of the ore, the chemical constitution of the latter and the conditions of 
 locality. It is usually a complicated procedure. The important methods alone will be 
 indicated briefly here ; details must be sought in special works upon chemical technology 
 (p. 331). It is only the richest ores from which the silver can be directly recovered ; 
 most of the metal is obtained either (i) by first preparing an argentiferous lead and then 
 separating these two metals subsequently ; or {2) by the amalgamation process ; or (3) by 
 converting the silver into soluble compounds and precipitating it from the solutions. 
 
 1. The ores are worked together with lead ores for metal or, if they contain more than 
 yL of silver, they are fused directly with lead. Various methods are pursued to enrich 
 the silver-lead alloy with the first metal, {a) Pattinsouiziug. When lead containing 
 silver is fused and cooled slowly at first pure lead crystallizes out which is removed with 
 sieves; an alloy rich in silver and readily fused remains. { 1 )) Parkes's process. This 
 method consi.sts in adding zinc to the silver-lead alloy of low silver content and obtaining 
 thereby an alloy of lead-zinc-silver, fusible with difficulty. When the fusion cools this 
 alloy appears on the surface as “zinc crust.” The zinc is expelled from it by distillation. 
 The lead now rich in silver (obtained by one or the other method) is cupelled, i. e., it is 
 heated with air-access in a furnace the bottom of which consists of a porous mass. The 
 lead is oxidized to lead oxide which partly escapes and is in part absorbed. The silver 
 is not oxidized and remains in the cupel. 
 
 2. The amalgamation process is based on the fact that mercury dissolves silver and 
 also decomposes its compounds with the formation of an amalgam. 
 
 The ore is roasted with sodium chloride, whereby silver chloride is produced. The 
 divided material is then mixed with iron scraps and water in rotating vessels. The iron 
 causes the precipitation of the metallic silver from its chloride : 
 
 2AgCl -|- Fe = FeCl2 -f 2Ag. 
 
 To free the metal from various impurities it is dissolved in mercury and the liquid 
 amalgam ignited, when mercury distils off and silver remains. Owing to scarcity of com- 
 bustible material, the conversion of silver ores into silver chloride is executed, in Mexico 
 and Peru, by mixing the ores with sodium chloride and copper sulphate in the presence 
 of water. In this way cuprous chloride is produced, which is transposed, with silver sul- 
 phide, into silver chloride and cuprous sulphide : 
 
 CU2CI2 + Ag2S = CU2S + 2AgCl. 
 
 3. By solution and precipitation. The ores are roasted alone and afterwards together 
 with salt. Silver sulphate is first produced and then silver chloride. The roasted mass 
 is extracted with a concentrated salt solution which dissolves the silver chloride, and 
 from this solution the silver metal is precipitated by metallic copper. The resulting 
 cuprous chloride solution is acted upon with iron to recover the copper (Augustin). Or 
 after roasting the ore alone silver sulphate is dissolved out with water and the solution 
 treated with copper (Ziervogel). The silver can also be dissolved out by sodium copper 
 thio.sulphate, the metal being then precipitated by means of a calcium sulphydrate solu- 
 tion (Russell). 
 
 Commercial silver (work-silver) is not pure, but invariably contains 
 copper and traces of other metals in greater or less quantity. To pre- 
 pare chemically pure metal, the work-silver is dissolved in nitric acid, 
 and from the solution of nitrates thus obtained hydrochloric acid pre- 
 cipitates the silver as chloride : 
 
 AgNOg -f HCl = AgCl + HNO3. 
 
 The latter is reduced by various methods; either by fusion with sodium 
 carbonate, or by the action of zinc or iron in the presence of water: 
 
 2AgCl -f Na-^CO, 2NaCl -f- 2Ag + COg | O. 
 
 2AgCl -f- Zn = ZnClj T 2Ag. 
 
 29 
 
338 
 
 INORGANIC CHEMISTRY. 
 
 Silver is a pure white, brilliant metal, of specific gravity 10.5. It is 
 tolerably soft and very ductile, and can be drawn out to a fine wire. It 
 crystallizes in regular octahedra. It fuses at 960°, and is converted into 
 a greenish vapor in the oxyhydrogen flame. It is the best conductor of 
 heat and electricity. Silver is not oxidized by oxygen ; by the action of 
 ozone it is covered with a very thin layer of silver peroxide. When in 
 molten condition, silver absorbs 22 volumes of oxygen without combin- 
 ing chemically with it ; the absorbed gas escapes again when the metal 
 cools [sprouting or si)itting of silver]. 
 
 Silver is capable of existing in three allotropic forms, which have 
 properties greatly different from those of ordinary silver. The first form 
 is soluble in water and has a blue color. The second variety is insoluble, 
 and somewhat resembles the first form. The third closely resembles gold 
 in color and luster. These allotropic varieties of silver are broadly dis- 
 tinguished from normal silver by color. They very likely are more active 
 conditions of silver, common silver being a polymerized variety (Am. Jr. 
 Science, 37, 476). 
 
 Silver unites directly with the halogens; by the action of hydrochloric 
 acid it becomes coated with an insoluble layer of silver chloride. Boiled 
 with strong sul])huric acid, it dissolves to sulphate: 
 
 ^■^8 T zIIjSO^ = Ag.^SO^ 4 “ SO2 T 2II2O. 
 
 The best solvent of silver is nitric acid, which even in a dilute state 
 and unaided by heat, converts it into nitrate. 
 
 As silver is rather soft, it is usually employed in the arts alloyed with 
 copper, whereby it acquires a greater hardness. Most silver coins consist 
 of 90 per cent, of silver and 10 per cent, of copper; the English shil- 
 lings contain 92.5 per cent, of silver. What is termed the fineness of 
 silver is the number of parts per thousand which the alloy contains. 
 
 Oxygen forms three compounds with silver, but only the oxide yields 
 corresponding salts. 
 
 Silver Oxide, AgjO, is thrown out of a silver nitrate solution by 
 sodium or potassium hydroxide as a dark-brown, amorphous precipitate. 
 It is somewhat soluble in water, and blues red litmus-paper. In this, and 
 in the neutral reaction of the nitrate, the strongly basic, alkaline nature 
 of silver and its oxide exhibits itself; the soluble salts of nearly all of 
 tlie other heavy metals show an acid reaction (p. 330). When heated to 
 250°, the oxide decomposes into metal and oxygen ; at too° it is reduced 
 by hydrogen. Silver hydroxide, AgOH, is not known ; the moist oxide 
 reacts, however, very much like an hydroxide. 
 
 On dissolving precipitated silver oxide in ammonium hydroxide, black 
 crystals (Ag.^O. 2NH3) separate when the solution evaix)rates, and when 
 dry these exi)lode ui)on the slightest disturbance — fulminating silver. 
 
 Silver Suboxide, Ag/), is said to l)e produced by heating silver citrate in a current 
 of hydrogen, and is a black, very unstable powder, which decomposes readily into silver 
 oxide and silver. 
 
SILVER. 
 
 339 
 
 Silver Peroxide, AgO or Ag202, is formed by passing ozone over silver or its oxide, 
 or by the decomposition of the nitrate by the electric current. It consists of black, 
 shining octahedra, and at ioo° decomposes into silver oxide and oxygen (Mulder). 
 
 The salts of silver correspond to the oxide Ag.^O, and are all constituted according to 
 the form AgX, hence are termed argentic. They are analogous to the cuprous and mer- 
 curous derivatives, and show a great resemblance to the former. It would, therefore, be 
 more correct to designate them argent£>«.y. Compounds of the bivalent form AgXj are 
 not known for silver. If, however, the mercurous and cuprous compounds are expressed 
 by double formulas (pp. 322, 332) : 
 
 CuCl Cu, llgCl Hg, 
 
 1 I )0 and I I )0 
 
 CuCl Cu-^ HgCl Hg/ 
 
 which view is supported by their chemical deportment, and is experimentally confirmed 
 in the case of cuprous chloride by its vapor density ; those of silver might be represented 
 by analogous formulas : 
 
 AgCl 
 
 I 
 
 AgCl 
 
 Ag\ 
 
 I )o 
 
 AgNOa 
 
 AgN 03 . 
 
 Then the silver atom would be bivalent and a complete parallelism would be estab- 
 lished with copper. In support of this view there exists the fact that when the silver 
 halides are dissolved in pyridine, methyl sulphide, etc., the colligative properties of 
 their solvents are influenced as if halide molecules possessing the formulas Ag2Cl2, Ag2- 
 Br2, Ag2l2, and possibly in part higher formulas, were present. See pp. 268, 332. The 
 chemical formulas of solid bodies do not generally designate their true molecular values 
 as in the case of gases, but only their simplest atomic composition. It is very probable 
 that even the simplest chemical compounds, e. g.^ potassium chloride and silver chloride, 
 consist in their solid condition of complex molecules corresponding to the formulas 
 (KCl)n, (AgCl)ni. An argument supporting this view is afforded by the existence of 
 different modifications of chloride and bromide of silver ; these differ from each other in 
 their external properties, and their varying susceptibilities to light. The doubling of 
 formulas, as .shown above with CU2CI2, Hg2Cl2, etc., is mainly due to the tendency to 
 deduce all the compounds of an element from a constant value, according to the doctrine 
 of con.stant valence. This is, however, impossible (p. ifp, 322). According to present 
 notions of valence, and as it is presented in the periodic sy.stem, compounds (MeCl, 
 MeCl2, MeCl.j, etc. ) are constituted according to definite forms or types that may materi- 
 ally determine their properties (p. 322). So far as the similarity of metallic compounds 
 is concerned, it is of .secondary importance whether the quantities corresponding to the 
 simple formulas, in the solid or gaseous state, do unite to larger, complex molecules 
 (compare HgCl and CU2CI2, — AICI3, A^CHg).^ and AL^Clf;, GaClj and Ga2Cl^,, — SnCl2, 
 80201^, PbCl2, etc. ). In case of the se.squioxides Me2G3 it is also immaterial whether 
 they are derived from supposed trivalent elements, or from those that are quadrivalent. 
 The .same may be remarked of the metallic compounds Me30^ ^ { MeO. 0)2Me (see 
 Spinels, 351). 
 
 The use of simple or of double formulas for the metallic compounds is therefore of no 
 special importance. 
 
 Silver Chloride, AgCl, exists in nature as horn -silver. When hydro- 
 chloric acid is added to solutions of silver salts, a white, curdy precipitate 
 separates ; the same fuses at 490° to a yellow liquid, which solidifies to a 
 horn-like mass. The chloride is insoluble in dilute acids; it dissolves 
 somewhat in concentrated hydrochloric acid and in sodium chloride, 
 
340 
 
 INORGANIC CHEMISTRY. 
 
 very easily in ammonium hydroxide, i)otassium cyanide, and sodium 
 hyposulphite. It crystallizes from ammoniacal solutions in large, regular 
 octahedra. Dry silver chloride absorbs ammonia gas, forming a white 
 compound of the formula 2AgC1.3NH3, which at 38° gives up its 
 ammonia. 
 
 Silver Bromide, AgBr, is precipitated from silver salts by hydro- 
 bromic acid or soluble bromides. It has a bright yellow color and dis- 
 solves with more difficulty than the chloride in ammonium hydroxide; 
 in other respects it is perfectly similar to the latter. Heated in chlorine 
 gas it is converted into silver chloride. 
 
 Silver Iodide, Agl, is distinguished from the chloride and bromide 
 by its yellow color, and its insolubility in ammonia, bused silver iodide 
 at first solidifies in isometric crystals, which gradually change to hexag- 
 onal forms, but when the latter are heated to 146°, they suddenly revert to 
 the isometric forms. It dissolves readily in hydriodic acid, to Agl. HI, 
 which, upon evaporation of the solution, separates in shining scales. 
 Heated in chlorine or bromine gas, it is converted into chloride or 
 bromide; conversely, chloride and bromide of silver are converted into 
 silver iodide by the action of hydriodic acid. 
 
 These opposite reactions are explained by the principle of the greatest evolution of 
 heat. Chlorine and bromine expel iodine from all iodides because the heat of forma- 
 tion of the latter is less than that of the bromides and chlorides. Again, hydriodic 
 acid (gaseous or in aqueous solution) converts silver chloride into the iodide according to 
 the equation : 
 
 AgCl 4- HI = Agl -f- IlCI, 
 
 because the heat modulus of the reaction is positive (for gaseous halogen hydrides 
 T12.5 Cal., for the .solution -(-10.6 Cal. See the Table at close of book). Yet it may 
 be noted here that Julius found that silver chloride and iodide were changed to bromide 
 upon heating them in bromine vapor, and further that the chloride and bromide were 
 also changed by iodine vapor into silver iodide, d'his is evidently an example of mass 
 action [Z. f. anal. Ch. (1883) 22, 523; also Blau, Monatsheft 17 (1896), 547]. 
 
 Sunlight, and also other chemically active rays (magnesium light, 
 phosphorus light) color silver chloride, bromide, and iodide at first 
 violet, then gray-black, whereby they are probably converted into com- 
 pounds of the form Ag^X. Pure silver iodide is rather non-sensitive to 
 light, but exceedingly sensitive if it contains silver nitrate or substances 
 which can take it up (e.g. tannin); on this depends the application of 
 these salts in photography. 
 
 When silver plates, previously exposed to iodine, bromine or chlorine vapors, and con- 
 sequently covered with a thin layer of a silver halide, are placed in a camera obscura in- 
 visible images are produced. Mercury vapors condense on the spots which were exposed 
 to the light and visible pictures appear (Daguerreotype, 1839). This was the first photo- 
 graj)hic process but it has been replaced by others: (i) Wet collodion process. A 
 glass plate is covered with collodion (a .solution of pyroxylin iii an ethereal solution of 
 alcohol) holding in solution cadmium iodide and ammonium iodide together with about 
 one-fourth j)art of ammonium bromide. After the evaporation of the ether the glass 
 plate is immersed in a solution of silver nitrate, whereby silver iodide and bromide are 
 ])recipitated uj)on the surface. I'he j)late thus prepared is expo.sed to light in the camera 
 obscura, and, after the action, dipi)ed into a .solution of pyrogallic acid or ferrous sul- 
 phate. 'fhese reducing substances separate metallic silver in a finely divided state, which 
 
SILVER. 
 
 341 
 
 is precipitated upon the places where the light has acted and the picture, before invisible, 
 is “ developed.” The plate is now introduced into a solution of potassium cyanide or 
 sodium hyposulphite, which dissolves the silver salts not affected by the light, while the 
 metallic unaltered silver remains (fixing). (2) Broviide-gelati}ie />rocess. The plates are 
 coated with silver bromide, finely distributed in the molten gelatine. When the layers 
 have solidified the dry plates can be preserved for a long time. The action of the light 
 produces silver sub-bromide which is more rapidly reduced by alkaline pyrogallol, hydro- 
 quinone, hydroxylamine, by potassium ferrous oxalate and eikonogen (amido-/?-naphthol- 
 /^-monosulphonate of sodium) than unaltered silver bromide, which in this case is also 
 removed by sodium hyposulphite. The negative thus formed is covered at the places 
 upon which the light shone by a dark layer of silver, while the places corresponding to 
 shadows of the received image are transparent. The copying of the glass negative on 
 paper sensitized with silver nitrate is executed in a similar manner. 
 
 Silver Nitrate, (A rgenfum nitricmn'), is obtained by dissolving 
 
 pure silver in dilute nitric acid, and crystallizes from its aqueous solution 
 in large rhombic tables, isomorphous with potassium saltpeter. At 
 ordinary temperatures it is soluble in one-half part of water or in four 
 parts of alcohol, the solution has a neutral reaction, and in this respect 
 differs from the salts of almost all metals, which react acid (p. 330). 
 It fuses at 200°, and solidifies to a crystalline mass. When perfectly pure 
 it is not affected by light. Organic substances turn it black. Silver 
 nitrate is employed for cauterizing wounds {lunar caustic). 
 
 By dissolving work-silver in nitric acid a mixture of silver and copper nitrates is 
 obtained. To separate the silver salt from such a mixture it is heated to redness, the 
 copper being thus converted into oxide and the unaltered silver nitrate extracted with 
 water. 
 
 Silver Nitrite, AgN02, is precipitated from concentrated silver nitrate 
 solutions by potassium nitrite. It crystallizes in needles, dissolves with 
 difficulty in water, and decomposes above 90°. 
 
 Silver Sulphate, Ag.^SO^, is obtained by the solution of silver in hot sulphuric acid, 
 and crystallizes in small rhombic prisms which are difficultly soluble in water. It is 
 isomorphous with anhydrous sodium sulphate. 
 
 Silver Sulphite, AggSOg, is precipitated as a white, curdy mass, if sulphurous acid 
 be added to the solution of the nitrate. It blackens in the light and decomposes at 100°: 
 
 2 Ag 2 S 03 = Ag^SO^ A Ag2 + SO. 2 . 
 
 Silver Nitride, AgNg, is very similar to silver chloride. It is more stable towards 
 light and when heated or struck explodes with great violence. Angeli claims that it can 
 be readily obtained by adding a saturated aqueous solution of hydrazine sulphate to a 
 concentrated aqueous solution of silver nitrite (Ber. 26 (1893) ill, 885 ; see pp. 132, 324). 
 
 Silver Sulphide, Ag2S, occurs in regular octahedra, as argentite. 
 Hydrogen sulphide precipitates it as a black amorphous sulphide from 
 silver solutions. By careful ignition in the air it is oxidized to silver 
 sulphate. It is insoluble in water and ammonium hydroxide. 
 
 Silver Disulphide, Ag2S2, is a dark-brown amoridious powder. It is produced on 
 mixing a .solution of silver nitrate in benzonitrile with one of sulphur in carbon bisulphide. 
 Dilute hydrochloric acid converts it into a mixture of sulphur and silver chloride, with 
 evolution of hydrogen sulphide : 
 
 I 2IICI = zAgCl -f H2S -]- S. 
 
342 
 
 INORGANIC CHEMISTRY. 
 
 Silvering. — When silver contains more llian 15 per cent, of copper it has a yellowish 
 color. To impart a pure white color to objects made of such silver they arc healed to 
 redness with acce.ss of air. 'I'he copper is thus superticially oxidized, and may be removed 
 by dilute sulphuric acid. 'I'he surface of pure silver is then polished. 
 
 'I'he silvering of metals and alloys ((German silver, argentan ) is executed in a dry or 
 wet way. In the first, the objects to be silvered are coated with liipiid silver amalgam, 
 with a brush, and then heated in an oven ; the mercury is volatilized, and the silver 
 surface then ]X)lished. 
 
 At present, the galvanic process has almost comjdetely superseded the other i)rocesses. 
 It depends on the electrolysis of the .solution of the double cyanide of silver and potas- 
 sium (AgCN.KCN), whereby the silver is thrown out at the electro-negative pole and 
 deposits upon the metallic surface in connection with that electrode. 
 
 'I'o silver glass, cover it with a mixture of anammoniacal .silver .solution, with reducing 
 organic substances like aldehyde, milk-sugar, and tartaric acid. Under definite condi- 
 tions, the reduced silver deposits upon the gla.ss as a regular metallic mirror. 
 
 Recognition of Silver Compounds. — Hydrochloric acid throws 
 down a white, curdy precipitate of silver chloride, which dissolves readily 
 in ammonium hydroxide. Zinc, iron, copper, and mercury throw out 
 metallic silver from solutions of silver salts, and from many insoluble 
 compounds, like the chloride. 
 
 GOLD. 
 
 Au = 197.2. 
 
 Gold {auriini) usually occurs in the native state, and is found dissemi- 
 nated in veins in some of the oldest rocks. Gold sands are formed by the 
 breaking and disintegration of these. It occurs, in slight quantity, in 
 the sand of almost every river. Combined with tellurium it forms sylvan- 
 ite, found in Transylvania and California. It is present in minute quan- 
 tity in most varieties of pyrites and in many lead ores. For the sepa- 
 ration of the gold grains the sand or pulverized rocks are washed with 
 running water, which removes the lighter particles and leaves the specific- 
 ally heavier gold. 
 
 The method of MacArthur and Forrest has of late found extensive appli- 
 cation in the extraction of gold from its ores. The latter after reduction 
 to a coarse powder and roasting are extracted with potassium cyanide 
 solution. The gold dissolves as potassium aurocyanide and is precipitated 
 with metallic zinc : 
 
 2Au + 4KCN + H.p -f O = 2 KAu(CN)., + 2KHO. 
 
 The following localities are important gold producers: United States, 
 Australia, Russia, South Africa, and the Klondyke in Alaska. 
 
 Native gold almost invariably contains silver, copper, and various other 
 metals. To remove these, the gold is boiled with nitric or concentrated 
 suliihuric acid. The removal of the silver by the latter acid is only com- 
 I)lete if that metal jiredominates ; in the reverse case a portion of it will 
 remain with the gold. Therefore, to sejiarate pure gold from alloys poor 
 in silver they must first be fused with about three-fourths their weight of 
 
GOLD. 
 
 343 
 
 the latter metal. Gold may be separated from coi)per and lead by cnpel- 
 lation (p. 337). 
 
 Pure gold is rather soft (almost like lead) and has a specific gravity of 
 19 32. It is the most ductile of all metals, and may be drawn out into 
 extremely fine wire and beaten into thin leaves, which transmit green 
 light. At about 1060° it meltsto a greenish liquid. It is not altered by 
 oxygen, even upon ignition ; acids do not attack it. It is only in a mix- 
 ture of nitric and hydrochloric acids (aqua regia), which yields free 
 chlorine, that it dissolves to gold chloride, AuClg Free chlorine pro- 
 duces the same. Most metals, and many reducing agents (ferrous sulphate, 
 oxalic acid) precipitate gold from its solution as a dark-brown powder. 
 
 As gold is very soft it wears away rapidly, and is, therefore, in its prac- 
 tical applications, usually alloyed with silver or copper, which have 
 greater hardness. The alloys with copper have a reddish color, those 
 with silver are paler than pure gold. The German, French, and Ameri- 
 can gold coins contain 90 per cent, of gold and 10 per cent, of copper. 
 A 14-karat gold is generally employed for ornamental objects; this con- 
 tains about 58.3 per cent, of pure gold (24 karats representing pure gold). 
 
 Gold, according to its atomic weight, belongs to the group of copper 
 and silver; and, on the other hand, forms the transition from platinum 
 to mercury. Its character is determined to a high degree by these double 
 relations (p. 328). Like the other elements of high atomic weight, mer- 
 cury, thallium, lead, and bismuth, belonging to the same series of the pe- 
 riodic system, it varies considerably in character from its lower analogues. 
 
 Gold, like silver and copper, yields compounds of the form AuX, 
 aurous, analogous to the cuprous and argentous. Besides, it has those of 
 the form AuXg, auric derivatives, in which it is trivalent. These show 
 the typical character of the trivalent combination form, which expresses 
 itself in the acidity of the hydroxides (p. 330); auric hydroxide, 
 Au(OH) 3, unites almost solely with bases. On the other hand, they 
 show many similarities to the highest combination forms of the metals 
 with high atomic weight: platinum (PtXJ, mercury (HgX2), thallium 
 (TIX3), and lead (PbXJ (p. 357). 
 
 AUROUS COMPOUNDS. 
 
 Aurous Chloride, AuCl, is produced by heating auric chloride, 
 AuClg, to 180°, and forms a white powder insoluble in water. When 
 ignited, it decomposes into gold and chlorine; cold water decomposes it 
 slowly, more quickly on heating with formation of auric chloride and gold. 
 
 Aurous Iodide, Aul, separates as a yellow powder, if potassium 
 iodide be added to a solution of auric chloride : 
 
 AUCI3 + 3KI = Aul T L 4 - 3KCI. 
 
 When heated it breaks up into gold and iodine. 
 
344 
 
 INORGANIC CHEMISTRY. 
 
 When auric oxide or sulpliicle is dissolved in potassium cyanicU', largo 
 colorless prisms of the double cyanide, AuCN.KCN, crystallize out iijx)!! 
 evaporation. The galvanic current and many metals ])recipitate gold 
 from this compound ; hence it serves for the sej)aration of gold from 
 ores and for electrolytic gilding, which, at })rescnt, has almost entirely 
 superseded the gilding in the dry way (see ]). 342). 
 
 Aurous Oxide, Au^O, is formed by the acticm of potassium hydroxide 
 upon aurous chloride. It is a dark-violet j)owder which at 250° decom- 
 poses into gold and oxygen. It is changed to auric chloride and gold 
 by the action of hydrochloric acid. 
 
 Only a few double salts of the oxygen derivatives of univalent gold are 
 known. 
 
 AURIC COMPOUNDS. 
 
 Auric Chloride, AuClg, results from the action of chlorine upon the 
 metal. It is a lemon-yellow deliquescent mass, which dissolves readily in 
 alcohol and in ether. On evaporating the solution long, yellow-colored 
 needles of the composition HAuCl^. qli/), hydrochlor-auric acid, remain. 
 It forms beautifully crystallized double salts with many metallic chlorides, 
 g., KAuCl^ -}- 2H2O and NH^AuCl^ -f- 2^H20. When a solution of 
 auric chloride is heated with magnesium oxide a brown precipitate is 
 obtained, from which all the magnesia is removed by concentrated nitric 
 acid, leaving Auric Oxide (AiqOg). This is a brown powder which 
 decomposes, near 250°, into gold and oxygen. If the preci})itate con- 
 taining the magnesia be treated, not with concentrated, but with dilute 
 nitric acid, Auric Hydroxide, Au(OH)g, remains as a yellowish-red 
 powder: Both the oxide and hydroxide are insoluble in water and acids; 
 they possess, however, acid properties, and dissolve in alkalies. There- 
 fore the hydroxide is also called auric acid. Its salts, the aurates, are 
 
 I 
 
 constituted according to the formula MeAuOg, and are derived from the 
 meta-acid, HAuOg = AuO. OH. 
 
 Potassium Aurate, KAuOg + 3H2O, crystallizes in bright yellow 
 needles, from a potassium hydroxide solution of auric oxide. These are 
 readily soluble in water ; the solution reacts alkaline. The corresponding 
 aurates are precipitated from this solution by many metallic salts, e. g. ; 
 
 KAuOg + AgNOg = AgAuO^ + KNOg. 
 
 The precipitate produced by magnesia in a solution of auric chloride 
 (see above) consists of magnesium aurate, Mg(Au02)2. ’ Oxygen salts of 
 auric oxide are not known. 
 
 Auric Sulphide, AiqSg, is preci])itated as a blackish-brown com- 
 ]>ound, from gold solutions, by liydrogen sulphide. Its composition is 
 either AiqSg or AU2S2 (mixed with sulphur), depending on the condi- 
 tions prevailing at the time. It dissolves in alkaline sulphides with forma- 
 tion of sulj)ho-salts, which are only derived from Au2S(<f. ^., KgAuSj). 
 Hydrogen sulphide })reci})itates aurous sulphide, AiqS, from hot gold 
 
METALS OF GROUP III. 345 
 
 solutions. It is steel-gray in color, and soluble in pure water, from 
 which hydrochloric acid precipitates it. It is used in gold plating. 
 
 Stannous chloride (SnClJ added to an auric chloride solution pro- 
 duces, under certain conditions, a purple-brown precipitate, purple of 
 Cassius, which is employed in glass and porcelain painting. Alumina 
 and magnesia yield similar purples, and it appears that their red colora- 
 tion is due to finely divided metallic gold. 
 
 On pouring ammonium hydroxide over auric oxide a brown compound 
 is produced — fubninating gold. When this is dried and heated or struck 
 it explodes very violently. 
 
 METALS OF GROUP III. 
 
 The trivalent elements, affording derivatives mainly of the form MeX3, 
 belong to group 111 of the periodic system (p. 246) : 
 
 Sc = 44.1 Y 89 La = 138 Yb = 173 
 
 B =:r II A 1 = 27.1 
 
 Ga =go In ^-114 ... T 1 = 203.7. 
 
 These bear the same relation to one another as do the elements of 
 group II (p. 299). Boron has the lowest atomic weight, and the basic, 
 metallic character in it is reduced very much or does not appear at all. 
 In its exclusively acidic hydroxide, B(OH)3, it approaches the metal- 
 loids, and is therefore treated with them (p. 241). 
 
 Aluminium is a perfect metal; its hydroxide, Al(OH)3, exhibits a pre- 
 dominating basic character, and yields salts with acids. Its relations to 
 boron are like those of silicon to carbon, or of magnesium to beryllium. 
 The connection of aluminium and boron with the same group plainly 
 shows itself in the entire character of the free elements, and in their 
 compounds. Thus aluminium and boron are not dissolved by nitric acid, 
 but by boiling alkalies : 
 
 A 1 -p 3KOH = A1(0K)3 + 3H. 
 
 There is only a gradual difference between their hydrates. Boron 
 hydroxide, B(OH)3, not only acts as a feeble acid, but we also find that 
 aluminium hydroxide manifests an acidic character, inasmuch as it is 
 capable (pp. 258, 330) of forming metallic salts with strong bases (chiefly 
 the alkalies) ; but owing to the higher atomic weight of aluminium the 
 basic character exceeds the acidic. 
 
 Scandium, yttrium, lanthanum and ytterbium attach themselves to 
 aluminium as the first sub-group. These constitute the third members of 
 the great periods, and hence exhibit a pronounced basic character. As 
 light metals, they are very similar to aluminium in their compounds, so 
 that they are all embraced in one group, which (corresponding to the 
 earthy nature of their oxides) is designated the group of earth metals. 
 Cerium, praseodymium and neodymium bear a peculiar relation to lan- 
 thanum ; their atomic weights are nearly alike and their properties very 
 
346 
 
 INORGANIC CHEMISTRY. 
 
 similar. They, like erbium, terbium, llmlium, holmium aud samarium, 
 are by no means so accurately known that they can l)e assigned definite 
 ])laces in the periodic system. Most of them are i)robably mixtures of 
 unknown elements. 
 
 The second sub group is more distinctly characterized and accurately 
 investigated; it consists of the heavy metals, gallium, indium and thal- 
 lium. 'I'hese belong to the right side of the great periods, ])ossess, there- 
 fore, a less basic character, and bear the same relation to one another as 
 zinc, cadmium and mercury. 
 
 I. GROUP OF THE EARTH METALS. 
 
 ALUMINIUM. 
 
 A1 = 27.1. 
 
 Aluminium, both in quantity and distribution, is one of the most 
 important constituents of the earth’s crust. As oxide, it crystallizes as 
 ruby, sapphire and corundum ; less pure as emery. It is commonly 
 found as aluminium silicate (clay, kaolin), and in combination with other 
 silicates, as felds])ar, mica, and also in most crystalline rocks. It occurs, 
 too, united with fluorine and sodium, as cryolite, AlFl3.3NaFl, in large 
 deposits, in Greenland. 
 
 Metallic aluminium is at present obtained almost exclusively by the 
 electrolysis of alumina dissolved in fused cryolite. The oxygen burns 
 the carbon anode to carbon monoxide and aluminium separates in the 
 molten condition. 
 
 Metallic aluminium was first obtained as a gray powder by Wohler in 1827 when he 
 heated the chloride with potassium. Later he conducted the vapors of aluminium chlo- 
 ride over heated sodium or potassium, obtaining thereby globules of the metal. Bunsen, 
 1854, reduced the double chloride of sodium and aluminium with the electric current and 
 separated aluminium also in powder form. St. Claire-Deville announced similar methods 
 almost simultaneously. It is due to his efforts that the metal was produced on a com- 
 mercial scale (1856). This he accomplished by reducing aluminium sodium chloride, 
 partly with introduction of cryolite, by means of sodium : 
 
 AlCb.NaCl + 3Na = 4NaCl -f Al. 
 
 However, in 1888 this method was in use in but one factory. A new impulse was given 
 to the manufacture of the metal after the di.scovery of Ca.stner’s and of Netto’s methods of 
 preparing metallic .sodium. Equal parts of cryolite and sodium chloride were fused 
 together and blocks of sodium of 5-7 kilograms in weight were dipped into the then 
 molten mass, the slag, separated from the aluminium and containing .sodium fluoride, was 
 treated with aluminium sul[ihate and artificial cryolite prepared. This sodium method 
 was abandoned as a conse(|uence of the develojmient of electrolysis. In America 
 natural gas is used for heat and jiower purpo.ses ; and the power of Niagara Falls was 
 turned to the production of electric currents. A portion of the Rhine Falls at Schaff- 
 liauscn in Switzerland is applied to the same purpo.se. Another large aluminium plant 
 exists at hroges near Orenoble. 'I'he electrolysis is conducted in accordance with the 
 method of 1 Itroult, using a .solution of alumina in fused cryolite and fluorspar. T he oxide 
 breaks down (juite readily into aluminium and oxygen. The Hall method, so largely 
 
Al^UiMINlUM. 
 
 347 
 
 used in America, is practically the same. (Compare Cl. Winkler, Ch. Zt., 1892, 349, 
 and Elbs; ibid., 1894, 1637; 1897, 995; especially W. Borcher’s Elektrometallurgie, 
 1896, and his Elektrische Schmelzofen, Halle, 1897.) 
 
 Aluminium is an almost silver-white metal of strong luster, very ductile, 
 and it may be drawn out into fine wire and beaten into thin leaflets. Its 
 specific gravity is 2.583 ; it belongs, consequently, to the light metals (see 
 pp. 252,313). It fuses at 660° but will not vaporize. It changes very little 
 in the air at ordinary temperatures, and even when heated. If, however, 
 thin leaves be heated in a stream of oxygen, they will burn with a bright 
 light. Pure aluminium is attacked with difficulty by water. The commer- 
 cial article, containing sodium, is more readily affected, while the amal- 
 gamated product is very easily and quickly acted upon by the reagent. 
 Nitric acid affects aluminium only superficially ; sulphuric acid dissolves 
 it on boiling, while it is readily soluble, even in the cold, in hydrochloric 
 acid, and in a solution containing 5 per cent, of acetic acid and some 
 sodium chloride it is attacked at once. It dissolves in potassium and 
 sodium hydroxide with evolution of hydrogen, and forms aluminates: 
 
 A 1 4- 3KOH = K3 AIO3 -f 3H. 
 
 Owing to its stability in the air and beautiful luster, aluminium is some- 
 times employed for vessels and ornaments. 
 
 It has been found that vessels of aluminium are attacked by most foods and drinks. 
 The effect is slight and diminishes rapidly with continual use. The relish of foods pre- 
 pared in such vessels appears to suffer nothing by continuance nor does there appear to 
 be any harmful influences to the general health (Plagge). 
 
 Alloys of copper with 5-12 per cent, of aluminium are distinguished 
 by their great hardness and durability. They make good castings, and 
 possess a gold-like color and luster. Under the name of aluminium 
 bronze, they are used for the composition of various articles, as watches, 
 spoons, etc. Their firmness and elasticity render them suitable for physical 
 instruments (arms of balances) and watch springs. 
 
 Aluminium has a most remarkable affinity for oxygen. It has conse- 
 quently been in use for quite a while in isolating other elements from their 
 oxides (p. 136). Magnesium oxide is the only oxide not reduced by 
 metallic aluminium. Alumina (corundum) always results in such reduc- 
 tions : 
 
 Cr203 -p 2AI = AI2O3 + Cr2. 
 
 When metallic oxides are heated with aluminium powder so much heat is liberated 
 that violent explosions take place. II. Gold.schmidt has recently discovered a means of 
 controlling these rearrangements and of utilizing the immense heat of the reaction by ignit- 
 ing the mixture of metallic oxide and aluminium powder within by means of a percussion 
 point at but one place and letting the transpositions advance from there. The “ percus- 
 sion point” consists of a magnesium band supplied at its end with a mixture of aluminium 
 powder and barium peroxide or some other sub.stance rich in oxygen. In such reactions 
 the aluminium is converted into a slag of fused corundum, which on cooling crystallizes 
 in part like the natural product. 
 
348 
 
 INORGANIC CHKMISTRV. 
 
 Sulphides iiisiciul of oxides, can be used ; tliis is sonietitnes advantageous because 
 aluuiiniuin sulphide melts easily and can be readily separated from melals : 
 
 3CuS 4 2AI = AlgSg -f 3CU. 
 
 [See II. Goldschmidt, Ann. Chem. (1898) 301, 19.] 
 
 d'he heat evolved in the formation of aluminium hydroxide is 388.8 Cal. This far 
 exceeds that of any other oxides : 
 
 Cu,0 = 37 ; Fe„03,3llg0 = 191. 
 
 ddie heat corresponding to one atom of oxygen, is, therefore, 129.6 Cal.; since that of 
 water is far less (11.2,0 — 69.0), it must be decompo.sed by aluminium, with liberation 
 of hydrogen (p. 273). If this does not transpire under ordinary conditions, the reason 
 must be sought for in the insolubility of aluminium hydroxide. Indeed, the reaction 
 occurs if aluminium chloride, or another salt, in which the aluminium oxide is soluble, 
 be added to the water. Conversely, the high heat of formation of aluminium oxide 
 explains why it is not reduced by carbon. 
 
 Aluminium unites in but one proportion with oxygen, Al^Og; this is 
 the source of salts of the type AIX^ (p. 349). Water decomposes them 
 hydrolytically in part and that explains the acid reaction of their aqueous 
 solutions; they have a sweet, acrid taste. 
 
 The trivalence of aluminium is shown in its organic derivatives, e.g. : 
 
 A1(CH3)3, A\{C,H,)„ etc. 
 
 Such compounds like those with hydrogen are best adapted for the determination of the 
 valence of an element (p. 247). At sufficiently high temperatures the density of the 
 vapors of- aluminium chloride correspond to the formula AlCl,, but at lower temperatures 
 this, like other metallic chlorides (<?. g., tin dichloride, gallium chloride, etc., p. 339), 
 shows a tendency to form larger molecules — to polymerize. 
 
 Aluminium Chloride, AICI3, is produced by the action of chlorine 
 upon heated aluminium; also by heating a mixture of aluminium oxide 
 and carbon in a current of chlorine: 
 
 AlA + 3C -p 6C1 = 2 AICI 3 + 3CO. 
 
 Chlorine and carbon do not act separately upon the oxide ; by their 
 mutual action, however, the reaction occurs in consequence of the affinity 
 of carbon for oxygen, and of chlorine for aluminium. The oxides of 
 boron and silicon sliow a similar deportment. 
 
 In order to get chloride free from iron and oxychloride, it is advisable 
 to heat the metal in dry hydrochloric acid gas (Seubert and Pollard, Ber. 
 (1891) ?4, 2575). 
 
 Aluminium chloride may be obtained in white, hexagonal leaflets by 
 sublimation. It sublimes readily, but will only fuse when subjected to 
 high ])ressure (in a sealed tube) and at 193-194°. At 175-179° the 
 fused, cooling mass exhibits the jihenomenon of boiling and subliming 
 (j). 229). The boiling point under the ordinary pressure is 183°. 
 
ALUMINIUM. 
 
 349 
 
 Deville and Troost (1857) first determined the vapor density of aluminium 
 chloride, bromide, and iodide at 440°, and found it to correspond to the 
 formulas Al2Clg, Al2Brg, and Later, Friedel and Crafts redetermined 
 
 that of the chloride at 218-440° with like results. This led to the sup- 
 position that aluminium was quadrivalent in all its compounds (Jahresb. 
 f. Ch. 1888, I, 131). The most recent investigations of Nilson and 
 Peterson (Z. f. phys. Ch. i (1887), 459) have shown that at higher 
 temperatures (above 760°) the vapor density of the chloride corresponds 
 to the formula AICI3 (p. 339). 
 
 Aluminium chloride absorbs moisture from the air, and deliquesces. 
 It crystallizes from a concentrated hydrochloric acid solution, with six 
 molecules of water. On evaporating the aqueous solution, the chloride 
 decomposes into aluminium oxide or hydroxide and hydrogen chloride : 
 
 2AICI3 + 3H2O = AI2O3 + 6 HC 1 . 
 
 It forms double chlorides with many metallic chlorides, viz. : AICI3. NaCl, 
 AICI3.KCI. The solutions of these may be evaporated to dryness with- 
 out decomposition. It also unites with many halogen derivatives of the 
 metalloids : 
 
 A 1 C 13 . PCI., AICI3 . POClg, AICI3 . SCh. 
 
 Aluminium Bromide, AlBrj, is obtained like the chloride, and consists of shining 
 leaflets which fuse at 90° and boil at 265-270°. 
 
 Aluminium Iodide, AII3, is formed on heating aluminium filings with iodine. It 
 is a white, crystalline mass, fusing at 185°, and boiling at 360°. It is best prepared by 
 covering sheet aluminium with carbon bisulphide, and then adding the calculated amount 
 of iodine gradually, letting the whole stand for some time, and then distilling off the car- 
 bon bisulphide. The reaction occurring between aluminium iodide and oxygen is inter- 
 esting. If the vapor of the former be mixed with the latter, and then brought in contact 
 with a flame, or if acted upon by an electric spark, a violent detonation will ensue ; 
 aluminium oxide and iodine result : 
 
 2AII3 3O = AI2O3 -f- 312- 
 
 This deportment is due to the great difference in the heats of formation of the aluminium 
 oxide (about 380 Cal.), and the iodide (140.6 Cal.). The chloride and bromide are 
 similarly decomposed, but with less violence. 
 
 Aluminium Fluoride, AIFI3, obtained by conducting hydrogen fluoride over heated 
 aluminium oxide or hydroxide, sublimes at a red heat in colorless rhombohedra. It is 
 insoluble in water, unaltered by acids, and is very stable. It yields insoluble double 
 fluorides with alkaline fluorides. The compound, AlFl3.3NaFl, occurs in Greenland, 
 in large deposits, as cryolite, and is employed in the soda and alumina manufacture 
 (p. 292). 
 
 Aluminium Oxide, AI2O3, is found crystallized in hexagonal prisms 
 in nature, as ruby, sapphire, and corundum, colored by other admixtures. 
 Impure corundum, containing aluminium and iron oxides, is called 
 emery, and .serves for polishing glass. The specific gravity of these 
 minerals is 3.9 ; their hardness is only a little below that of the diamond. 
 Artificial aluminium oxide — corundum — may be obtained by reducing 
 metallic oxides with aluminium — following Goldschmidt’s method; in 
 making chromium the slag will contain little crystals of the ruby. The 
 crystallized or strongly ignited aluminium oxide is almost insoluble in 
 
350 
 
 INORGANIC CHEMISTRY. 
 
 acids; to decompose it, it is fused with caustic alkalies or with ])rimary 
 potassium sulphate, flKSO^. 
 
 Aluminium Hydroxides. — The fior?nal hydroxide, Al(OH)3, occurs 
 in nature as hydrargillite. The hydroxide, Al202(()H)^, is diaspore. 
 Bauxite is a mixture of the hydroxide, Al20((i)H)^, with ferric oxide. It 
 occurs in large dei)osits in France, in the Caucasus, etc. The normal 
 hydroxide is artificially obtained as a white voluminous i)recipitate, 
 by adding ammonium hydroxide or an alkaline carbonate (in the latter 
 case carbon dioxide escapes, p. 352) to a soluble aluminium salt. It 
 dissolves, when freshly precipitated, in acids and in potassium and 
 sodium hydroxides, liy long standing under water, or after drying, 
 it is, without any alteration in comi)osition, difficultly soluble in 
 acids. When carefully heated, the normal hydroxide first passes into 
 the hydrate, AlO.OH. 
 
 The freshly preeijiitated hydroxide dissolves readily in a solution of aluminium chloride 
 or acetate. On dialyzing (p. 237) this solution the aluminium salt or crystalloid diffuses, 
 and in the dialyzer remains the pure acjueous solution of the hydroxide. This has a faint 
 alkaline reaction and is coagulated by slight quantities of acids, alkalies and many salts ; 
 the soluble hydrate passes into the insoluble gelatinous modification. 
 
 Gelatinous aluminium hydroxide possesses the property of precipitating many dyestuffs 
 from their solutions, forming colored insoluble compounds (lakes) with them. On this 
 is based the application of aluminium hydroxide as a mordant in dyeing. The acetate is 
 generally used for this purpose. Goods saturated with this salt are heated with steam, 
 which causes the decomposition of the weak acetate ; acetic acid is driven off, while 
 the separated aluminium hydroxide adheres to the fiber of the material. If the latter 
 now be introduced into the solution of coloring matter the color will be fixed by the 
 aluminium hydroxide upon the fiber. At present, sodium aluminate is employed in- 
 stead of the acetate. 
 
 Aluminium hydroxide has a feeble acid character, and can form salt- 
 like compounds with strong bases. On carefully evaporating its solution 
 in sodium or potassium hydroxide, or upon addition of alcohol, white 
 amorphous compounds of KAIO2, NaA102 and Al(ONa)3 are obtained. 
 The potassium compound can be obtained in crystalline form. These 
 derivatives, known as aliwiinates, are not very stable, and are even 
 decomposed by carbon dioxide, with elimination of aluminium hydroxide : 
 
 2NaA10., 4 - CO2 + 3H,0 = 2A1(0H)3 + 
 
 d'he aluminium hydroxide obtained in this manner, in distinction from 
 that precipitated from acid aluminium solutions by the alkalies, is not 
 gelatinous, and is more difficultly soluble in acids, especially acetic. It 
 comprises the ordinary alumina of commerce. 
 
 On adding calcium chloride, strontium chloride, or barium chloride to 
 the solution of potassium or sodium aluminate, white insoluble aluminates 
 are precipitated : 
 
 2NaA102 4- CaCl2 = Ca(A102)., 4 - 2NaCl. 
 
 Similar aluminates fre(piently occur as crystallized minerals, in nature. 
 Thus the spinels consist chiefly of magnesium aluminate. 
 
ALUMINIUM. 
 
 351 
 
 chrysoberyl is beryllium aluminate, gahnite is zinc alu- 
 
 . , AlO.O. .7 
 
 minate, ^10.0^^^^' 
 
 Nearly all these minerals, commonly called spinels, crystallize in regu- 
 lar octahedra, like the corresponding chromium compounds (see these); 
 the exceptions are chrysoberyl, crystallizing in the rhombic system, and 
 
 hausmannite, = MnO' quadratic system. 
 
 Technically, alumina is obtained from cryolite, bauxite and other minerals containing 
 aluminium. The pulverized bauxite is heated ’with dry sodium carbonate in furnaces, 
 and the resulting sodium aluminate extracted with water. P'rom the clear solution car- 
 bon dioxide precipitates the hydroxide, while sodium carbonate remains dissolved, and is 
 afterward recovered. The dried aluminium hydroxide occurs as a white powder in trade. 
 
 K. j. Baeyer treats the finely ground bauxite with caustic soda and stirs into this solu- 
 tion freshly precipitated, crystalline alumina (as obtained on conducting carbon dioxide 
 into a cold aluminate solution). All of the alumina separates after some hours, and the 
 alkaline liquid is again run through the same course. The method is based on the decom- 
 position of the aluminates by water ; the alumina which is mixed in with it acts like a 
 crystal of Glauber salt upon supersaturated solutions (p. 288). 
 
 The gelatinous, readily soluble (colloidal) aluminium hydroxide (see above) precipi- 
 tated from acid solutions by alkalies, has lately been prepared upon a large scale, accord- 
 ing to the method of Lowig, by treating the sodium aluminate solution with milk of lime ; 
 calcium aluminate precipitates, while sodium hydroxide remains in solution : 
 
 2NaA102 -b Ca(0H)2 = Ca(A102)2 + 2NaOH. 
 
 The calcium aluminate is dissolved in hydrochloric acid : 
 
 Ca(A102)2 + 8IIC1 = 2AICI3 -f CaCq + 4H2O, 
 
 and to the solution now containing the alumina as chloride the corresponding amount of 
 calcium aluminate added, and aluminium hydroxide is precipitated : 
 
 2AICI3 4- 3Ca(A102)2 4- 12H2O == 8A1(0H)3 + 3CaCl2. 
 
 According to this procedure, the sodium hydroxide formed in the first reaction is obtained 
 together with the alumina. 
 
 On conducting carbon dioxide into a solution of alkaline carbonates, and adding a 
 solution of an alkaline aluminate at the same time, white aluminiian-alkali carbonates 
 are precipitated : 
 
 AI2O3.K2O 4- 2KHCO3 = Al203.K,0.2C02 4 2K0fI. 
 
 The caustic alkali that is formed in this way is converted again into bicarbonate by car- 
 bon dioxide. In a dry state the precipitates are white, chalk-line masses which at 90° 
 contain five molecules of water : 
 
 Al2O3.K2O.2CO2 f 5 H.A 
 
 Their constitution may be expressed by the formula : 
 
 K-O-CO-O-AI ^ O. 
 
 They dissolve readily in dilute acids, even acetic, with evolution of carbon dioxide, and 
 are suitable for the preparation of pure alumina mordants and antiseptic solutions (Lowig). 
 
352 
 
 INORGANIC CHEMISTRY. 
 
 Tlie basic cliaracter of altiniinium liydroxide exceeds the acid ; but it 
 is so feeble that it is not ca])able of forming salts with weak acids, as car- 
 bon dioxide, suliihnrons acid, and hydrogen sulphide. When sodium 
 carbonate is added to solutions of aluminium salts, aluminium hydroxide 
 is precipitated, while carbon dioxide is set free : 
 
 2AICI3 + sNa^COj + 311,0 = 2AI (011)3 I 6NaCl + 3CO,. 
 
 The alkaline sulphides behave similarly : 
 
 2AICI3 -1 3(Nnd,S + 611,0 = 2A1(0II)3 + 6NIbCl -|- 3H,S. 
 
 That the hydrate is a feeble acid is shown in the decomposition of the 
 alkali salts by carbon dioxide. 
 
 Aluminium Sulphate, Al2(SOj3, crystallizes from aqueous solution 
 with eighteen molecules of water in thin leaflets with ])early luster. These 
 dissolve readily in water; when heated, they melt and lose all their water 
 of crystallization. The sulphate is obtained by dissolving the hydroxide 
 in sulphuric acid, or by the decomposition of pure aluminium silicate 
 (clay) with the same acid ; the residual silicic acid is removed by filtra- 
 tion, and the solution of the sul])hate evaporated. When a quantity of 
 ammonium hydroxide, insufficient for complete precipitation, is added 
 to the sulphate, basic sulphates separate out. Salts similar to the latter 
 are also found in nature; thus, aluminite, used to prepare alum, has the 
 composition : 
 
 + 7H3O or (A 10 . 0 ),S 0 , + 9 H, 0 . 
 
 Aluminium sulphate can combine with the alkaline sulphates and 
 affords double salts, termed alu??is, e. g., potassium alum: 
 
 Al2(SOj3.K,SO, q 24H,0 or KAl(SOd, + I 2 H, 0 . 
 
 Their constitution is expressed by the following formula : 
 
 III 
 
 SO^ = A 1 — SO, — K + i 2 H, 0 . 
 
 In this compound the potassium may be replaced by sodium, ammonium, 
 rubidium, Ccesium, and also by thallium hydroxylamine, and some organic 
 bases. Iron, chromium and also vanadium, titanium and manganese 
 afford like derivatives: 
 
 Fe,fS0,)3.K,S0, + 24H,0 Cr,(SO,)3. (NT-Tp,(SO), + 2411,0. 
 
 Potassium iron alum. Ammonium chromium alum. 
 
 All these alums crystallize in regular octahedra or cubes, and can form 
 isomorphous mixtures. 
 
 'I'he most im|)ortant of them is pofassiinn aluminium sulphate or ordi- 
 nary alum, KA1(S0,)2 -|- 12! 1,0. Jt crystallizes from water in large, 
 transjiarent octahedra, soluble in 8 parts of water of ordinary tempera- 
 
ALUMINIUM. 
 
 353 
 
 tiire, or in 3<3 part of boiling water. The solution has an acid reaction and 
 a sweetish, astringent taste. When })laced over sulphuric acid, alum loses 
 three-fourths of its water. When heated it melts in its water of crystalliza- 
 tion, loses all the latter, and becomes a white, voluminous mass — burnt 
 alum. Upon adding a little sodium or potassium carbonate to a hot 
 alum solution, the hydroxide first produced dissolves, and when the liquid 
 cools, the alum crystallizes out in cubes, as cubical alimi. The addition 
 of more sodium carbonate causes the precipitation of the basic salt — 
 KA1(S0^)2. A1(0H)3. Alunite, found in large quantities near Rome and 
 in Hungary, has a similar composition, K(A\ 0 ').J SO^)^ 3H2O. 
 
 Commercial alum is obtained according to various methods: (i) From alunite, by 
 heating and extracting with hot water. In this way alum dissolves while the hydroxide 
 remains ; from such solutions the former crystallizes in combinations of the octahedron 
 with cube faces — Roman alum. (2) The most common source of alum was formerly altan 
 shale, a clay containing pyrite and peat. This is roasted and after moistening with water 
 is exposed for a long time to the action of the air. By this means pyrite, FeS2, is con- 
 verted into ferrous sulphate, FeS04, and free sulphuric acid, which, acting upon the 
 clay, forms aluminium sulphate. The mass is extracted with water, potassium sulphate 
 added, and the whole permitted to crystallize. (3) At present clay is treated directly with 
 sulphuric acid, and to the solution of aluminium sulphate potassium or ammonium sul- 
 phate is added. (4) Bauxite and cryolite are admirable material for the preparation of 
 alum. The working of cryolite for alumina and soda is described on p. 292, and that of 
 bauxite, p. 351). 
 
 Ammonium Alum, NH4A1(S04)2 4 - 12H2O, crystallizes, like potas- 
 sium alum, in large crystals, and at present, owing to its cheapness, is 
 applied almost exclusively for technical purposes. Sodium alum is much 
 more soluble, and crystallizes with difficulty. As the alum employed in 
 dyeing must contain no iron, we understand why this salt is not applica- 
 ble. At present the alum is being more and more supplanted by alu- 
 minium sulphate and sodium aluminate in all practical operations, because 
 these chemicals can be procured perfectly free from iron. 
 
 Aluminium Phosphate, AlPO^-j- 4H2O, is thrown out of alumin- 
 ium salt solutions by sodium phosphate, as a white gelatinous precipi- 
 tate; this is readily soluble in acids, acetic excepted. Turquoise is a 
 basic phosphate of aluminium colored blue by a copper compound. 
 
 Aluminium Silicates. — The most important of the aluminium 
 double silicates, so widely distributed in nature, are : leucite, KAl(Si03)2, 
 albite or soda feldspar, NaAlSi30g, ordinary feldspar — orthoclase, KAlSi30g, 
 and the various micas, which, with quartz, compose granite. When 
 these disintegrate under the influence of water and the carbon dioxide 
 of the air, alkaline silicates are dissolved and carried away by water, 
 while the insoluble aluminium silicate, clay, remains. Perfectly pure 
 clay is white, and is called kaolin, or porcelain clay ; its composition 
 mostly corresponds to the formula, Al/SiOJg -f 2H2O or Al2Si207 -}- 
 2H2O. When clay is mixed with water a tough, kneadable mass is 
 obtained. By drying and burning, it becomes compact and hard, and is 
 the more fire-proof, the purer the clay. On this depends the use of clay 
 for the manufacture of earthenware, from the red brick to porcelain. 
 
 30 
 
354 
 
 INOkCJANIC CHEMIS'IRY. 
 
 "I'o produce porcelain ii very ('me mixture of kaolin, feldspar and quartz is cinj)loyed. 
 On strong ignition, the feldspar fuses, (ills the pores of the clay and thus furnishes a fused 
 transparent mass — porcelain. When it is not so strotigly ignited, it remains porous — 
 faience — serving for liner clay vessels. 'I'o render these impervious to water, they are 
 covered wdth glazing. 'I'liis consists of various readily fusible silicates. Rough earthen- 
 ware vessels are constructed from impure clay, and they are u.sually glazed by throwing 
 salt into the ovens at the time of burning, d'he salt vaj)orizes and forms an easily 
 fusible silicate on the surface of the clay, hydrochloric acid being liberated. 
 
 Ultra 7 narine . — Tlie rare mineral Lapis lazuli, which was formerly em- 
 l)loyed as a very valuable blue color under the name of Ultrainai'ine, is 
 a compound of altiminitim sodium silicate with sodium j)olysuli)hides. 
 At present ultramarine is prepared artificially, in large quantities, by 
 heating a mixture of clay, dry soda (or sodium stilj)hate), sulphur and 
 charcoal, away from air. Green ultramarine is the product. This is 
 then washed with water, dried, mixed with powdered sulphur and gently 
 heated with air contact until the desired blue color has appeared — blue 
 ultramarine. The chemical constitution of this blue coloring material 
 has not yet been explained. On pouring hydrochloric acid over the blue 
 product, the color disapi)ears with liberation of sulj)hur and hydrogen 
 sulphide — this would point to the existence of a polysulphide. Violet 
 and red ultramarines are prepared at present by conducting dry hydrogen 
 chloride gas and air over common ultramarine at 100-150°. 
 
 RARE EARTH METALS. 
 
 In some very rare minerals, like cerite, gadolinite, euxenite, orthite, samanskite, thorite, 
 and monazite, occurring principally in Sweden, Norway, Greenland, and the United 
 States, is found a series of metals which, in their entire deportment, closely resemble 
 aluminium (p. 345). The.se are scandium, yttrium, cerium, lanthanum, neodymium, 
 praseodymium, samarium, ytterbium, and the more recent erbium, terbium, gadolinium, 
 decipium, thulium, holmium and dysprosium. These generally form difficultly soluble 
 oxalates, and. are, therefore, precipitated from solution by oxalic acid. The anhydrous 
 sulphates are quickly and easily soluble in cold water ; in warm water they are slowly and 
 difficultly soluble. They form double salts with the alkali sulphates ; the potassium 
 double salts are constituted according to the formula 1^02(804)3. 3X280.^. Their solu- 
 bility in water differs ; some of them are insoluble in saturated solutions of potassium 
 sulphate. The formates are partly sparingly soluble and in part very soluble. The 
 hydrous nitrates and chlorides are very soluble in water and in alcohol. The nitrates of 
 these metals are decomposed at different temperatures ; this property affords an excellent 
 means for their isolation and separation. Fractional precipitation with ammonia is also 
 used for this purpose. The salts and solutions of the individual metals are in part 
 colorless and ])artly colored. Most of the latter show absorption spectra. The position 
 of the ab.sorjilion lines in the spectrum — their wave-length — is characteristic for that 
 particular element and is a sure and excellent means for its recognition. The oxides of 
 many of these metals when heated in the Bunsen flame diffu.se great quantities of light, 
 d'he glowing oxides of .some of the elements whose .salt .solutions show absorption spectra 
 give a di.scontinuous spectrum with bright bands whose positions corre.spond exactly to the 
 ab.sorptioti lines. 
 
 8candium, yttrium, lanthanum, cerium, samarium, and ytterbium have been most accu- 
 rately investigated. 'Fheir atomic weights are very approximately correct. The most in- 
 teresting of the grouj) is scandium, atomic weight 44.1. It fills out the gap between 
 calcium and titanium. It coincides in all its properties with those deduced theoretically 
 from the periodic system by Mendelejeff for the element ekaboron (compare Gallium). 
 
RARE EARTH METALS. 
 
 355 
 
 Scandium, Sc = 44.1, discovered by Nilson in 1879, is contained in euxenite and 
 gadolinite ; it has not yet been obtained in a free condition. Its oxide, Sc^Oy, is obtained 
 by igniting the hydroxide or nitrate, and is a white, infusible powder (like magnesia and 
 oxide of beryllium). Its specihc gravity equals 3.86. The hydroxide, Sc(OM).j, is 
 precipitated as a gelatinous mass from its .salts by the alkalies, and is insoluble in an 
 excess of the latter. The nitrate crystallizes in little prisms, and is easily decomposed by 
 heat. The potassium double sulphate, Sc.^(SO^)3. 3K2SO4, is soluble in warm water, but 
 not in a solution of potassium sulphate. The chloride affords a characteristic spark spectrum. 
 
 Yttrium — Y = 89, discovered by Mosander in 1843 in the yttria of gadolinite, has 
 been obtained by electrolyzing the anhydrous fused chloride. It is a gray powder which 
 decomposes water in the cold. Cleve and Hoglund first obtained many of its compounds 
 in a pure state. It occurs principally in the gadolinite from ytterby (about 35 per cent. 
 Y2O3). Its oxide is a white, infusible, strong base, which is capable of expelling ammonia 
 from its salts. Its specific gravity is 5.04. Its potasshan double sulphate, Y2(S()^)3. 3X280^, 
 is soluble in a potassium sulphate solution, and in this manner it can be readily separated 
 from cerium, lanthanum, didymium and samarium. \X.^nitrate, Y(N03)3 “h forms 
 
 large, deliquescent crystals. The chloride, YCI3 -j- 6H2O, forms large prisms, and gives 
 a brilliant spark spectrum rich in lines. 
 
 Lanthanum, La = 138 (from ?iav6dveiv, to be hidden), was discovered in cerite by 
 Mosander (1839). It can be obtained by the electrolysis of its fused chloride. It resem- 
 bles iron in color and in luster. It oxidizes in the air and in a flame burns with a bright 
 light. Specific gravity = 6. 16. The oxide La203 is brilliant white ; it combines directly 
 with water to the hydrate lLa.{OH).^, which reacts alkaline and absorbs carbon dioxide from 
 the air. It is a powerful base which expels ammonia from ammonium salts in the cold. 
 The anhydrous sulphate, La2( 80^)3, dissolves readily in water of 4°. When its saturated 
 solution is heated to 40° most of it separates as the hydrous salt. The following sul- 
 phates are known : La2(804)3 -\- 6H2O and La2(804)3 -f 9H2O. The potassium double 
 salt is sparingly soluble in water and insoluble in a saturated potassium sulphate solution, 
 La2(804)3. 31^2^04. The nitrate, La(N03)3 • 6H2O, consists of large prismatic crystals, 
 which lose water on heating. The salt melts when more strongly heated, and on cooling 
 becomes like dust. It decomposes at high temperatures. The anhydrous chloride, 
 LaCl3, is a crystalline mass, which deliquesces in moist air, forming crystals of the hydrous 
 chloride, 2LaCl3 . 15H.2O. The chlo7'ide shows a very bright spark spectrum with numer- 
 ous lines. Pure lanthanum oxide exposed in a vacuum to kathode rays shines with a 
 powerful light and before the spectroscope shows a brilliant band spectrum. The lan- 
 thanum salts are colorless ; they do not have an absorption spectrum. 
 
 Cerium, Ce — 140, discovered simultaneously in 1803 by Klaproth and Berzelius 
 and named, as was also the mineral, from the newly discovered planet Ceres, occurs in 
 cerite (60 per cent.), and is also obtained by the electrolysis of the chloride. It is very 
 similar to lanthanum, but at ordinary temperatures is more stable in the air than the latter ; 
 it burns much more readily, so that .broken -off particles of it inflame of their own accord. 
 The specific gravity of the fused metal is 6.72. Besides the salts of the sesquioxide, 
 it forms some from the dioxide, The former in which cerium is trivalent are colorless, 
 
 while the latter — the ceric — are colored yellow or brown ; red ceric hydroxide, Ce(OH)4, 
 is precipitated from the former, on the addition of hypochlorites. A little aqueous hydro- 
 fluoric acid will convert the ceric hydroxide into cej-ium tet)'aJluo7'ide,Ce¥\^ T H2O. The 
 tetrachloride, CeCl4, is also known in double salts. These compounds indicate that cerium 
 is quadrivalent and that it probably belongs to the fourth group of the periodic system 
 (p. 246). 
 
 Didymium, Di — 142, was found in cerite by Mosander, and was regarded as an 
 element until 1885. Auer von Welsbach then succeeded in resolving it into two new 
 elements, which he called Neodymium = Nd (144?) and Praseodymium — Pr 
 (140?). The oxide of the old didymium (from 6i6v/w/, twins, a constant associate of 
 cerium and lanthatuim) has a bright gray color after strong ignition, that of neodymium 
 a bluish color and that of praseodymium a greenish-white color. While the salts of old 
 didymium had an amethyst-red color, those of neodymium are a rose- violet and those of 
 praseodymium are an intense leek-green. 'I'he salts and .solutions of the old didymium 
 .showed a magnificent absorption spectrum and its components share this with it. 'I'he 
 same is true of the spark spectra. 'I'he most important derivatives of praseodymium 
 

 INORGANIC CllKMISTRY. 
 
 were recenlly prepared and studied by von Scheie. I( lias two oxides: I’rO,^ aiul I’r,/ 
 
 Its salts follow the type I’rX, ; the praseodymium in them is pronouncedly Irivalent (Z. f. 
 anorfr. Ch. (1898) 17, 310 ; 18, 352). 
 
 Samarium, Sm 150, was discovered by M. 1 )elafontaine and almost simultaneously 
 by Lecoq de Hoisbaudran in 1878 on studying unknown absorption bands given by the 
 samarskile Irom North Carolina. It has also been found in cerite, gadolinite and orthite. 
 Marignac (1880) prepared the oxide and some salts in a pure state. 1 ’. 'I'. Cleve has made 
 an exhaustive investigation of it. d'he metal has not yet been prepared. 'I'he oxide is 
 pure white and its specific gravity is 8.38. In the lUmsen flame it shows a remarkable 
 light-emission jiower and a discontinuous .spectrum with bright bands. Samarium .salts 
 are light yellow in color. 'I'hey and their solutions exhibit a characteristic absorption 
 spectrum. 'The sulpItaiCy Sm.^(SOj).j | 811 .^ 0 , consists of small crystals, dis.solving with 
 difficulty in water. Its double .salt with potassium sulphate is sparingly soluble in water 
 and almost insoluble in a saturated potassium sulphate .solution, d'he ni/rafr, .Sm(N()3).„ 
 crystallizes, melts easily and is decomposed, d'he chloride, SmClj 611 ./), forms large 
 tabular crystals ; its .S])ark spectrum is very distinct. 
 
 Ytterbium, Vb — 173. Marignac (1878) found that its oxide was the chief constituent 
 of erbium oxide which Hahr aiul Bunsen had regarded as a distinct element. It is ob- 
 tained by the repeated incomplete decomposition of a mixture of the nitrates of these 
 earths with careful heating. Ytterbium and scandium oxides being the feeblest ba.ses, 
 are first eliminated and can then be se])arated by means of potassium sulphate. Ytterbium 
 oxide, Yb./).,, is a white, infusible ])owder of specific gravity 9.17. It is readily soluble 
 in acids; its salts are color]e.ss and do not show an absorption spectrum. Ytlerlduni 
 sulphate, Yb.^(.S()^)3 811 ^ 0 , forms large pri.sms, stable in the air. Its double salt with 
 
 potassium sulphate is readily .soluble in water and in a .saturated potassium sulphate 
 solution. Ytterbium chloride gives a spark spectrum rich in line.s. 
 
 Erbium — Er. Mosander (1843) demonstrated that the supposedly simple yttria was 
 in fact a mixture of the three earths — erbia, terbia and ytterbia. The erbia of to-day 
 is the rose-red oxide from the old yttria, which Mosander called terbia, and which Bahr 
 and Bunsen ( 1866), as well as Cleve and Hoglund (1873) prepared in large quantities 
 from gadolinite and regarded it as elementary. However, Marignac in 1878 obtained 
 ytterbium from it and Nilson in 1879 scandium. The investigations continued in that 
 line proved that it could be still further simplified. The old erbia gave an ab.sorption 
 spectrum very rich in lines and the fact that on partly decomposing its nitrate by fusion 
 or by the fractional precipitation of its nitrate .solution with dilute ammonia, fractions were 
 obtained in which the ab.sorption bands appeared in different number and strength, 
 induced P. T. Cleve to assume in it the presence of three elements — the real Erbium = 
 Er (166), Thulium = Tm and Holmium ^IIo. Then Lecoq de Boisbaudran by similar 
 observations discovered that holmium also contained Dysprosium = Dy {SvaTTpoatro^, 
 difficult to reach). Nilson and Kriiss regard erbium, thulium, holmium, and dysprosium 
 as capable of further division. 
 
 Terbium = Tr. The oxide (terbia) was discovered by Mosander in the old yttria, 
 but called erbia by him at that time. It occurs in small quantities in the gadolinite of 
 Ytterby in Sweden. The names, yttrium, erbium, terbium, and ytterbium, are transposi- 
 tions of the syllables of the word Ytterby. The samarskite from North Carolina contains 
 great quantities of terbia. The metal has not been prepared. Terbia is a dark orange-red 
 powder. It dis.solves in hydrochloric acid with evolution of chlorine. It loses its color 
 when heated in hydrogen, also suffers a slight loss in weight and becomes white. Very 
 probably the orange-colored oxide is a peroxide. The terbium salts are colorless ; they 
 have no absorption spectrum. Its formate dissolves with difficulty in water. The double 
 sulphates with potassium and sodium sulphate are .s[)aringiy soluble in saturated potassium 
 sulidiate and sodium sulphate solutions. The chloride does not yield a spark spectrum, 
 d'erbia is ])robably identical with the oxide of philippium, di.scovered by M. Delafontaine, 
 and the oxide of mosandrium, by |. Lawrence Smith. Lecoq de Boisbaudran considers 
 that terbia contains two additional earths, which he designates temporarily 7m and 7,3. 
 
 Gadolinium, Cd ^ 156? Marignac discovered its oxide in the North Carolina samar.s- 
 kite. It is also |)resenl in so.me ortliites. It is an almost white ]-)owder, very .soluble in 
 acids. 'The salts and salt solutions do not exhibit an absorption spectrum ; the chloride 
 (l(jes not give a spark spectrum. Expo.sed to the kathode rays of an induction machine 
 
THE GALLIUM GROUP. 357 
 
 in an almost vacuous tube gadolinia gives out a fiery red color, and in the spectroscope 
 has a spectrum with a bright orange-red line, whose wave length A = 6094. 
 
 Decipium, Dp =; lyi ? Its oxide decipia (from decipere, to deceive) was discovered 
 in the samarskite of North Carolina by Delafontaine. It is colorless ; its salts also are 
 colorless and show no absorption spectrum. Its double sulphate with potassium sulphate 
 is insoluble in a saturated potassium sulphate solution. Decipium does not give a spark 
 spectrum. 
 
 2. THE GALLIUM GROUP. 
 
 The three heavy metals, gallium, indium, and thallium, bear the same 
 relation to aluminium which copper, silver, and gold bear to sodium 
 and zinc cadmium, and mercury to magnesium. 
 
 Cu 
 
 63.6 
 
 Zn 
 
 65.4 
 
 Ga 
 
 70 
 
 Ge 
 
 72 
 
 As 
 
 75 
 
 Ag 
 
 107-93 
 
 Cd 
 
 112 
 
 In 
 
 114 
 
 Sn 
 
 118.5 
 
 Sb 
 
 120 
 
 Au 
 
 197.2 
 
 Hg 
 
 200.3 
 
 Tl 
 
 204. 1 
 
 Pb 
 
 206.9 
 
 Bi 
 
 208.5 
 
 They constitute the corresponding members of the three great periods, 
 and as a second sub-group attach themselves to aluminium, while cerite 
 metals form the first, more basic sub-group (p. 354). The entire character 
 of the three elements under consideration is influenced by this position 
 in the periodic system, because regular relations appear in all directions, 
 as may be observed, for example, in the specific gravities, melting points, 
 and other physical properties in the free metals : 
 
 
 Ga 
 
 In 
 
 Tl 
 
 Atomic weight, 
 
 70 
 
 II4 
 
 204.1 
 
 Specific gravity, 
 
 5-9 
 
 7-4 
 
 II. 8 
 
 Melting point, 
 
 30° 
 
 176° 
 
 290° 
 
 Being members of group III of the periodic system, Ga, In, and T 1 
 yield compounds of the trivalent form, and these are analogous to those 
 of aluminium in many respects. 
 
 Thallium, like other elements with high atomic weights (Au, Hg, Pb), 
 exhibits great variations from the group properties (p. 321). It yields, 
 for example, not only derivatives of the form TIX3, but also those of 
 'FIX. If we include thallium as a member of the last great period (Pt, 
 Au, Hg, Tl, Pb, Bi), we will discover that, as in the case of the other 
 metals of this series, a remarkable regularity underlies all its forms of 
 combination — the highest as well as the lowest : 
 
 PtCb AuCl IlgCl TlCl PbCb BiCla 
 
 PtCl, AUCI3 IlgCl, TICI3 PbCl, BiXg. 
 
358 
 
 INORGANIC CHEMISTRY. 
 
 1 . GALLIUM. 
 
 Ga = 70. 
 
 Gallium was discovered iii zinc-blende from I’ierrefitte, in 1875, by Lecoq de lioisbaud- 
 ran, by means of the spectrosc()i)e. Since then it has been observed in other blendes, 
 and in the clay iron ore of Yorkshire, as well as in the pig-iron of Middlesborough. It 
 is identical with Linnemann’s austriuin. As early as the year 1869, Mendelejeff, taking 
 the table of the periodic system devised by him as basis, predicted tlie existence of a metal 
 (standing between aluminium and indium, with an atomic weight of nearly 69), which he 
 named Kka-aluminimn. Its i)roperties were necessarily deduced from its position in the 
 periodic system. All the jnoperties of gallium known at that time agreed with those of 
 eka-aluminium, and it seemed very j^robable that this element, which had been theoreti- 
 cally established, was in reality gallium. 
 
 As yet gallium has only been found in very small (juantity, and is but imperfectly 
 investigated. It is characterized by a spectrum consisting of two violet lines. Separated 
 by electrolysis from an ammotiiacal solution of its sulphate, it is a white, hard metal, of 
 specific gravity 5.9, with a fusing point of 4-30°. It is only .superficially oxidized in the 
 air, not altered by water, and is not volatile up to a red heat. Like aluminium, it is 
 scarcely attacked by nitric acid, but dis.solves readily in hydrochloric acid, as well as in 
 caustic pota.sh and ammonia with the evolution of hydrogen. 
 
 Gallium and aluminium unite to form alloys ; those rich in the former metal are liquid 
 at the ordinary temperature, and decompo.se water as energetically as metallic .sodium. 
 
 Gallium Oxide, GajOg, is obtained by igniting the nitrate. It is a white mass which 
 sublimes when heated in a current of hydrogen. 'I'he hydroxide, Ga(OII)3, is thrown 
 out of the solutions of its salts by the alkalies as a white flocculent precipitate, readily 
 soluble in an excess of the precipitant, but rather difficultly soluble in ammonium 
 hydroxide. 
 
 Gallium Chloride, GaCIg, is produced on heating gallium in a current of chlorine 
 gas or in hydrogen chloride ; it forms large colorless crystals which fu.se at 75°, sublime 
 at about 60° and boil at 215-220°. Its vapor density above 440° corresponds to the 
 formula GaCl^, at 270° very closely to Ga.^Clg. The chloride fumes and deliquesces in the 
 air, like aluminium chloride, and decomposes in the evaporation of its aqueous solution. 
 
 Gallium Nitrate, Ga(N03)3, and Gallium Sulphate, Ga.3(S04)3, are crystalline and 
 very deliquescent. The latter forms a double salt with ammonium sulphate — similar to 
 the alums : 
 
 Ga3(SOj3.(NH,)3SO, + 24H.P. 
 
 Hydrogen sulphide only precipitates gallium from acetic acid solutions. 
 
 Gallium Dichloride, GaCl2, is also known. It results on heating the trichloride 
 with metallic gallium. It consists of colorless crystals, melts at 164° and distils at 535 °- 
 Its vapor density at 1000° corresponds to the formula GaClj. 
 
 2 . INDIUM. 
 
 In = 114. 
 
 Owing to its resemblance to zinc, indium was regarded as a bivalent metal, and its 
 compounds were suj)po.sed to have the formula InX.^ ; this fixed its atomic weight at 75-6. 
 'file specific heat, however, made the atomic weight one and a half times as large (p. 
 251). Hence it is trivalent and its derivatives are constituted according to the form 10X3. 
 It belongs to the groiq) of aluminium, and, in its derivatives, manifests some similarity to 
 this metal. 
 
 It was discovered, in 1863, by Reich and Richter, by the aid of spectrum analysis. 
 Its S[)ectrum is characterized l)y a very bright indigo-blue line, hence its name. It only 
 occurs in very minute ([uantilies in .some zinc-blendes from Freiberg and the Hartz. 
 
 It is a silver-white, soft and tenacious metal, of s])(f(:ific gravity '] \ 2 . It melts at 
 176° and distils at a wliin? heat. At ordinary temperatures it is not altered in the air; 
 
THALLIUM. 359 
 
 heated, it burns with a blue flame to indium oxide. It is difficultly soluble in hydro- 
 chloric and sulphuric acids, but dissolves readily in nitric acid. 
 
 Indium Chloride, InClg, results from the action of chlorine on metallic indium, or 
 upon a heated mixture of indium oxide and carbon. It sublimes in white, shining 
 leaflets, which deliquesce in the air. Its vapor density at 600-800° corresponds to the 
 formula InClg, but above 840° it undergoes a gradual decomposition. It does not 
 decompose when its aqueous solution is evaporated. 
 
 Indium Oxide, In^Og, is a yellow powder resulting from the ignition of the 
 hydroxide. 
 
 Indium Hydroxide, In(OH)g, is precipitated as a gelatinous mass, by alkalies, from 
 indium solutions. It is soluble in sodium and potassium hydroxides. 
 
 Indium Nitrate, In(NOg)g, crystallizes with three molecules of water, in white 
 deliquescent needles. *, 
 
 Indium Sulphate, In2(S04)g, remains on evaporating a solution of indium in .sul- 
 phuric acid as a gelatinous mass, with three molecules of water. It forms an alum w'ith 
 ammonium sulphate. 
 
 Indium Sulphide, lugSg, is precipitated by hydrogen sulphide as a yellow-colored 
 compound from indium solutions. It combines to sulpho-salts, InS . SK (from 
 
 InO.OH) with the sulphides of the alkalies. 
 
 Indium also forms a dichloride and monochloride. 
 
 Indium Dichloride, InClg, is produced when metallic indium is heated in a current 
 of hydrogen chloride. It is a white crystalline mass, which on exposure to a more 
 intense heat becomes a yellow liquid and sublimes. Its vapor density at 1000-1400° 
 corresponds to the formula InClg. Water decomposes it at once into indium trichloride 
 and metallic indium : sInClg = 2lnCl3 + In. 
 
 Indium Monochloride, InCl, like gallium dichloride (p. 358), results when the 
 dichloride is heated with metallic indium. It is a crystalline, reddish-yellow mass, which 
 has a reddish-black color when fused. Its vapor density at 1100-1400° corresponds to 
 the formula InCl, Water decomposes it into the trichloride and metallic indium : 
 3lnCl = InClg -)- 2ln. 
 
 3. THALLIUM. 
 
 T1 = 204 . 1 . 
 
 Thallium is rather widely distributed in nature, but in very small quantity. The very 
 rare mineral crookesite contains 17 per cent, of the metal, together with copper, selenium 
 and silver. It is often found with potassium in sylvite and carnallite, in mineral springs, 
 and in some varieties of pyrite and zinc-blendes. When these pyrites are roasted for the 
 production of sulphuric acid, according to the chamber process, the thallium deposits as 
 soot in the chimney and in the chamber sludge, and was discovered in the latter, almost 
 simultaneously, by Crookes (1861) and Lamy (1862 ), by means of the spectroscope. 
 
 To get the thallium, the chimney-dust is boiled with water or sulphuric acid, and 
 thallous chloride precipitated from the solution by hydrochloric acid. The chloride is 
 then converted into sulphate, and the metal separated from the latter by means of zinc or 
 the electric current. 
 
 Thallium is a white metal, as .soft as sodium, and has the specific gravity 11.8. It 
 melts at 290°, and distils at a white heat. It oxidizes very rapidly in moist air. It 
 does not decompose water at ordinary temperatures. It is, therefore, best preserved 
 under water in a clo.sed vessel. By air access it gradually di.ssolves in the water, forming 
 thallous hydroxide and carbonate. Heated in the air it burns with a beautiful green 
 flame who.se spectrum shows a veiy inten.se green line, hence the name thallium, from 
 i^a^Aor, green. Thallium dissolves readily in sulphuric and nitric acids, but is only 
 .slightly attacked by hydrochloric acid, owing to the insolubility of thallous chloride, 
 
 I III 
 
 Thallium forms two series of com[)ound.s : thallous, TlX, and thallic, TlXg. The 
 first are very similar to the compounds of the alkalies (and also to those of silver). The 
 solubility of the hydroxide and carbonate in water shows this ; their solutions have an 
 alkaline reaction. Again, many thallous salts are i.somorphous with those of potassium. 
 
360 
 
 INORGANIC CIIKMIS’IRV. 
 
 iiiul form similar (loul)le sails (see Alums, p. 352). In the insoluhilily of ils sulplmr 
 and haloj^en compounds, univalent thallium api)roaches silver and lead. 
 
 In its comi)oundsof the form 'I'lXj thallium is trivalent, like aluminium, hut otherwise 
 shows .scarcely any similarity to the latter. 
 
 'rhallium compounds are poi.sonou.s. 'i'hey are used in making thallium glass, which 
 refracts more strongly than lead glass. 
 
 THALLOUS COMPOUNDS. 
 
 Thallous Oxide, Tl.^ 0 , is formed by the oxidation of thallium in the air, or by heat- 
 ing the hydroxide to ioo°^ It is a black powder which di.ssolves in water with formation 
 of the hydroxide. 
 
 Thallous Hydroxide, 'ri(OII), may be prepared by decomposing thallous sulphate 
 with an equivalent amount of barium hydroxide, and crystallizes with one molecule of 
 water in yellowish prisms. It di.ssolves readily in water and alcohol, yielding strong 
 alkaline solutions. 
 
 Thallous Chloride, TlCl, forms on heating, the metal in hydrochloric acid gas and 
 is thrown down from solutions of thallous salts by hydrochloric acid as a white, curdy 
 precipitate, which is difficultly soluble in water. It .separates in .small crystals from the hot 
 solution. It fu.ses at 427°, and boils at about 715°. Like ])ota.ssium chloride, it forms an 
 insoluble double salt with jdatinic chloride, rt(jl^.2'riCl. 'J'/iallous hrofnide forms a white, 
 and thallous iodide a yellow ]u-eci{)itate. 'I'he latter is not .soluble in potassium iodide. 
 
 Thallous Sulphate, 'Fl.^SO^, crystallizes in rhombic pri.sm.s, isomorphous with potas- 
 sium sulj)hate. It dissolves in 20 parts of water at ordinary temperatures. It affords 
 double salts with the sulphates of the metals of the magnesium group, of ferrous oxide, 
 of cupric oxide, etc. (p. 315), e. g., MgSO^. Tl^SO^ + 6 H ,^0 ; these are perfectly similar 
 and analogous to the corresponding double salts of potassium and ammonium. It 
 forms thallium alum with the .sulphates of the .sesquioxides of the iron group, e. g., 
 T1A1(S04)2 -f- 12H2O ; these are similar to potassium alum, KA1( 80^)2 + I 2 li 20 . 
 
 Thallous Carbonate, Tl^COg, is obtained from the oxide by the absorption of carbon 
 dioxide ; it crystallizes in needles, which dissolve at ordinary temperatures in 20 parts of 
 water. The solution has an alkaline reaction. 
 
 Thallous Sulphide, Tl^S, is precipitated from thallous salts by hydrogen sulphide 
 as a black compound, insoluble in water. 
 
 THALLIC COMPOUNDS. 
 
 Thallic Chloride, TICI 3 , is produced by the action of chlorine upon thallium or 
 thallous chloride in water, and is very .soluble in the latter. It decomposes at 100° into 
 thallous chloride and chlorine. The alkalies precipitate from its solutions thallic 
 hydroxide, 'riO.OII, a brown powder, which, at 100°, passes into thallic oxide, 
 Tl./ >3. Further heating decomposes the latter into thallous oxide and oxygen. 
 
 'The oxide and hydroxide are soluble in hydrochloric, nitric, and sulphuric acids, 
 forming 'I'lCl.^, Tl(NO.j).,, 
 
 On conducting chlorine through a solution of thallic hydroxide in jiotassium hydroxide, 
 it assumes an inten.se violet color, due ju’obably to the formation of the potassium salt of 
 thallic acid, the compo.sition of which is yet unknown. 
 
METALS OF THE FOURTH GROUT. 
 
 361 
 
 METALS OF THE FOURTH GROUP. 
 
 The elements of group IV in the periodic system (p. 246), 
 
 Ti = 48. 1 Zr = 90.6 Ce = 140 Th = 282 
 
 C = 12.00 Si = 28.4 
 
 Ge = 72 Sn =118.5 Pb = 206.9, 
 
 show the same analogies that were observed with the members of group 
 III (p. 345). Their character is, however, non-metallic ; their deri\ a- 
 tives are chiefly of the types MeX^ and Me02, of which the latter a;e 
 acid (p. 258). 
 
 The first two elements, carbon and silicon, with low atomic weights, 
 belong to the two small periods and are true metalloids. Their oxides and 
 hydroxides are acid in nature. The first more basic sub-group comprises 
 titanium, zirconium, cerium, and thorium. They constitute the fourth 
 members of the large periods. Their compounds are almost exclusively 
 of the type Me02, similar to the silicon derivatives, and are 
 
 usually discussed (with the exception of cerium) with the metalloids after 
 silicon (pp. 238, 355). The other sub-group consists of more electro- 
 negative heavy metals : germanium, tin, and lead. These constitute the 
 transition from the elements in group III, corresponding to them, to those 
 of group V : 
 
 Ga 70 Ge 72 As 75 
 
 In 1 14 Sn 118.5 Sb 120 
 
 Tl 204.1 Pb 206.9 Pi 208.5. 
 
 Their intermediate position accounts for their metalloidal character. In 
 this group, as in all other groups, it is noticed that as the atomic weight 
 rises (from germanium to lead) there is a successive rise in metallo-basic 
 character. All three members form dioxides, 
 
 GeOg SnOj Pb02, 
 
 which may be viewed as anhydrides of the acids 
 
 H 2 Ge 03 H 2 Sn 03 H 2 Pb 03 . 
 
 These are perfectly analogous to silicic acid, but their stability and acidity 
 diminish as the atomic weights of their basal elements increase. 
 
 Lead dioxide, Pb02, combines with bases (especially the alkalies), form- 
 ing salts of plumbic acid, , potassium plumbate, K2Pb03. These are 
 not very stable ; water decomposes them into their components. Lead 
 dioxide unites with difficulty with acids to yield salts. When digested with 
 sulphuric acid it liberates an atom of oxygen, and forms salts of lead mon- 
 oxide, PbO. It yields chlorine with hydrochloric acid, but in the cold 
 the unstable tetrachloride, PbCl^, can be obtained. In this respect lead 
 dioxide resembles the peroxides, mangane.se peroxide, MnOj, and is 
 commonly known as lead peroxide. Plowever, the salts, Me2Pb()3, PbCl^, 
 and the organo-metallic compounds, such as Pb(CH3)^, argue in favor of 
 quadrivalent lead, and make it perfectly analogous to tin (p. 258). 
 
 31 
 
362 
 
 INORGANIC CHEMISTRY. 
 
 The elements of this grou}) yield monoxide derivatives, 
 
 GeO SnO PbO. 
 
 These are commonly known as ous compounds. They are basic and 
 only form salts with acids, 'bhe basicity and stability of their deriva- 
 tives increase as the atomic weights rise. The german^?//^ and stann^7;/j 
 compounds are readily oxidized to derivatives of the dioxide type, while 
 lead monoxide, PbO, is a strong base, and forms very stable salts. 
 
 1. GERMANIUM. 
 
 Ge = 72. 
 
 This element was discovered in 1886 by Cl. Winkler, of Freiberg. As 
 early as 1871 MendelejeflT, with the periodic system as his basis, predicted 
 the existence of an element with an atomic weight of about 73, which 
 corresponded to the then existing gap between silicon and tin; he 
 called it ekasilicon (the first analogue of silicon). The perfect agree- 
 ment of the essential properties of germanium with those of the theoretical 
 ekasilicon constitutes a brilliant confirmation of the law of periodicity 
 (P- 357 )- 
 
 Winkler discovered germanium in the very rare mineral, argyrodite. The latter is a 
 double sulphide of germanium and silver, Ge^S . 3Ag2S. Penfield claims the formula 
 GeS^ . 4Ag,,S for argyrodite as well as for confieldite. The first is monoclinic and the 
 second isometric. It is also present in minute quantities in euxenite (together with 
 titanium and zirconium) (Krtiss), in samarskite and frankeite. It maybe separated from 
 these minerals by fusing them with sulphur and soda. Sodium sulphogermanate is then 
 produced, and it is soluble in water (p. 363). 
 
 To obtain free germanium, its dioxide is heated in a current of hydrogen or reduced with 
 carbon. The product is a dark-gray powder, which melts at 900°, and upon solidifying 
 readily crystallizes into beautiful, grayish-white, metallic octahedra. Its specific gravity 
 at 20° equals 5.469. Its specific heat was found equal to 0.0737 at 100°, and at 440°, 
 0 0757. Therefore, its atomic heat at 100° is 5.33 and at 440°, 5.45. It increases very 
 slightly (like those of aluminium and silicon) with rise of temperature and is a little less 
 than the mean atomic heat (p. 254). 
 
 Germanium is very stable in the air. When ignited it burns with the production of 
 white vapors of germanium dioxide, GeO.^. The metal (like silicon) is insoluble in hydro- 
 chloric acid. Nitric acid converts it (like tin) into the hydrate of the dioxide. It is solu- 
 ble in alkalies upon fu.sion. When heated in the non-luminous gas flame, germanium 
 and its compounds do not impart a color to the same. Its spectrum can only be pro- 
 duced by the action of the induction .spark. 
 
 Germanium, like tin, forms derivatives of the oxides GeO and Ge02 ; the first are called 
 germanowj compounds, the latter germanzV, or derivatives of germanic acid. 
 
 • GERMANOUS COMPOUNDS. 
 
 'I'liese are not very stable, and are readily oxidized to the higher form. 
 
 Germanous Oxide, GeO, is formed when the hydroxide is ignited in a current of 
 carbon dioxide. It is a grayish-black powder. Germanous Hydroxide, Ge(0H)2, is 
 |)rccij)itated as a yellow-colored compound iq)on the addition of caustic alkali to the solu- 
 tion of Ibc (lichloride. It is soluble in hydrochloric acid. 
 
TIN. 363 
 
 Germanous Chloride, GeClg, has not been obtained pure. It is formed when hydro- 
 chloric acid gas acts upon heated germanous sulphide. 
 
 Germanous Sulphide, GeS, is a reddish-brown precipitate produced by the action of 
 hydrogen sulphide upon the solution of the dichloride. It may be obtained in grayish- 
 black crystals by heating germanium sulphide in hydrogen gas. It is soluble in hot 
 hydrochloric acid, forming the corresponding chloride. 
 
 GERMANIC COMPOUNDS. 
 
 Germanium Tetrachloride, GeCl^, is formed by the direct union of germanium with 
 an excess of chlorine. The metal, when gently heated, burns in an atmosphere of chlo- 
 rine, with a bluish color. When in powder form it inflames at the ordinary temperature. 
 The tetrachloride is also produced if the sulphide, GeSg, be heated together with mercuric 
 chloride. It is a colorless, mobile liquid, of specific gravity 1.887 18°. It boils at 
 
 86°. It fumes strongly in moist air, and is decomposed by water into hydrochloric acid 
 and germanic hydroxide, Ge(OII)4. It is not decomposed by concentrated sulphuric 
 acid. Its vapor density, from 300-740°, corresponds to the molecular formula, GeCl^. 
 
 Germanium Chloroform, GeHClg, corresponding to ordinary chloroform, CHCI3 
 (see p. 159), is produced when metallic germanium is heated in a current of hydrochloric 
 acid gas. It is a mobile liquid, boiling at about 72°. Its vapor density approximates 
 the molecular formula GeHClg. It becomes cloudy on exposure to the air, and colorless, 
 oily drops of Germanium Oxychloride, GeOCl2 (?), separate. 
 
 Germanium Bromide, GeBr^, is a strongly fuming liquid, which solidifies at 0° to a 
 crystalline mass. 
 
 Germanium Iodide, Gel^, results upon heating germanium chloride with potassium 
 iodide, or more readily by conducting iodine vapor over heated and finely divided metal. 
 It is an orange-colored solid, melting at 144°, and boiling at 400°. 
 
 Germanium Dioxide, Ge02, germanic anhydride, is formed upon roasting the metal 
 or the disulphide, or by treating the latter with nitric acid. It is a stable, white powder, 
 of specific gravity 4. 70 at 18°. It is slightly soluble in water (i part in 95 parts at 100°) 
 and imparts to the latter an acid reaction. Germanic Hydroxide, Ge(OH)^, or 
 GeO(OH)2, Germanic Acid, is produced by directly transposing the chloride with 
 water. It has not been obtained perfectly pure, as it loses more or less water. Like 
 silicic acid, it is wholly acid in its character, and only forms salts with bases. It is 
 soluble in the hydroxides and carbonates of the alkalies, especially on fusion, while it is 
 almost insoluble in acids. 
 
 Germanic Sulphide, GeS2. Concentrated hydrochloric acid or sulphuric acid will 
 precipitate it from solutions of its sulpho-salts. It is also formed when hydrogen sulphide 
 is conducted through strongly acidulated solutions of the oxide. It is a white, voluminous 
 precipitate, insoluble in acids, but readily soluble in water. If the precipitate is washed 
 with water it dissolves. It is reprecipitated by acids, especially if hydrogen sulphide be 
 conducted through the solution. The sulphide dissolves readily in the fixed alkaline 
 hydroxides and ammonia. It forms sulpho-sa\X.s with the alkaline sulphides. These are 
 perfectly analogous to the sulpho-stannates. Argyrodite is an example of this class, 
 Ag^GeSg + 2Ag2S (p. 362). 
 
 2 . TIN. 
 
 Sn = 118.5. 
 
 Tin {Sfannum) occMYs in nature principally as dioxide (cassiterite, tin- 
 stone) on the Malay Peninsula, in the islands of Banca and Bilitong, 
 and in England (Cornwall), Saxony, India, and in Australia. To pre- 
 pare the metal the oxide is roasted, lixiviated, and heated in a furnace 
 with charcoal : 
 
 SnOj -f- 2C — Sn -p 2CO. 
 
3^4 
 
 INORGANIC CHEMISTRY. 
 
 Thus obtained, it usually contains iron, arsenic, and other metals; to 
 purity it the metal is fused at a low temperature, when the pure tin flows 
 away, leaving the other metals. The tin obtained in the East Indian 
 isles is almost chemically pure, while that of England and of Saxony con- 
 tains traces of arsenic and coj)per. 
 
 Tin is an almost silver-white, strongly lustrous metal, with a specific 
 gravity of 7.3. It possesses a crystalline structure; and when a rod of 
 it is bent it emits a peculiar sound (tin cry), due to the friction of the 
 crystals. Ui)on etching a smooth surface of tin with hydrochloric acid, 
 its crystalline structure is recognized by the appearance of remarkable 
 striations. At low temiieratures i)erfectly pure compact tin ])asses grad- 
 ually into an aggregate of small quadratic crystals. The metal is tolera- 
 bly soft, and very ductile, and may be rolled out into thin leaves (tin-foil). 
 It becomes brittle at 200°, and may then be powdered. It fuses at 231°, 
 and distils at a white heat (about 1700°) ; it burns with an intense white 
 light when heated in the air, and forms tin dioxide. It does not oxidize 
 in the air at ordinary temperatures, and withstands the action of many 
 bodies, hence is employed in tinning copper and iron vessels for house- 
 hold use. 
 
 The most interesting of the tin alloys, besides bronze and soft solder, 
 is britannia metal. It contains 9 parts of tin and i i)art of antimony, 
 and frequently, also, 2-3 per cent, of zinc and i per cent, of copper. 
 
 Tin dissolves in hot hydrochloric acid, to stannous chloride, with 
 evolution of hydrogen gas : 
 
 Sn + 2 HCI =r SnCb -f 211. 
 
 Concentrated sulphuric acid, when heated, dissolves tin, with forma- 
 tion of stannous sulphate. 
 
 Nitric acid, depending upon the temperature and the concentration of 
 the acid, forms soluble stannous nitrate or solid stannic nitrate, which 
 separates and is converted by the increasing dilution of the acid or by 
 hot water into a basic salt and stannic acid. Different stannic acids 
 result in accordance with the conditions of experiment. Anhydrous 
 nitric acid, HNO3, does not change tin. It dissolves when boiled with 
 potassium or sodium hydroxides, forming stannates: 
 
 Sn + 2 KOH + H^O = K^SnOg + 2 !!,. 
 
 There are two series of compounds : the stannous, and stannic or com- 
 pounds of stannic acid. The first readily oxidize to stannic compounds. 
 
 STANNOUS COMPOUNDS. 
 
 Tin Dichloride or Stannous chloride^ SnCl.^, results when tin dissolves 
 in concentrated hydrochloric acid. When its solution is evaporated it 
 crystallizes with two molecules of water (SnCl^ 2H3O), which it loses at 
 100°. It is used in dyeing, as a mordant, under the name of tin salt. 
 The anhydrous chloride, obtained by heating the metal in dry hydro- 
 
STANNOUS COMPOUNDS. 
 
 365 
 
 chloric acid gas, fuses at 250° and distils without decomposition at 606°. 
 Its vapor density at 900° agrees with the formula SnCl2, while at lower 
 temperatures the molecules Sn^Cl^ also seem to exist. 
 
 Stannous chloride dissolves readily in water. Its solution is strongly 
 reducing, and absorbs oxygen from the air with the separation of basic 
 stannous chloride: 
 
 SSnCh + O + H2O = 2 Sn<Q^ + SnCh* 
 
 In the presence of hydrochloric acid, only stannic chloride is pro- 
 duced. Stannous chloride precipitates mercurous chloride or metallic 
 mercury from solutions of mercuric chloride (p. 325). It unites with 
 chlorine to form stannic chloride, and with many chlorides to yield 
 double salts, e. g. : 
 
 SnCl2. 2KCI and SnC^ . 2 NH,C 1 . 
 
 It is decomposed by concentrated hydrochloric acid with evolution of 
 hydrochloric acid gas. 
 
 Tin Monoxide or Sta^inous oxide, SnO, is obtained by heating its 
 hydroxide, Sn(OH)2, in an atmosphere of carbon dioxide ; it is a blackish- 
 brown powder, which burns when heated in the air, and becomes stannic 
 oxide. Sodium carbonate added to a solution of stannous chloride pre- 
 cipitates white 
 
 Stannous staiino-hydrate, Sn(OH)2: 
 
 SnCl2 + Na2C03 + H2O = Sn(OH)2 + 2NaCl + CO2. 
 
 It is insoluble in ammonium hydroxide, but is readily dissolved by 
 potassium hydroxide. Upon slow evaporation of the alkaline solution, 
 dark crystals of stannous oxide, SnO, separate; but, on boiling the solu- 
 tion, the hydrate decomposes into potassium stannate, K2Sn03, which 
 remains dissolved, and metallic tin : 
 
 2 K 2 Sn 02 + H2O = K2Sn03 + 2KOH + Sn. 
 
 The hydroxide forms salts by its solution in acids. Stannous chloride, 
 SnClg, and stannous sulphate, SnSO^, are formed when tin is warmed with 
 concentrated hydrochloric or sulphuric acid. The sulphate separates in 
 small, granular crystals, when its solution is evaporated. 
 
 Tin Monosulphide, Stannous sulphide, SnS, is precipitated from 
 stannous solutions by hydrogen sulphide, as a dark-brown amorphous 
 precipitate. Obtained by fusing tin and sulphur together, it is a lead- 
 gray crystalline mass. It dissolves in concentrated hydrochloric acid, 
 with liberation of hydrogen sulphide, and forms stannous chloride. It 
 is insoluble in alkaline monosulphides, but, if sulphur be added and the 
 solution boiled, it will dissolve as a sulpho-stannate (p. 367): 
 
 SnS -f S + K2S = K2SnS3. 
 
366 
 
 INORGANIC CHEMISTRY. 
 
 STANNIC COMPOUNDS. 
 
 Tin Tetrachloride, Stannic chloride^ SnCl^, is produced by the 
 action of chlorine upon lieated tin or stannous chloride (see Lorenz, 
 Z. f. anorg. Ch. 10 (1895) 44)- is a colorless licpiid {Spiritus fumans 
 Libavii^, fuming strongly in moist air, of specific gravity 2.27, and boil- 
 ing at 114°; its va})or density corres|)onds to the molecular formula, 
 SnCl^. It attracts moisture from the air and is converted into a crystal- 
 line mass (butter of tin), SnCl^ -{- 3^^/^ readily soluble in water. 
 Boiling decomposes the solution into metastannic acid (H^SnO,) and 
 hydrochloric acid : 
 
 SnCh -f 3 1 BO _ 11281103 i 4IICI. 
 
 Stannic chloride possesses a salt-like nature, and combines with metallic 
 chlorides to the so-called double salts, e. g., Sn(d^. 2KCI and SnCI^.- 
 2NH4CI ; the latter compound is known as pink sa/t iind is used in calico 
 printing. It also yields crystalline double salts with chlorides of the 
 metalloids, c. g., SnCl^.PClg and SnCl^.aSCl^. It is not attacked by 
 even hot concentrated sulphuric acid. 
 
 Tin tetrachloride combines with hydrochloric acid, forming HjSnClg -p 6II2O, analo- 
 gous to the chlorplatinic acid, H2PtClg -[- 61 bO. It is formed when hydrochloric acid gas 
 is conducted into a concentrated solution of tin tetrachloride in water. In the cold it 
 solidifies to a leafy crystalline mass, melting at 4 9 °- These double compounds can also 
 be regarded as salts of hydrochlorstannic acid, e. g., (Nlb)2SnClg. 
 
 Tin Bromide, SnBr^, forms a white, crystalline mass that melts at 30° and boils at 
 200°. It unites with hydrogen bromide, forming H2SnBrg 4 8 H., 0 , crystallizing in 
 yellow needles and plates. 
 
 Tin Iodide, Snl^, formed upon heating tin and iodine at 50°, consists of orange-red 
 octahedra, fusing at 146° and boiling at 295°. A solution of tin iodide in arsenic 
 tribromide (p. 146) has the specific gravity 3.73 at 15° ; it is therefore the heaviest solu- 
 tion specifically of which we have knowledge. 
 
 Tin Fluoride, SnFp, is only known in combination with metallic fluorides {e. g., 
 K2SnF]g), which are very similar to and generally isomorphous with the salts of hydro- 
 fluosilicic acid (K28iFl6)- 
 
 Tin Dioxide, Stannic oxide, Sn02, is found in nature as tin-stone or 
 cassiterite, in quadratic crystals or compact, brown masses, of specific 
 gravity 6.8. It is prepared, artificially, by heating tin in the air, and it 
 then forms a white amorjihous powder, or long, rhombic leaflets. It 
 may be obtained crystallized, by conducting vapors of the tetrachloride 
 and water through a tube heated to redness. The dioxide fuses in the 
 electric furnace, and is not soluble in acids or alkalies. When fused with 
 sodium and potassium hydroxide it yields stannates soluble in water. 
 
 The dioxide forms two hydrates — the stannic acids, H.^SnOg and 
 H^Sn(b — corresponding to the hydrates H2CO3, H2Si03 and H^CO^ and 
 114^104(1). 226). Each of these two acids is known in two varieties 
 which differ very much from one another in their deportment. This 
 double appearance of similarly constituted stannic acids was the first 
 instance of isomerism (|). 87). Berzelius (1811) observed it, but he did 
 not comprehend its full imjiort until after (1817) the investigations of 
 Davy and Oay Lussac had been made. Tin and nitric acid of specific 
 
LEAD. 
 
 367 
 
 gravity 1.35 yield stannic nitrate and the /5- or vietastannic acid, of the 
 formula H2Sn03 or H^SnO^, depending on its method of preparation. 
 It is white, voluminous and insoluble in acids. It only dissolves when 
 freshly made in sodium hydroxide. An excess of soda precipitates 
 sodium metastannate from this solution ; it dissolves in pure water. 
 
 The acid upon which hydrochloric acid has been poured dissolves in pure 
 water. Sulphuric acid reprecipitates it. a-Stannic acid — also H2Sn03 or 
 H^SnO^ — is precipitated as a voluminous white powder on the addition of 
 alkali to solutions of the tetrachloride or bromide which have not stood 
 too long. It dissolves readily while moist in mineral acids and in sodium 
 hydroxide, even when the latter is in excess. Both acids separate on 
 heating their dilute hydrochloric acid solutions or upon the addition of 
 ammonia water. The presence of tartaric acid prevents the precipitation 
 of the a-acid by ammonia. The a-acid in hydrochloric acid solution 
 gradually passes into the /5-acid (Lorenz, Z. f. anorg. Ch. 9 (1895), 368). 
 The reason for this first observed case of isomerism has not been explained. 
 
 Most of the salts of stannic oxide with acids, e. g , the sulphate, are not 
 very stable, and washing with warm water decomposes them. The phos- 
 phate and arsenate are exceptions. They are insoluble in water and in 
 nitric acid, which is of importance in analysis. The metallic salts of the 
 stannic acids are more stable. The most important of these is 
 stannate, Na2Sn03 -(- 3H2O, which is employed in calico printing under 
 the name of preparing salts. It is produced upon a large scale by fusing 
 tin-stone with sodium hydroxide. On evaporating the solution, it crystal- 
 lizes in large, transparent, hexagonal crystals. 
 
 Tin Disulphide, Stannic sulphide, 81182, is precipitated as an amor- 
 phous yellow powder by hydrogen sulphide from acid solutions of stannic 
 salts. If a mixture of tin filings, sulphur, and ammonium chloride be 
 heated it is obtained in the form of a brilliant crystalline mass, consist- 
 ing of gold-yellow scales. It is then called mosaic gold, and is applied in 
 bronzing. Concentrated hydrochloric acid dissolves the precipitated di- 
 sulphide, forming stannic chloride ; nitric acid converts it into metastan- 
 nic acid. The sulphides and hydrosulphides of the alkalies dissolve tin 
 disulphide, forming sulphostannates (see p. 223). Sodhun sulphostannate, 
 Na28n83 -(- 2H2O, crystallizes in colorless octahedra. Acids decompose 
 the sulphostannates with the separation of tin disulphide and the evolu- 
 tion of hydrogen sulphide. Freshly precipitated tin disulphide dissolves 
 rather readily in aqueous ammonia forming a red liquid which becomes 
 colorless after standing in the air for some time. From this solution acids 
 precipitate a white tin oxysulphide, 8n82 + 8n80, soluble in aqueous am- 
 monium carbonate. 
 
 3 . LEAD. 
 
 Pb = 206.9 
 
 I.ead {Plumbuni) is found in nature principally as galenite, Pb8. The 
 other more widely distributed lead ores are cerussite, PbCOg, crocoisite 
 (PbCrOJ, wulfenite (PbMoOJ ; also pyromorphite , PbClz -|- 3Pb3(POj2- 
 
368 
 
 INORCJANIC CHKMISrJ<Y. 
 
 anglesilc, PbSO^, and bournonitr, PhS.Sb^S.^. The countries of 
 importance in tlie industry of lead are the Unite(i States, Si)ain, (Ger- 
 many, Clreat Pritain, and New South Wales, (lalenite isthecln’ef source 
 of lead; the i)rocess of its sei)aration is very simple, d'he galenite is 
 first roasted in the air and then strongly ignited away from it. In the 
 roasting, a portion of the lead sulphide is oxidized to oxide and sulphate : 
 
 and 
 
 1‘1)S I 30 = PbO + SOj 
 ri)S I o^ ^pbscv 
 
 Upon ignition, these two substances according to Percy react with the 
 lead sulphide according to the following expiations: 
 
 and 
 
 2 PbO + PbS = 3Pb + SO., 
 PbSO, -I- PbS = 2Pb + 2SO,.- 
 
 Galenite may also be fused with iron : PbS -j- Fe = FeS -|- Pb. Iron 
 sulphide, containing lead sulphide, is used in making sul[)hur dioxide for 
 the lead-chamber process, d'he resulting work- lead is contaminated with 
 nearly all the metals which occur in ores. They are eliminated by rather 
 complex methods. 
 
 Pure metallic lead has a bluish-white color, is very soft, and tolerably 
 ductile. A freshly cut surface has a bright luster, but on exposure to air 
 becomes dull by oxidation. Its specific gravity is 11.37. It melts at 
 about 300°, and distils at a white heat (about 1700°). It burns to lead 
 oxide when heated in the air. 
 
 In contact with air and pure water lead oxidizes to lead hydroxide, 
 Pb(OH)2, which is somewhat soluble in water. If, however, the water 
 contain carbonic acid and mineral salts — even in slight quantity, as in 
 natural waters — no lead goes into solution, but it is covered with an 
 insoluble layer of lead carbonate and sulphate. When much carbon 
 dioxide is present the carbonate is somewhat soluble in water. This 
 behavior is very important for practical purposes, as lead pipes are fre- 
 quently employed in conducting water. Lead also dissolves in ammonia 
 and liine-water in the presence of air. 
 
 Sulphuric and hydrochloric acids have little effect on the compact 
 metal, owing to the insolubility of its sulphate and chloride; yet, if the 
 lead be in the form of powder, both acids will dissolve it. It forms lead 
 nitrate with dilute nitric acid. Zinc, tin, and iron precipitate it, as 
 metal, from its solution ; a strip of zinc immersed in a dilute solution of 
 lead acetate is covered with an arborescent mass, consisting of shining 
 crystalline leaflets (lead tree). 
 
 Alloys. — An alloy of e(]ual i)arts of lead and of tin fuses at 186°, and is 
 used for soldering (soft solder). An alloy of 4-5 parts of lead and i ])art 
 of antimony is very hard, and answers for the manufacture of type (hard 
 lead or type-metal). Shot are lead containing 0.2-0.35 per cent, of 
 arsenic. 
 
 d'he usual lead compounds are constituted according to the type PbX.,, 
 
LEAD. 
 
 369 
 
 and are called plumbic (p. 357). Many of the lead salts are isomorphous 
 with those of barium ; the sulphates of both metals are insoluble in water. 
 In addition to the salts of the type PbX2 lead also forms that of the type 
 PbX^, which are, however, very unstable. 
 
 Lead Oxide, PbO, is produced when lead is heated in air. It melts 
 at a red heat and after fusion it solidifies to a reddish-yellow mass of 
 rhombic scabs (litharge). When lead is carefully roasted, or the hydrox- 
 ide or nitrate ignited, we obtain a yellow amorphous powder called mas- 
 sicot. Lead oxide hasstrong basic properties; it absorbs carbon dioxide 
 from the air, and imparts an alkaline reaction to water as it dissolves as 
 hydroxide. Like other strong bases it saponifies fats (lead plaster). It 
 dissolves in hot potassium hydroxide, and on cooling crystallizes from 
 solution in yellow rhombic or red tetragonal forms, depending upon 
 existing conditions [Geuther, Ann. Chem. 219 (1883), 56]. 
 
 Lead Hydroxide, Pb(OH)2. Alkalies throw it out of lead solutions 
 as a white, voluminous precipitate. 
 
 It dissolves slightly in water, and absorbs carbon dioxide with forma- 
 tion of lead carbonate. When heated to 130° it decomposes into lead 
 oxide and water. It is solul)le in caustic potash or soda. 
 
 If lead or the amorphous oxide be heated to 300-400°, for some time, 
 in the air, it will absorb oxygen and be converted into a bright red pow- 
 der, called red lead, or minium. Its compositi(m corresponds to the 
 formula Pb^ ; it is considered a compound of lead monoxide with lead 
 peroxide : 
 
 Pb, 0 , ^ 2PbO + Pb 02 . 
 
 When minium is treated with dilute nitric acid, lead nitrate passes into 
 solution, while a dark-brown amorphous powder, lead peroxide, Pb02, 
 remains. 
 
 This oxide is more conveniently obtained by the interaction of lead 
 carbonate and sodium hypochlorite, e , if a solution of lead acetate be 
 mixed with an excess of sodium carbonate and chlorine be conducted 
 into the pasty mass : 
 
 PbCOg + Na2C03 + CI2 = Pb02 + 2NaCl 4 2CO2. 
 
 Lead peroxide dissolves in ice-cold concentrated hydrochloric acid to 
 a reddish-yellow liquid containing lead teirachloiide, PbCl^. 
 
 Oxygen is disengaged when sulphuric acid acts upon it, and lead sul- 
 phate (PbSOJ is formed. When dry sulphur dioxide is conducted over 
 it, glowing sets in and lead sulphate results : 
 
 Pb 02 + SO2 = PbSO,. 
 
 When ignited lead dioxide breaks down into lead monoxide and 
 oxygen. 
 
 Lead dioxide, when warmed with potassium hydroxide, dissolves, and on cooling, 
 large crystals of potassinvi pluvibate, K2Pb03 -j 3 HgO, separate out ; these are perfectly 
 analogous to potassium stannate, K2Sn<')3 4- 3H2O. An alkaline lead oxide solution 
 added to a solution of potassium plumbate produces a yellow precipitate (Pl^O^ I^O), 
 
370 
 
 INORGANIC CHEMISTRY. 
 
 which loses water upon gentle wanning, amt is converted into red lead. Therefore, the 
 latter must be considered as the lead salt of a nonnal plumbic acid, Pb(OIl)^, which cor- 
 responds to stannic, Sn(Oll)4, and silicic, Si(Oll)^, acitls : 
 
 Pl).,()^ - rb,rb(),. 
 
 The calcium salt, — calcium orthoplumhate, Ca.^I’bC)^, — is produced upon igniting lead 
 peroxide with lime or calcium carl)onate, or when lead oxide and lime are heated with 
 air access, /. <?., the absorjjtion of oxygen. It is a yellowish-red substance, which breaks 
 down when heated in carb(m dioxide into lead oxide, calcium carbonate and oxygen : 
 
 Ca,Pb(), t 2CO, = PbO 4- 2CaC03 + (),. 
 
 The resulting mixture of lead oxide and chalk regenerates calcium plumbate when it is 
 heated in an air current (Kassner’s oxygen methcKl). 
 
 Lead peroxide and other oxides of lead j)lay an important part in charging accumulators 
 with electricity. In the process of charging the lead oxide at the negative plate is decom- 
 ])osed into lead and oxygen, while at the positive pole it absorbs oxygen and becomes 
 peroxide. In discharging these changes are reversed. 
 
 Another oxide, Pb./).^, which is precipitated as a reddish yellow powder on the addi- 
 tion of sodium hydrocldorite to an alkaline lead .solution, is very probably the lead salt of 
 metaplumbic acid, ll2pb()., : 
 
 Pb 203 PbPbO.,. 
 
 Nitric acid decomposes it into lead nitrate and peroxide. It dissolves in cold hydro- 
 chloric acid without liberation of chlorine ; this gas escapes, however, when the solution 
 is heated. 
 
 Lead Tetrachloride, PbCl^, was prepared by H. Friedrich in 1893 
 [Monats. f. Ch. 14, 505]. It is formed on adding lead peroxide to cold 
 concentrated hydrochloric acid, or upon conducting chlorine into con- 
 centrated hydrochloric acid containing lead chloride in solution or in 
 suspension. On adding ammonium chloride to the clear liquid the 
 double salt, PbCl^. 2NHCI, separates in yellow isometric crystals. It is 
 decomposed, when added to ice-cold concentrated sulphuric acid, into 
 ammonium sulphate, hydrochloric acid and lead tetrachloride, which 
 collects under the sulphuric acid in the form of a yellow oil. 
 
 Lead tetrachloride becomes crystalline at — 15°. It decomposes gradu- 
 ally at the ordinary temperature into lead chloride and chlorine. The 
 decomposition is explosive at 105°; its specific gravity at 0° is 3.18. It 
 fumes in moist air. It is soluble in water. Concentrated solutions when 
 warmed liberate chlorine and lead chloride separates ; in dilute solutions 
 lead peroxide is precipitated. Sodium and potassium hydroxides act 
 similarly. The group similarity with carbon, silicon, germanium and 
 tin exi)resses itself in this chloride ; the chlorides of these other members 
 are liquids at the ordinary temperature. 
 
 Lead Chloride, PbCl^, separates as a white precipitate, when hydro- 
 chloric acid is added to the solution of a lead salt. It is almost insoluble 
 in cold water; from hot water, of which it requires 30 parts for solution, 
 it crystallizx's in white, shining needles. It melts at about 500° and solidi- 
 fies to a horn-like mass. It is volatile at a white heat; its vapor density 
 corresponds to the formula P1)C1.2. 
 
 Lead Iodide, Pbl2, is thrown down as a yellow precipitate from lead 
 solutions by j)otassium iodide ; it crystallizes from a hot solution in shin- 
 ing, yellow leaflets, melting at 383°. It is difficultly soluble in water. 
 
LEAD. 371 
 
 Lead tetraiodide, Pbl4, and tetrabromide, PbRr,, are known in the form of double 
 salts with salts of organic bases (Classen and Zahorski). 
 
 Lead Nitrate, Pb(N 03 ).^, obtained by the solution of lead in dilute 
 nitric acid, crystallizes in regular octahedra (isomorphous with barium 
 nitrate) and dissolves in two jtarts of water at the ordinary temperature. 
 It melts at a red heat, and is decomposed into lead monoxide, nitrogen 
 dioxide, and oxygen. When boiled with lead oxide and water, it is con- 
 verted into the basic nitrate, which separates in white needles. 
 
 Lead Sulphate, PbSO^, occurs in nature as ang/esiie, in rhombic 
 crystals, isomorphous with barium sulphate. It is precipitated from lead 
 solutions as a white crystalline mass by sulphuric acid. It dissolves with 
 difficulty in water, more readily in concentrated sulphuric acid, very easily 
 in concentrated sodium hydroxide (distinction from barium sulphate). 
 When ignited with carbon, it is decomposed according to the following 
 equation : 
 
 PbSO^ + 2C = PbS + 2CO.,. 
 
 Lead Carbonate, PbCOg, occurs as cerussite in nature. It is pre- 
 cipitated by ammonium carbonate from lead nitrate solutions. Potassium 
 and sodium carbonates precipitate basic carbonates, the composition of 
 which varies with the temperature and concentration of the solution. 
 A similar basic salt, whose composition agrees best with the formula : 
 
 2 PbC 03 . Pb(OH)2* 
 
 is prepared on a large scale by the action of carbon dioxide upon lead 
 acetate. It bears the name white lead. 
 
 White lead was formerly manufactured by what is known as the Dutch process. Leaden 
 plates rolled up into a spiral were moistened in earthenware pots with acetic acid, and then 
 covered with manure and permitted to stand undisturbed for some time. In this way the 
 action of the acetic acid and air upon the lead produced a basic acetate, which the carbon 
 dioxide, evolved from the decaying manure, converted into basic lead carbonate. At 
 present it is prepared by dissolving litharge in acetic acid, and converting the resulting 
 basic acetate into a carbonate by conducting carbon dioxide into it. 
 
 White lead is employed for the manufacture of white oil colors. As it is poisonous, 
 and blackened by the hydrogen sulphide of the air' (formation of lead sulphide), it is 
 being replaced more and more by zinc white and permanent white (BaSO^). 
 
 Lead Sulphide, PbS, occurs crystallized in metallic, shining cubes 
 and octahedra. Hydrogen sulphide precipitates it as an amorphous black 
 y^owder. It is insoluble in dilute acids. Hydrogen sulphide preciydtates 
 red chloro-sulyffiide, PbS . PbClg, from solutions of lead salts containing 
 much strong hydrochloric acid. 
 
 The soluble lead compounds are very poisonous. They have a sweet- 
 ish, astringent taste. They are readily recognized by the following reac- 
 
 *Its chemical structure may be expressed by the following formula : 
 
 HO-Pb-COg-Pb-COg-Pb-OH. 
 
372 
 
 INORGANIC CHEMISTRY. 
 
 tions : sulpliuric acid i)recipitates wliite lead siili)hale, wliich is soluble 
 in sodium hydroxide; from this solution hydrogen sulphide ])reci])itates 
 black lead sulphide. Potassium iodide [)recii)itates yellow lead iodide. 
 
 4. BISMUTH. 
 
 \V\ = 208.5. 
 
 Bismuth constitutes a natural group with antimony, arsenic, phosphorus 
 and nitrogen. We observed that, with increase of atomic weight, the 
 metalloidal character of the lower members becomes more metallic ; the 
 acid nature of the oxides becomes basic. Antimony oxide (SI)./),) is a 
 base, while the higher oxide, Sb.p^^, represents an acid anhydride. In 
 bismuth, the metallic nature attains its full value. This is manifest in 
 its inability to unite to a volatile compound with hydrogen and in the 
 basic oxyhydrate BiO(OH). Bismuth trioxide is a base, and the pent- 
 oxide j)ossesses a very feeble acid character ; the latter behaves more like 
 a metallic peroxide. 
 
 Bismuth usually occurs native, and in combination with sulphur, as 
 bismuthinite ; vvitli tellurium as tetradymite ; also as bisjnuih ocher ^ Bia^g? 
 and in many cobalt, nickel and silver ores. 
 
 Bismuth is obtained principally from Saxony, and in London from Bolivian and Aus- 
 tralian ores. In Saxony the roasted ores are extracted with concentrated hydrochloric 
 acid, and the solution then diluted with water, when the oxychloride is precipitated. It 
 is redi.ssolved in hydrochloric acid and freed from iron by precipitation with water, and 
 then melted, together with coal, lime, and slag, in graphite crucibles to obtain the 
 bismuth. 
 
 To obtain the metal, the sulphide is roasted in the air, and the result- 
 ing oxide reduced with charcoal. 
 
 Bismuth is a reddish-white metal, of specific gravity 9.9. It is brittle 
 and may be easily pulverized. It crystallizes in rhombohedra. It fuses 
 at 265° and distils at a white heat (about 1300°). It does not change in 
 the air at ordinary temperatures. When heated it burns to bismuth 
 oxide, Bi203. It is insoluble in hydrochloric acid, but dissolves in boil- 
 ing sulphuric acid with formation of sulphate of bismuth, and the evolu- 
 tion of sulphur dioxide. Nitric acid dissolves it readily in the cold. 
 
 Water decomposes bismuth solutions in the same manner as those of 
 antimony; insoluble basic salts are precipitated, but these are not dis- 
 solved by tartaric acid (distinction from antimony). 
 
 Bismuth Chloride, BiCl^, arises from the action of chlorine upon 
 heated bismuth, and by the solution of the metal in aqua regia. It is a 
 white crystalline mass, which fuses at 220° and boils at about 447°. It 
 deliquesces in the air. Water renders its hydrochloric acid solution turbid, 
 a white, crystalline ])recipitate of bismuth oxychloride, BiOCl, separating 
 at the same time ; 
 
 BiCIg + lIjO = BiOCl T 2HCI. 
 
 The metalloidal character of bismuth is indicated bv this reaction. 
 
CHROMIUM GROUP. 
 
 373 
 
 The compounds BiBr^ (orange-yellow) and Bilg (black or dark brown) 
 are very similar to bismuth chloride. All three combine with many 
 metallic haloid salts to form double halogen derivatives. 
 
 Halogen derivatives of quinquivalent bismuth are unknown. 
 
 Bismuth Oxide, Bi203, prepared by burning bismuth or heating the 
 nitrate, is a yellow powder, insoluble in water and the alkalies. 
 
 Normal bismuth hydroxide, Bi(OH)3, is not known in a free state. 
 Potassium hydroxide added to a bismuth solution precipitates a white 
 amorphous 7 )ietahydrate, BiO. OH. 
 
 Chlorine conducted through a concentrated potassium hydroxide solu- 
 tion in which bismuth oxide is suspended precipitates red bismuthic 
 acid (HBiOg or H.2Bi20g), which, when gently heated, becomes Bi203, 
 bis 7 iiuthic oxide. Strong ignition converts the latter into bismuth tri- 
 oxide and oxygen; hydrochloric acid dissolves it to bismuth chloride, 
 with liberation of chlorine. 
 
 Bismuth Nitrate, Bi(N03)3, is obtained by the solution of bismuth 
 in nitric acid, and crystallizes with five molecules of water in large, trans- 
 parent tables. In a little water it dissolves without any change; much 
 water renders it turbid, owing to the precipitation of white, curdy basic 
 CNO, CNO, 
 
 salts: Bi < NO^and Bi < OH. The precipitate is emploved in medicine 
 
 (oh (oh 
 
 under the name of Bis 7 }iuthu 7 n sub 7 titricu 77 i {sub 7 iitrafe'). It frequently con - 
 tains tellurium. 
 
 Bismuth Sulphate, Bi2(S04)3, is formed when bismuth dissolves in 
 sulphuric acid. It crystallizes in delicate needles. Bismuth Sulphide, 
 1^283, occurring as bismuthinite, is thrown down as a black i)recipitate 
 from bismuth solutions by hydrogen sulphide. Unlike antimony and 
 arsenic sulphides, it does not form sulpho-salts. 
 
 The alloys of bismuth are nearly all readily fusible. An alloy of 4 
 parts of bismuth, i part of cadmium, i part of tin, and 2 ])arts of lead, fuses 
 at 65° (Wood’s metal). The alloy of 9 parts of bismuth, i part of lead 
 and I part of tin (Rose’s metal) fuses at 94°. 
 
 Alkaline stannous solutions precipitate metallic bismuth from solutions 
 of its salts : 
 
 2BiO . OH + 3X2800, = 2Bi + 3X28003 + H2O. 
 
 CHROMIUM GROUP. 
 
 We observed that a group of more metallic analogous elements : titan- 
 ium, zirconium and thorium attached itself to the metalloidal elements, 
 carbon, silicon and tin (p. 238); and further that their was an analogous 
 group of more metallic elements : vanadium, niobium and tantalum, cor- 
 responding to the metalloidal group of phosphorus (|). 226). We now 
 meet a group of metals, consisting of chromium, molybdenum, tungsten, 
 and probably uranium, that bears a like relation to the elements of the 
 
374 
 
 INORGANIC CHEMISTRY. 
 
 sulphur group (see Periodic System of the Elements). The resemblance 
 of these elements to sulphur and its analogues is plainly manifest in their 
 highest oxygen compounds X()^, and their derivatives (see also Man- 
 ganese). As the elements of the sulphur group in their highest oxygen 
 compounds are sexivalent, so chromium, molybdenum, tungsten and 
 uranium form acid oxides: CrO,, M0O3, UO3. Many of the salts 
 
 corresponding to these are very similar to and isomorphous with the salts 
 of suli)huric acid. Sodium chromate, like sodium sulphate, crystallizes 
 with ten molecules of water ; the potassium salts of both groups form iso- 
 morphous mixtures; their magnesium salts, as well as that of tungstic 
 acid, have the same constitution : 
 
 MgSO, + 7H2O and MgCrO^ + 7II2O. 
 
 Corresponding to the acid oxides are the chlorine derivatives: 
 
 SO2CI2, CrO,Cl2, M0O2CI2, MoOCl,, WOCl, and WClg, 
 
 which are perfectly analogous, so far as chemical deportment is con- 
 cerned. 
 
 The most important basic oxide of chromium is its sesquioxide. This 
 affords salts with the acids, and they are perfectly similar to those of the 
 sesquioxides of iron (Fe203), manganese (Mn203), and aluminium (AI2O3) 
 (P- 33 °)- 
 
 II 
 
 Finally, compounds of chromium, CrX2, are known in which the metal 
 figures as a dyad. These so-called chromoiis compounds are very unstable, 
 and are oxidized by the air into chromic compounds. In this respect 
 they also resemble the ferrous derivatives FeX2 (p. 398). 
 
 Chromium compounds containing more oxygen than the chromates are 
 found in the salts of perchromic acid ; however, nothing definite is known 
 of their constitution at present. The same is true of tungsten and molyb- 
 denum (pp. 381, 383, 384). 
 
 Salts of molybdenum and tungsten, corresponding to the states of 
 lowest oxidation, are not known, because these metals occur as hexads 
 in most of their derivatives. Uranium, which has the highest atomic 
 weight of the group, shows some variations from its analogues. 
 
 1. CHROMIUM. 
 
 Cr = 52.1. 
 
 Chromium is found principally as chromite in nature. This is a com- 
 bination of chromic oxide with ferrous oxide, Cr203.Fe0, and occurs in 
 North America, Sweden, Hungary, and in large quantities in the Urals. 
 Crocoisite, or lead chromate (PbCrOJ, is not met with so frequently, 
 (diromite is used almost exclusively for the preparation of all other 
 chromium derivatives, as it is first converted into potassium chromate 
 by fusion with potassium carbonate and nitrate (p. 379) 
 
 Metallic chromium was first obtained free from cai bon and fused in large 
 
CHROMIC COMPOUNDS. 375 
 
 quantities in 1898 by Goldschmidt, who ignited a mixture of chromic 
 oxide and aluminium powder (p. 347) : 
 
 Itis bright gray in color, very brilliant, and on cleavage faces shows strong 
 crystallization. It is very hard, extremely difficult to melt and very stable 
 in the air. W. Hittorf claims that chromium behaves like a noble metal at 
 the ordinary temperature in that it is inactive electromotively and in the 
 tension series arranges itself with platinum. Hydrochloric acid at 0° 
 has no action upon it and when it is made the anode it dissolves in the 
 acid as chromic acid. At more elevated temperatures it decomposes 
 hydrochloric acid with the liberation of hydrogen ; and in an electro- 
 motive respect follows zinc. It dissolves as chromous chloride. 
 
 Four series of chromium compounds are known : chromous (CrX2), 
 chromic (CrX.^), the derivatives of chromic acid, called chromates, and 
 the perchromates. All chromium compounds are brightly colored, hence 
 the name chromium (from color). 
 
 CHROMOUS COMPOUNDS. 
 
 These are very unstable, and by oxidation pass readily into chromzV compounds. Like 
 ferrous salts, they are produced by the reduction of the higher oxides. The following may 
 be mentioned : Chromous Chloride, CrClg. This is obtained by heating chromic 
 chloride, CrCl.^, in a stream of hydrogen or by dissolving chromium in hydrochloric acid. 
 It is a white crystalline powder. It volatilizes without decomposition. At 1300-1600° 
 the vapor density corresponds to a mixture of the molecules CrClg and CrgCh. It dis- 
 solves in water with a blue color ; the solution absorbs oxygen with avidity, and becomes 
 green in color [Ann. Chem. 228 (1885), 113]. The alkalies precipitate yellow chro- 
 mous hydroxide, Cr(OH)2, from it. This is readily oxidized. When heated it parts 
 with hydrogen and water and becomes chromic oxide : 
 
 2Cr(OH)2 = Cr 203 + H 2 + H 2 O. 
 
 CHROMIC COMPOUNDS. 
 
 Chromic Chloride, CrClg, like AICI3, is obtained by ignition of the 
 oxide and charcoal or, better, the metal in a current of chlorine. When 
 raised to a red heat in this condition it sublimes in shining violet leaflets, 
 which are transformed into chromic oxide by ignition in the air. Its 
 vapor density at 1200-1300° corresponds to the formula CrClg; it vapor- 
 izes very slowly below 1000°. 
 
 Pure chromic chloride only dissolves in water after long-continued 
 boiling; if, however, it contains traces of chromous chloride, CrCl2, it 
 dissolves readily at ordinary temperatures. Green crystals of CrClg -|- 
 6H2O separate from the green solution on evaporation ; these deliquesce 
 in the air. The same crystals may be obtained from solutions of chromic 
 
376 
 
 INORGANIC CHEMISTRY. 
 
 hydroxide, Cr(OH)3, in hydrocliloric acid. When they are dried inter- 
 mediate oxychlorides, CrCl2(OH) and Cr2Cl(OI I)^, are formed ; finally 
 Cr(OH)3 remains. 
 
 Chromic Hydroxide, Cr(OH),,. It is i)recii)itated by ammonium 
 hydroxide from chromic solutions as a voluminous bluish gray, hydrous 
 mass. The green preciju tales i)roduced by sodium and i)Otassium hydrox- 
 ides contain alkali that cannot be removed even by boiling water. 
 They dissolve readily with an emerald-green color, in an excess (T jiotas- 
 sium or sodium hydroxide (slightly in ammonia), but they are reprecij)!- 
 tated upon boiling their solutions. When chromic hydroxide is healed 
 to 200° in a current of hydrogen, the i)roduct is the hydroxide, CrC). Oil, 
 which is a grayish-blue i)owder, insoluble in dilute hydrochloric acid. 
 
 Chromic Oxide, Cr^O^, is a green, amorphous powder. It is also 
 formed by the ignition of chromium trioxide: 
 
 2Cr03 -= C1-2O3 -f- 3O, 
 or of ammonium dichromate: 
 
 (NIb),Cr.A = Cr 203 + 41^0 + N^. 
 
 It may be obtained in black, hexagonal crystals, by conducting the 
 vapors of the oxychloride (p. 380) through a tube heated to redness: 
 
 2Cr02Cl2 = Cr203 + 201^ + O. 
 
 Ignited chromic oxide is insoluble in acids. When fused with silicates, 
 it colors them emerald-green, and serves, therefore, to color glass and 
 porcelain. 
 
 Giiignef s green is a beautifully green-colored chromium hydroxide, 
 which is applied as a paint. It is obtained by igniting a mixture of one 
 part of potassium dichromate with three parts of boric acid ; after treating 
 the mass with water, which dissolves potassium borate, there remains a 
 green powder, the composition of which corresponds to the formula : 
 
 Cr20(0H)^. 
 
 The predominating properties of chromic oxide are basic, as it readily 
 affords salts with acids; yet its basic nature, like that of all sequioxides, 
 is but slight, so that it does not afford salts with weak acids (p. 374). 
 In addition to all this it possesses a slightly acidic character, and metallic 
 salts are derived from it, generally from the hydroxide, CrO.OH, which 
 are analogous to the aluminates (p. 351). Such salts — chromites — like 
 (OO . 0)2Mg and (CrO . can be obtained crystallized in regular 
 
 octahedra by fusing chromic oxide with metallic oxides and boron trioxide 
 (as flux). The mineral chromite is such a salt : 
 
 Cr2O3.Fe0-^ (CrO.O)2Fe. 
 
 Chromium Sulphate, Cr2(SO,)3, is obtained by dissolving the 
 hydroxide in concentrated sulphuric acid. 'I'he solution, green at first. 
 
DERIVATIVES OF CHROMIC ACID. 
 
 377 
 
 becomes violet on standing, and deposits a violet-colored crystalline 
 mass. This may be purified by solution in water and precipitation by 
 alcohol. This salt crystallizes from very dilute alcohol in bluish-violet octa- 
 hedra containing fifteen molecules of water. If the aqueous solution of 
 the violet salt be heated, it assumes a green color, because the salt breaks 
 down into free acid and a basic salt which, upon evaporation, separates 
 as a green amorphous mass, soluble in alcohol. When the green solution 
 stands, it reverts to the violet of the neutral salt. The other chromic 
 salts, the nitrate and the alum, behave in a similar manner. 
 
 It is very probable that in the green solutions the chromium exists as 
 a part of peculiar “complex” acids: chrom-sulphuric acid, chrom-nitric 
 acid ; it no longer appears as ion, but as a part of the latter. Hence the 
 salts of green solutionsdo not show the reactions of salts of chromic oxide 
 nor of sulphuric acid. 
 
 Chromium sulphate forms double salts with the alkaline sulphates — the 
 chromium alums (p. 352). 
 
 Potassium Chromium Alum, K2SO^ . Cr2(SOj3 24H2O, crys- 
 tallizes in large, dark-violet octahedra. It is most conveniently ])repared 
 by conducting sulphur dioxide through a solution of potassium dichromate 
 containing sulphuric acid ; 
 
 K2Cr20, + H2SO, + 3SO2 = Cr2(SOd3 . K2SO, + H2O. 
 
 At 80° the violet solution of the salt becomes green, and on evapora- 
 tion yields an amorphous green mass. 
 
 As chromium hydroxide possesses only a slightly basic nature, salts 
 with weak acids, like carbonic, sulphurous, and hydrosulphuric acids (see 
 Aluminium, p. 325) do not exist. Therefore, the alkaline carbonates and 
 sulphides precipitate chromium hydroxide from solutions of chromium 
 salts : 
 
 Cr2(SOd3 + 3Na3C03 -f 3H2O = 2Cr(OH) + 3Na3SO, + 3CO2 
 
 and 
 
 -1- 3 (NHd 2 S + 6H2O = 2Cr(OH)3 + 3 (NHd 2 SO, + 3H2S. 
 
 Ammonium sulphide produces a black precipitate, chromous sulphide, 
 CrS, in solutions of chromous salts. 
 
 DERIVATIVES OF CHROMIC ACID. 
 
 In its highest oxygen derivative, CrOg, chromium possesses a complete 
 metalloidal, acid-forming character. Chromic acid, HgCrO^, is perfectly 
 analogous to sulphuric acid, H2SO^, but has not been obtained free, since 
 when liberated from its salts it at once breaks down into the oxide and 
 water : 
 
 HgCrO^ = CrOg + H2O. 
 
 The chromates are often isomorphous with the corresponding sulphates 
 (P* 379)- Polychromaies also exist, and are derived from polychromic 
 32 
 
INORGANIC CHEMISTRY. 
 
 
 acids produced by tlie condensation of several molecules of the inninal 
 acid (see Disulphuric Acid, p. 193) : 
 
 Il^CrO^, IbCr./)., I etc. 
 
 Chromic acid. Dichromic acid. Trichromic acid. 
 
 The constitution of their salts is e.xprcssed by the following formulas; 
 
 Cr(),< 
 
 OMe 
 
 OMe 
 
 Cr(),< 
 
 Cr().,< 
 
 OMe 
 
 O 
 
 ( )Me 
 
 Cr(),,< 
 
 CrO.,< 
 
 CrO,< 
 
 OMe 
 
 O 
 
 o 
 
 OMe. 
 
 The free polychromic acids are not known, because as soon as they are 
 sejiarated from their salts, they immediately break down into the acid 
 oxide and water : 
 
 = 3CrO, + II.p. 
 
 The polychromates are fretpiently, but incorrectly, called salts ; 
 true acid or ])rimary salts, in which only one hydrogen atom is replaced 
 by metal (KHCrOJ, are unknown for chromic acid. 
 
 The salts of normal chromic acid are mostly yellow colored, while the 
 polychromates are red. The latter are produced from the former by the 
 action of acids ; 
 
 2K,CrO, + 2IIXO3 = K,Crp^ -f 2KNO3 + 
 
 Conversely, by the action of the alkalies, the polychromates pass into the 
 normal salts : 
 
 K2Ci' 207 + 2KOII = 2K2Cr04 -f IbO. 
 
 The formation of the polychromates maybe explained as follows : The 
 chromic acid liberated from its salts by stronger acids breaks down into 
 water and the acid oxide, which combines with the excess of the normal 
 chromate : 
 
 K 2 Cr 04 -h CrOg = K2Cr20^. 
 
 When there is an excess of acid the anhydride (CrOg) is set free. 
 
 Chromium Trioxide, or Chromic acid a?thydride {Acidnm chro- 
 f?iicu?n), CrOg, consists of long, red, rhombic needles or prisms, ob- 
 tained by adding sulphuric acid to a concentrated potassium dichromate 
 solution. The crystals are readily soluble in water and deliquesce in 
 the air if not perfectly free from sulphuric acid. When heated, they 
 blacken, melt, and at about 250° decompose into chromic oxide and 
 oxygen : 
 
 2Cr03 = Ci-gOg + 3O. 
 
 It volatilizes without decomposition in slight quantities and this even 
 occurs below its jjoint of fusion. 
 
 Chromium trioxide is a poweiful oxidizing agent, and destroys organic 
 matter; hence its solution cannot be filtered through paiier. When alco- 
 hol is iioured on the crystals, detonation takes })lace, the alcohol burns, 
 and green chromic oxide remains. By the action of acids, e. g., sul- 
 phuric, the trioxidc deports itself like a peroxicle; oxygen escapes and a 
 
DERIVATIVES OF CHROMIC ACID. 
 
 379 
 
 chromic salt results. When heated with concentrated hydrochloric acid 
 chlorine is evolved : 
 
 2 Cr 03 + 12HCI = Ci-jClg + 6II2O + 3CI2. 
 
 Reducing substances, like sulphurous acid and hydrogen sulphide, con- 
 vert chromic acid into oxide : 
 
 2Cr03 -|- 3112^ — ^^2^3 ~f' “1“ 3S, 
 
 2Cr03 + 3SO2 + 3H2O == Cr203 + 3H2SO^ = €12(80^3 + 3H2O. 
 
 Alcohol acts similarly, being oxidized to aldehyde : 
 
 ^CrOg + 3H2SO4 -j- 3C2HgO = Crg (804)3 + 6H2O 3C2H4O. 
 
 Potassium Chromate, K2Cr04, is obtained by adding potassium 
 hydroxide to potassium dichromale. It forms yellow rhombic crystals, 
 isomorphous with potassium sulphate (K2SO4) ; isomorphous mixtures crys- 
 tallize out from the solution of the two salts. 
 
 Neutral alkali chromates are always produced when any chromium com- 
 pound is fused with an alkaline carbonate and some oxidizing agent (salt- 
 peter, potassium chlorate). The addition of the oxidant is unnecessary 
 if the oxygen of the air can act sufficiently upon the fusion. Chromic 
 iron ore is roasted in reverberatory furnaces with soda and lime. Neutral 
 sodium chromate, ferric oxide, and calcium carbonate are the products : 
 
 2^203. FeO -j- 4Na2C03 -f- 4CaO -j- 7O = 4Na2Cr04 4CaC03 -f- ^^203- 
 
 The neutral chromate is extracted with water. When this solution is 
 concentrated and mixed with sulphuric acid anhydrous sodium sulphate 
 separates even from the hot liquid, from which, on further evaporation. 
 Sodium Dichromate, Na2Cr20^ -f 2H2O, deposits in red crystals on 
 cooling : 
 
 2Na2Cr04 + H28O4 = Na2Cr207 + Na2S04 + H2O. 
 
 All other chromium preparations are made from this salt [see Hausser- 
 mann, Jahrb. f. Ch. (1891), 1, 327]. Chromates are also produced if 
 chromium or its alloys are made the anode in an alkaline bath. 
 
 Potassium Dichromate, or Bichromate of potash, called 
 
 acid potassium chromate, is manufactured on a large scale by trans- 
 posing the sodium salt with potassium chloride : 
 
 Na2Cr20.j -f 2KCI = K2Cr20^ -f 2NaCl. 
 
 The great difference in solubility in water renders the separation of the 
 sodium chloride and potassium dichromate an easy matter. 
 
 Potassium dichromate (called red chromate of potash in commerce) 
 crystallizes in large, red, triclinic prisms, soluble at ordinary tempera- 
 tures, in ten parts of water. When heated, the salt fuses without change ; 
 at a very high heat it decomposes into potassium chromate, chromic 
 oxide, and oxygen : 
 
 2K2Cr207 = 2K2CrO^ -j- Cr203 3O. 
 
380 
 
 INORGANIC CHEMISTRY. 
 
 When the salt is warmed with sulphuric acid, oxygen escapes and potas- 
 sium chromium alum is produced : 
 
 lY2Cr207 4 " 411280^ — Cr2(S()j)3. 1^280^ -j- 41120 -f- 3O. 
 
 This reaction answers for the preparation of iierfectly pure oxygen. 
 P\irther, the mixture is made use of in laboratories, as an oxidizing agent. 
 
 Sodium Chromate, Na2CrO^ -f- 10H2O, forms deliciuescent crys- 
 tals, and is analogous to ('dauber’s salt (Na2S()^ -f- 10H2O). Barium 
 and Strontium Chromates, BaCrC)^ and SrCrO^, are almost insoluble 
 in water. Calcium Chromate, CaCrO^, dissolves with difficulty in 
 water, and crystallizes like gypsum with two molecules of water, d'he 
 7 ?iagnesiu?n salt, MgCrO^ -|- 7H2^> dissolves readily and corresponds to 
 Ep.som salts, MgSC)^ -j- 7H2(). I'he chromates of the heavy metals are 
 insoluble in water, and are obtained by transposition. 
 
 Lead Chromate, PbCrO^, is obtained by the precipitation of 
 soluble lead salts with potassium chromate. It is a yellow amorphous 
 powder which serves as a yellow paint — chrome yellow. When heated it 
 melts undecomposed, and solidifies to a brown, radiating crystalline 
 mass. It oxidizes all the carbon compounds at a red heat, and is there- 
 fore used in their analysis. In nature lead chromate exists as crocoisite. 
 
 Chromic Acid Chloranhydrides. — Chromic acid forms chloranhy- 
 drides similar to those of sulphuric acid (p. 195). Corresponding to sul- 
 phuryl chloride, SO2CI2, we have chromyl chloride, Cr02Cl2 ; and for the 
 
 r Cl (Cl 
 
 first sulphuric acid chloranhydride, SO2 j is the salt, CrOg j 
 
 VI r\ VI r\ 
 
 Cr02<oK 
 
 VI 
 
 CrO, 
 
 ,OK 
 
 OK- 
 
 Chromyl Chloride, CrOgCla, Chromhwi oxychloride, is produced by 
 heating a mixture of potassium (or sodium) dichromate (or monochro- 
 mate) and sodium chloride with sulphuric acid : 
 
 Na2Cr04 -p 2NaCl -P 4H2SO4 = Cr02Cl2 -p 4NaH804 4 - 2H2O, 
 
 and 
 
 Na2Cr207 -p 4NaCl + 611280^ = 2Cr02Cl2 + 6NaH804 + 3H2O. 
 
 The water produced at the same time must be absorbed by the excess 
 of sulphuric acid. 
 
 To prepare chromyl chloride, first fuse salt (10 parts) with potassium dichromate (12 
 parts) or with potassium monochromate (17 parts). The yellowish-brown mass is broken 
 into coarse ])ieces, jdaced in a retort provided with a condenser, and anhydrous or slightly 
 fuming sul])huric acid (30 ])arts) poured over them. When a gentle heat is applied 
 chromyl chloride distils over aiul is purified by further distillation. 
 
 Chromyl chloride is a red, transparent liquid, of specific gravity 1.91 
 at 25°, and fumes strongly in the air. It boils at 1 16-118° ; its vapor 
 density corres])onds to the molecular formula Cr0.2Cl2. 
 
DERIVATIVES OF CHROMIC ACID. 381 
 
 With water it is decomposed according to the following equation; 
 
 CrO.Cl., + H 2 O = CrOj + 2 HCI. 
 
 Cl 
 
 Chloro-chromic Acid, CiO.Xqp^ (see p. 380), is only known in 
 
 its salts. The potassium salt is formed by heating potassium dichromate 
 (3 parts) for a short period, with concentrated hydrochloric acid (4 parts) : 
 
 K 2 Cr 207 -f 2 HCI = 2Cr02<^j^ + H 2 O. 
 
 It crystallizes from the solution on cooling in flat, red needles. Heated 
 to 100° it gives up chlorine. It is decomposed by water into hydro- 
 chloric acid and potassium dichromate : 
 
 2Cr02<Qj^ + H 2 O == K 2 Cr 207 + 2 HCI. 
 
 Bromides and iodides of chromic acid are not known, but Chromyl Fluoride, Cr02Fl2, 
 a red-colored and very volatile liquid, which is decomposed by water with the formation 
 of chromic and hydrofluoric acids, is obtained by heating a mixture of lead chromate and 
 calcium fluoride with concentrated sulphuric acid in platinum or lead vessels. It etches 
 glass. It was formerly thought to be chromium hexafluoride, CrFlg. 
 
 The following reaction is very characteristic for chromic acid : On 
 adding hydrogen peroxide to a solution of chromium trioxide, or tlie 
 acidified solution of a chromate, the red liquid is colored blue. On 
 shaking the blue solution with ether, the latter withdraws the blue com- 
 pound which O. F. Wiede [Ber. 31 (1898), 516] considers to be the 
 anhydride of perchromic acid, 01*209, together with hydrogen per- 
 oxide. This chemist succeeded in preparing alkali salts of the perchromic 
 acid but always combined with hydrogen peroxide. They are violet in 
 color and are readily decomposed ; their formula is probably 
 
 MeO . CrO^ -f H 2 O 2 , 
 and that of the hypothetical acid : 
 
 CrO^. OH. 
 
 They lose oxygen and become dichromates. Silver nitrate produces a 
 brownish-red precipitate of silver chromate, AggCrO^, with aqueous 
 solutions of the chromates. This silver salt dissolves very readily in 
 ammonia and in the mineral acids. 
 
inor(;anic chemistry. 
 
 
 2. MOLYBDENUM. 
 
 Mo = 96.0. 
 
 Molybdenum is of comparatively rare occurrence in nature ; usually as molybdenite 
 (MoS^) andwulfenite (PbMo04). Scheele, in 1778, obtained molybdic acid from molyb- 
 denite which had until then been considered grapliite. lljelm, in 1790, prepared molyb- 
 denum as a silver-white metal, of specific gravity 8.6, by igniting the chlorides or oxides 
 in a stream of hydrogen. It is very hard, fuses at a higher temperature than platinum, 
 and is as ductile and as susceptible of polish as iron. Like the latter it takes up carbon at 
 1500° and becomes so hard that it will scratch glass. When heated in the air it oxidizes 
 to molybdenum trioxide. It is soluble in concentrated sulphuric and nitric acids. It is 
 also converted by the latter into in.soluble molybdenum trioxide, M0O3. 
 
 Molybdenum forms derivatives of the most different types, so that it is apparently a 
 bivalent, trivalent, ([uadrivalent, etc. , and even an octivalent metal. The vapor density of its 
 chloride, MoClg, has been determined and in it the metal is positively quinquivalent. In 
 its most stable oxide, MoO.^, it is apparently .sexivalent and in the stable sulphide MoSj 
 quadrivalent ; but we are ignorant as to the constitution of these bodies as well as of 
 their molecular magnitudes. Only the most important will be described. 
 
 Molybdenum Dichloride, MoClj, resulting from the trichloride, MoCl,, when 
 heated in a stream of carbon dioxide (together with M0CI4), is a bright yellow, non- 
 volatile powder. It is converted by potassium hydroxide into the hydrate, Mo(OH)2, 
 a black powder. 
 
 The molecular formula of this chloride (and that of the corresponding bromide) must 
 apparently be trebled — Mo^Clg, because dilute alkalies change it to chloromolybdic 
 hydroxide, Cl4Mo3(OII).2. The latter behaves like a diacid base having the radical 
 (MogCl^)!!. Concentrated haloid acids convert it into salts of the formula (MogCl^)!^.^, 
 in which only the two R-halogen atoms are precipitated by silver nitrate. In the 
 chloride (Mo3Cl4)Cl2 therefore only two chlorine atoms are precipitated as silver chloride. 
 On the other hand, chloromolybdic hydroxide (a bright yellow amorphous powder) is also 
 a feeble acid. It dissolves without change in dilute alkalies and is reprecipitated from 
 its yellow solution by weak acids. 
 
 The iron cyanide radical [Fe(CN)g] and many other radicals containing halogens, 
 derived from cobalt, chromium and platinum, conduct themselves like (M03CI4). In 
 these compounds, the halogens belonging to the radicals are obscured — protected — from 
 the reagents with which we are accustomed to see them react in their hydrogen and other 
 derivatives. This is true also of the other parts of the radicals, of the molybdenum in 
 (M03CI4), of the iron in [Fe(CN)g], etc., etc. These metals can be no longer recognized 
 by the reagents with which they are ordinarily detected. These new relations and their 
 explanation, by the theory of electrolytic dissociation, will be again considered in con- 
 nection with the iron-cyanogen compounds. 
 
 Molybdenum Trichloride, M0CI3, produced by gentle heating (at 250°) of molyb- 
 denum pentachloride, M0CI5, in a current of hydrogen or carbon dioxide, is a reddish- 
 brown powder, resembling red phosphorus, which, when strongly ignited, yields a dark- 
 blue vapor. It dissolves with a beautiful blue color in concentrated sulphuric acid, upon 
 heating, with an emerald-green color. Potassium hydroxide converts it into the hydroxide, 
 Mo( 011)3, which forms .salts with acids. The ignition of the hydrate affords the black 
 oxide, MogOg. Strong heating of the trichloride in a current of carbon dioxide leaves 
 molybdenum dichloride, MoCl.^, and it sublimes. 
 
 Molybdenum Tetrachloride, M0CI4, is a brown, crystalline powder, which appears 
 to break up by eva])oration into pentachloride, M0CI5, and trichloride, M0CI3. Molyh- 
 denimi disulphide, MoS^, is i)roduced by the ignition of the trisulphide, M0S3, away from 
 air. It is a .shining black powder, which occurs native as molybdenite, in hexagonal, 
 graphite-like crystals, with a specific gravity of 4.6. 
 
 Molybdenum Pentachloride, Mod,-, is prepared by heating molybdenite (MoSj), 
 or molybdenum in dry chlorine gas. It is a metallic, shining, black, crystalline mass, 
 fusing at 194° and distilling at 268° ; its vapor density corresponds to the molecular 
 formula MoC^. It fumes and deliipiesces in the air, and dissolves in water with hissing. 
 
MOLYBDENUM. 383 
 
 Its aqueous solution has a brown color. It dissolves in absolute alcohol and ether with 
 a dark-green color. 
 
 The hexachloride, MoClg, is not known, but the oxychlo 7 'ides, M0OCI4, and M0O2CI2, 
 are. 
 
 Bromine forms perfectly analogous compounds with molybdenum. 
 
 Molybdenum Trioxide, MoO^, results on roasting metallic molybdenum or the sul- 
 phide in the air. It is a white, amorphous mass, which turns yellow on heating ; it fuses 
 at a red heat and then sublimes in brilliant rhombic plates and needles. It is insoluble in 
 water and acids ; but dissolves readily in the alkalies and ammonium hydroxide. When 
 fused with the alkaline hydroxides or carbonates, salts are produced, partly derived from 
 the normal acid, H2MoO^, and partly from the poly-acids, and correspond to the poly- 
 chromates : 
 
 K2M0O4, K2M02OP K2Mo30jg, Na2]Mo^Oj3, KgMo^024, etc. 
 
 The ammonhwi salt, (NH^)2Mo04, is obtained by dissolving the trioxide in concen- 
 trated ammonium hydroxide. In the laboratory it serves as a reagent for phosphoric 
 acid. Alcohol causes it to separate out of its solution in crystals ; upon evaporation, 
 however, the salt, (NH^lgMo.Og^ + 4H2O, crystallizes out. Both salts are decomposed 
 by heat, leaving molybdenum trioxide. 
 
 Nitric acid gradually .separates a yellow crystalline powder from the alkali molybdate.s, 
 MoOi^OH)^, which is only slightly soluble in water. It loses water on drying and 
 changes to the normal hydrate Molybdic Acid, H2MoO^. By the action of reducing 
 agents, such as zinc, stannous chloride, and suljrhur dioxide, molybdic acid solutions 
 are colored first blue, then green, then brownish-red, and finally brown. This is due to the 
 formation of different lower oxides. The first blue-colored reduction product is formed 
 by warming a molybdic acid solution with metallic molybdenum ; it has the formula 
 Mo.^Og ( = 2M0O3 . Mo 02)‘ product, by the action of tin and hydrochloric 
 
 acid, is molybdenum sesquioxide, MO2O3. Its compounds are usually yellow in color. 
 Potassium permanganate converts all these lower oxides into molybdic acid. 
 
 Hydrogen peroxide converts molybdic acid and its salts into derivatives richer in 
 oxygen ; these were supposed to be derived from a permolybdic acid, HMoO^. The 
 latest researches have demonstrated that their oxygen content varies and that they can- 
 not be referred to a common parent body. It is thought that they contain, like ozone 
 and hydrogen peroxide, oxygen atoms = ©2 in union with one another, therefore they have 
 been designated ozo molybdates (Muthmann and Nagel, Ber. 31 (1898), 1836). 
 
 Molybdic acid can also form poly-acids with phosphoric and arsenic acids, e. g\, 
 H3PO4 . 11M0O3. These complex pho.sphomolybdic acids are distinguished by the fact 
 that they form salts insoluble in dilute acids with the metals of the potassium group, with 
 ammonia, with thallium and with organic bases. Sodium and lithium salts do not yield 
 precipitates with these acids. On adding a solution containing phosphoric (or arsenic) 
 acid to the nitric acid solution of ammonium molybdate, there is produced a yellow 
 crystalline precipitate of ammonium phospho molybdate, (NH^)3PO^ . 11M0O3 6H2O. 
 
 This reaction serves for the detection and separation of phosphoric acid. 
 
 Molybdenum Trisulphide, M0S3, is thrown down as a brown precipitate from acidu- 
 lated molybdenum solutions by hydrogen sulphide. It dissolves in alkaline sulphides 
 forming sulpho-.salts. Ignited away from air it is converted into black molybdenum disul- 
 phide, M0S2, which occurs native as molybdenite. 
 
 In addition to these molybdenum compounds in which the element is sexivalent 
 there is a 
 
 Molybdenum Tetrasulphide, MoS^. From this are derived persulphomolybdic 
 acid, H2M0S5, and its salts, e. g., K2M0S5 and KHM0S5. 
 
3^4 
 
 INORGANIC CHEMISTRY. 
 
 3. TUNGSTEN. 
 
 W 184. 
 
 Tungsten is found in nature in tlie tungstates: as wolframite, (Fe,Mn)WO^, as 
 scheelite, CaWO^, and as stolzite, PbWO^. 
 
 d'he metal is obtained, like molybdenum, by the ignition of the oxides or chlorides in 
 a stream of hydrogen, in the form of a black powder, or in steel-gray crystalline leaflets, 
 having a specific gravity of 19. It is very hard and difficultly fusible. It becomes tri- 
 oxide when ignited in the air. 
 
 Tungsten forms the following chlorides: WCl,^, WCl,, WCl- and WCl^. 
 
 The Pentachloride, WCl., is obtained by the distillation of WCl^ in a current of 
 hydrogen or carbon dioxide, and consists of shining, black, needle-like crystals. It 
 melts at 248° and boils at 275°, forming a greenish-yellow vapor, with the density 360 
 (WClg— 361.25 ; — 32). It affords an olive-green solution and a blue oxide, W^O^, 
 
 with water. It dissolves with a deep-blue color in carbon bisulphide. 
 
 Tungsten Hexachloride, WCl,j, is produced when the metal or a mixture of wolf- 
 ramite with carbon is heated in a current of chlorine. It forms a dark-violet, crystalline 
 mass, which melts at 275° and boils at 346°. The vapor density corresponds to the for- 
 mula WClg. It dissolves in carbon bisulphide with a reddish-brown color; it forms 
 tungsten trioxide, W().„ with water. The other chlorides are obtained from it by heating 
 them alone or in a current of hydrogen gas. 
 
 The Oxychloride, WOCl^, consists of red crystals, fusing at 210° and boiling at 
 229° ; its vapor density equals 340 (WOCl^ = 341.8;. The Dioxychloride, W02C1.^, 
 sublimes in bright yellow, .shining leaflets. They are produced from the other chlorides 
 by the action of water. 
 
 Tungsten Trioxide, WO.^, is thrown out of the hot .solution of tungstates by nitric 
 acid, as a yellow precipitate, insoluble in acids, but dissolving readily in potassium and 
 sodium hydroxides. Tungstic Acid, WO(OH)4, is, however, juecipitated from the 
 cold solution, but on standing over sulphuric acid it becomes \V02(011)2 and at 100° 
 passes into ditungstic acid, H2W20^ := W205(0H)2. 
 
 When tungstic acid is reduced in hydrochloric acid solution by zinc it first becomes 
 blue (formation of W'^205), and then brown, when the salt of the dioxide, WO2, is formed. 
 Potassium permanganate oxidizes this to tungstic acid. 
 
 The salts of tungstic acid are perfectly analogous to the molybdates, and are derived 
 from the normal acid or the poly-acids. The normal sodium salt, Na2W04 -f- 2H2(), 
 and the so-called sodium meta-tungstate, Na.^W^O^, -t- loIL^O, are applied practically. 
 Materials saturated with their solutions do not burst into a flame, but smoulder away 
 slowly. 
 
 Tungstic acid and the tungstates behave tow'ard hydrogen peroxide just like molybdic 
 acid. 
 
 The reduction of the alkaline tungstates (by fu-sion with tin, etc.) affords peculiar com- 
 pounds, e. g., or Nax( W03)y ; these have various colors, possess metallic luster, 
 
 and are applied as tungsten bronzes. 
 
 Tungstic acid also combines with phosphoric and arsenic acids, forming derivatives 
 analogous to those of molybdic acid with the same acids. 
 
 d'he metal is used in the manufacture of tungsten steel ; a slight quantity of it increases 
 the hardness of the latter very considerably. It was first prepared in 1783 by J. and F. 
 d’Elhujar. 
 
 4. URANIUM. 
 
 U = 239.5. 
 
 In nalure it occurs chiefly as uraninite, a compound of uranic and uranous oxides, 
 
 U0,2.2U()3= u/v 
 
 d’he metal, first obtained in 1840 by Peligot on heating uranous chloride with sodium, 
 lias a silver-while color and a sjiecific gravity of 18.7. When heated in the air it burns to 
 uranous-uraiiic oxide. Its specific heat equals 0.0267, 4 s atomic volume is therefore 
 
URANIUM. 
 
 385 
 
 6.6. It melts at about I5cx)°. 'Iliere are two series of uranium compounds. In the one, 
 the metal is a tetrad UX^ ; these uranous or urano-compounds are very unstable, and pass 
 readily into the uranic or derivatives of sexivalent uranium. Uranous oxide is of a basic 
 nature, and only forms salts with acids. 
 
 'J'he compounds of sexivalent uranium are called the uranic compounds. The oxide, 
 UC);^, and the hydroxide, U().^(OH)2, have a predominant basic character, but are also 
 capable of forming salts witli bases which are called loaiiates. In the salts derived from 
 acids, e. g.^ U0.2(N03)2 and UO^. SO^, the group UO.^ plays the role of a metal ; it is called 
 tiranyl, and its salts are termed iiranyl salts. They may also be regarded as basic salts. 
 
 URANOUS COMPOUNDS. 
 
 Uranous Chloride, UCI4, is obtained by heating metallic uranium in a stream of chlo- 
 rine, or uranous oxide in hydrochloric acid. It consists of dark-green octahedra with 
 metallic luster. It volatilizes at a red heat, forming a I'ed vapor, whose density agrees 
 with the formula UCI4. It deliquesces in the air, and dissolves with hissing in water. 
 Uranous hydroxide remains when the solution is evaporated. 
 
 Uranous Oxide, UO2, is formed when the other oxides are heated in a current of 
 hydrogen. It is a black powder, which becomes uranous-uranic oxide, UO2. 2UO3, when 
 heated in the air. 
 
 Uranous oxide dissolves with a green color in hydrochloric and concentrated sulphuric 
 acids. Uranous sulphate, U( 804)2 + ^^12^5 consists of green crystals. From the salts 
 the alkalies precipitate the voluminous, bright green uranons hydroxide, U(OH)4, which 
 becomes brown on exposure. Uranous salt solutions are distinguished by a very remark- 
 able absorption spectrum. 
 
 SEXIVALENT URANIUM COMPOUNDS. 
 
 Uranium Hexachloride, UClg, has not been obtained, but the oxychloride, 
 UO2CI2 (uranyl chloride), exists ; it is made by heating UO2 in dry chlorine gas, or 
 by the evaporation of uranyl nitrate with hydrochloric acid. It is a yellow crystalline 
 mass, deliquescing in the air. 
 
 Uranic Oxide, UO3 or Uranyl Oxide, UO2. O, is a yellow powder, and is obtained 
 by heating uranyl nitrate to 250°. When warmed with nitric acid it becomes uranyl 
 hydrate or tira?iic acid, U02(OH)2, which is also yellow-colored. 
 
 Uranyl Nitrate, U02(N03)2, results from the solution of uranous or uranic oxide, or 
 more simply of uraninite in nitric acid. It crystallizes with six molecules of water, in 
 large, greenish-yellow prisms, which are readily soluble in water and alcohol. On add- 
 ing sulphuric acid to the solution, Uranyl Sulphate, UO2SO4 -f 6H2O, crystallizes out, 
 on evaporation, in lemon-yellow needles. 
 
 If sodium or pota.ssium hydroxide be added to the solutions of uranyl salts, yellow pre- 
 cipitates of the uranates, K2U20^ and Na2U207, are obtained. These are soluble in 
 acids. In commerce the sodium salt is known as iiraniiiin yellow, and is employed for 
 the yellow coloration of glass (uranium glass) and porcelain. The uranates can be 
 obtained in crystalline form, by igniting uranyl chloride with alkali chlorides in the pres- 
 ence of ammonium chloride. Nascent hydrogen reduces uranic to uranous compounds. 
 
 The so-called uranic-uranous oxide, which constitutes uraninite, and is formed by the 
 ignition of the other oxides in the air, must be viewed as uranous uranate, 2U()3. UO., = 
 VI IV 
 
 (U02.0.2)2U. _ 
 
 Many uranium salts exhibit magnificent fluorescence. The oxide colors glass fluxes a 
 beautiful greenish-yellow (uranium gla.ss). Uranous oxide, UO2, imparts a beautiful 
 black color to glass and porcelain. 
 
 33 
 
386 
 
 INORGANIC CHEMISTRY. 
 
 Besides these compounds, in which uranium aj)pears to be quadrivalent and sexivalent, 
 it also yields a pentachloride, UCI5, like molybdenum and tungsten. 'I he same 
 results on conducting chlorine gas over a moderately heated mixture of carbon with one 
 of the uranium oxides. It consists of dark needles which, in direct light, are metallic 
 green, but in transmitted, ruby-red. It delicpiesces in the air to a yellowish-green liquid ; 
 upon heating it is dissociated into uranous chloride and chlorine (at 120-235°). 
 
 d'here is also a Ictroxide, U()^, which, like the trioxide, UO^, yields salts with the 
 bases. It corresponds to molybdenum tetrasulphide MoS^ (p. 383). 
 
 MANGANESE. 
 
 Mil = 55.0. 
 
 According to its atomic quantity, manganese bears the same relation to 
 the elements of the chlorine group as chromium does to the elements of the 
 sulphur group. The relationship manifests itself distinctly in the higher 
 states of oxidation. Permanganic acid, HMnO^, and its anhydride, 
 Mn^O^, are perfectly analogous to perchloric, HCIO4, and periodic, HIO4, 
 acids and their anhydrides, CbO^ and l20^. The permanganates and the 
 perchlorates are very similar, and for the most part are isomorphous. The 
 manganese in them appears to be septivalent, like the halogens in their 
 highest state of oxidation. The similarity of manganese to the halogens 
 is restricted to this one point of resemblance. In the rest of its deriva- 
 tives, manganese shows great resemblance to the elements standing in the 
 same horizontal series of the periodic system, viz., with chromium and 
 iron (p. 393). Like these two elements, it forms compounds of the types : 
 
 MO (M( 0 H) 2 .MX 2 ), M.Pa (M(On)3.MX3, MO3 (MOp^). 
 
 1. In the manganous derivatives, MnX2, the metal is bivalent. These 
 salts are the more stable, and comprise the most common manganese 
 compounds. They resemble and are usually isomorphous with the oi4S 
 salts of iron and chromium, and the salts of metals of the magnesium 
 group (p. 313). 
 
 2. The manganic compounds, MnX3, are similar to and isomorphous 
 with the ferric, chromic and aluminium derivatives; they are, however, 
 less stable, and easily reduced to the manganous state. 
 
 3. The derivatives of manganic acid, 
 
 II^MnO^ = Mn02(0H)2, 
 
 in which manganese is sexivalent, are analogous to those of ferric 
 (IL^FeOJ and chromic acid (H2C1O4), and, of course, to those of 
 sulphuric acid (H2SOJ. 
 
 4. Manganese also yields derivatives which can be referred to the type 
 M.2O7. These are the perma}iganates,^\\\(^ correspond in formula to the 
 persulphates. 'Phe manganese in them is apparently sei)tivalent, but we 
 liave no definite knowledge of its structure, not even as to whether it cor- 
 res|)(;nds to that of the ])crsul])hates (]). 297). 
 
 Conseciueiill V, in manganese we plainly observe how the similarity of 
 the elements in their compounds is influenced by the valence (see p. 
 
MANGANOUS COMPOUNDS. 
 
 387 
 
 329). In the ma.nga.uous condition, manganese, like the elements of the 
 magnesium group, has a rather strong basic character, which diminishes 
 considei'ably in the inangan/V state. Sexivalent manganese has a metal- 
 loidal acidic character, and, in manganic acid, approaches sulphur. By the 
 further addition of oxygen, manganese finally (in permanganic acid) ac- 
 quires a metalloidal character and resembles the halogens. We have 
 already noticed that many other metals, especially chromium and iron, 
 exhibit a similar behavior. Osmium in its tetroxide, OsO^, wholly resem- 
 bles the halogens. 
 
 On the other hand, the metalloidal and the weak basic metals acquire a strong basic, 
 alkaline character, by the addition of hydrogen, or hydrocarbon groups (CH.^, C2Hg). 
 The univalent groups, NH^ (ammonium), P(CH3)^ (tetramethyl phosphonium), 
 
 (triethyl sulphine), Sn(C2H5)3 (tin triethyl), etc., are of metallic nature, because their 
 hydroxides, P(CH3)4.0H, 8(02^13)3. OH, 80(02113)3 . OH, are perfectly similar to the 
 hydroxides (KOH, NaOH) of the alkali metals. 
 
 Manganese is widely distributed in nature. It is found native in 
 meteorites. Its most important ores are pyrolusite, Mn02, hausmannite, 
 Mn304, braunite, Mn203, manganite, Mn203.H20, and rhodochrosite, 
 MnC 03 . 
 
 Metallic manganese is obtained by igniting the oxides with aluminium 
 powder (p. 347). It has a grayish-white color, is very hard, and fuses 
 with difficulty (at about 1900°), and a little above this temperature it 
 volatilizes. Its specific gravity is 7. 2-8.0. It oxidizes readily in moist 
 air. It decomposes water on boiling, and, when dissolved in acids, 
 forms manganous salts. 
 
 The heat of formation of the most important manganese compounds corresponds to 
 the symbols : 
 
 (Mn,0,H20) = 95 (Mn,Cl2) =^112 (Mn,Cl2,Aq) = 128 
 
 (Mn,02,H20) = 1 16 (Mn,S,OJ = 245 (Mn, 0 „K) = 195. 
 
 MANGANOUS COMPOUNDS. 
 
 Manganous Oxide, MnO, results from ignition of the carbonate, 
 with exclusion of air, and by heating all manganese oxides in hydrogen. 
 It is a greenish, amorphous powder, which, in the air, readily oxidizes 
 to manganous-manganic oxide, MiigO^. 
 
 Manganous Hydroxide, Mn(OH)2, is a voluminous, white precipi- 
 tate, formed by the alkalies in manganous solutions. When exposed to 
 the air, it oxidizes quickly to brown manganic hydroxide, Mn2(OH)3. 
 
 Manganous salts usually have a ])ale, reddish color, and are formed by 
 the solution of manganese or oxides of manganese in acids. 
 
 Manganous Chloride, MnCl2, crystallizes with four molecules of 
 water in reddish tables. When heated, it decomposes into the oxide, 
 
388 
 
 INORGANIC CHEMISIRV. 
 
 liydrochloric acid and water. Anhydrous manganous cldoride is ])re- 
 pared by igniting the double salt 
 
 MnCl,.2Nll,Cl 4 up 
 
 (see Magnesium Chloride), or by heating manganese oxides in hydro- 
 chloric acid gas; it is a crystalline, reddish mass, which deliquesces in 
 the air. 
 
 Manganous Sulphate, MnSO^, crystallizes below 6° with seven 
 molecules of water (like magnesium and ferrous sulphates), at ordinary 
 temperatures with sH.p (like copper sulphate); and above 25° with 
 qHjO ; upon drying, the salt 
 
 Mn.SO, + 11,0 
 
 remains; the last molecule of water does not escape until at 280°. It 
 forms double salts with the alkaline sulphates, e. g. : 
 
 MnSO^. K,SO^ -I- 611,0. 
 
 Manganous Carbonate, MnCO,, exists in nature as rhodochrosite, 
 and is precipitated by alkaline carbonates from manganous solutions, as 
 a white powder, which turns brown on exposure to the air. 
 
 Manganous Sulphide, MiivS, is found in nature as alabandite or 
 manganese-blende. Alkaline sulphides precipitate a flesh colored sul- 
 phide from manganous solutions. It becomes brown in the air. 
 
 MANGANIC COMPOUNDS. 
 
 Manganic Oxide, Mn203, manganese sesquioxide, is a brown or black 
 powder produced by the ignition of all the manganese oxides in a current 
 of oxygen gas. It occurs as braiinite in dark-brown quadratic crystals. 
 
 IV III 
 
 Manganic Hydroxide, Mn2(0H)g or Mn(OH)3, manganic hydrate 
 separates as a dark-brown precipitate on exposing ammoniacal solutions 
 of manganous salts to the air. It dissolves in cold hydrochloric acid to 
 a dark-brown liquid, containing, in all probability, 7 fiangani€ chloride, 
 MnCl3. When this is heated it decomposes into manganous chloride 
 and chlorine. 
 
 IV 
 
 Manganite, occurring in iron-black crystals, is the hvdroxide, Mn202- 
 
 III 
 
 (0H)2 or MnO.OH. 
 
 Manganous-manganic Oxide, MnO. ^111,03. consti- 
 
 tutes the mineral hausmannite, crystallized in dark-gray quadratic octa- 
 hedra, and is obtained as a reddish-brown powder by the ignition of all 
 other manganese oxides in the air. It reacts with hydrochloric acid, 
 according to the equation : 
 
 MiqO^ + 8IIC1 .iMnCI, ] 4ll,0 -f- Cl,. 
 
MANGANIC COMPOUNDS. 
 
 389 
 
 Since manganic oxide is quadratic in its crystallization, while all other sesquioxides 
 (like corundum and hematite) are rhombohedral, and since the first is decomposed by 
 dilute nitric and sulphuric acids into manganese dioxide, MnOg, nnd a manganous salt, 
 it has been generally supposed that manganic oxide is not a sesquioxide, but rather a 
 compound of the dioxide with manganous oxide : 
 
 MnOj. MnO = MnO<( 0 ^]yjj^ 
 
 Hausmannite is quadratic, while other metallic oxides of the type (the spinels, 
 
 Y>. 351 and p. 377, and magnetite) are isometric; therefore the former is not considered 
 a compound of manganese sesquioxide and protoxide ; 
 
 III 
 
 Mn.A-MnO = “;;0;g>Mn, 
 but as manganous oxide and the dioxide : 
 
 Mn02.2MnO = MnO<Q-^”>0. 
 
 This is shown by its behavior toward dilute nitric and sulphuric acids, which decom- 
 pose it into manganese dioxide and two molecules of manganous oxide. Consult Phanke, 
 Jr. f. prakt. Ch. 36 (1887), 451. Chrysoberyl, unlike other .spinels, is trimetric, and 
 other reactions clearly prove (chiefly their deportment wdth concentrated sulphuric acid) 
 that manganic and mangano-manganic oxides are to be regarded as sesquioxide derivatives. 
 
 Manganic oxide, like the other sesquioxides, is a very feeble base, 
 which does not form salts with dilute or weak acids, and by separation 
 of oxygen reverts to the manganous condition. Its salts are very un- 
 stable. 
 
 Manganic Sulphate, Mn2(SOj3, is obtained by the solution of man- 
 ganic oxide, hydroxide, or, better, manganous-manganic oxide in con- 
 centrated sulphuric acid. When the last oxide is employed manganous 
 sulphate also results. The best procedure is to heat the hydrate of man- 
 ganese dioxide (see p. 390) with concentrated sulphuric acid to 168°, 
 when the sulphate will separate as an amorphous, dark-green powder. It 
 dissolves with a dark-red color in a little water. It is said to form alu 7 ns, 
 with potassium and ammonium sulphates, e. g., Mn2(SOj3 . K^SO^ -|- 
 24H2O, but Christensen’s researches make this rather doubtful. 
 
 Manganese Dioxide, Mn02, peroxide. This is the mineral pyro- 
 lusite, occurring in dark-gray radiating masses, or in almost black rhom- 
 bic prisms, which possess metallic luster. When gently heated it is con- 
 verted into oxide, by strong ignition into manganous-manganic oxide: 
 
 3Mn03 = Mn 30 , d- O 2 . 
 
 It is used for making oxygen. Chlorine escapes when it is warmed 
 with hydrochloric acid (p. 49) : 
 
 Mn02 + 4IICI = MnClj + 2II2O + Cl,. 
 
390 
 
 INORGANIC CFIKMISTRY. 
 
 'I'he dioxide may be obtained artificially by heating manganous nitrate 
 to 150-160°. Its hydrates, MnO.^.H./) and Mn(b. 21 1 , 0 , are jirodiired 
 on adding a hypochlorite to the solution of a manganous salt, or if chlo- 
 rine be conducted through a solution of manganese containing sodium 
 carbonate, or by adding potassium permanganate to a boiling solution of 
 a manganous salt, d'he precijiitated dioxide dissolves in cold hydro- 
 chloric acid, without liberating chlorine, as manganese tetrachloride, 
 MnCl^, is probably formed; when heat is applied it breaks down into 
 manganous chloride, MnCIa, and chlorine, Clj. This deportment would 
 indicate that manganese is quadrivalent in the dioxide (p. 259). Man- 
 ganese dioxide also unites with bases, yielding the so-called fnangam'tes, 
 e.g., BaMn^O^and K.^MiigO^. 
 
 Manganese ])eroxide (also Mn.,03 and Mn30^) serves chiefly for the manufacture of 
 chlorine gas, and it is, therefore, important from a technical ])oint to estimate the ([uan- 
 tity of chlorine which a given dioxide of manganese is able to set free. This is done by 
 boiling the oxide with hydrochloric acid, conducting the liberated chlorine into a potas- 
 sium iodide solution, and determining the separated equivalent amount of iodine by means 
 of sodium hyposulphite. Or the oxide is heated in a flask with oxalic and sulidiuric 
 acids, when the oxalic acid is oxidized to carbon dioxide, and from the quantity of this 
 set free we can calculate the quantity of active or available oxygen in the manganese oxide : 
 
 MnO^ + H2C,0, + H.,SO^ = MnSO^ + 2II2O d- 2CO.,. 
 
 In the preparation of chlorine the manganese is found in the residue as manganous 
 chloride. With the relatively high value of pyrolusite, it is important for trade that the 
 peroxide be recovered from the residue. This regeneration is at present largely executed 
 by the method proposed by Weldon, according to which the manganous chloride, contain- 
 ing an excess of hydrochloric acid, is neutralized with lime, the clear liquid brought into 
 a tall iron cylinder (the oxidizer), milk of lime added and air forced in. The mixture 
 becomes warm, and so-called calcium manganite, CaMnOg = CaO . Mn02, is precipi- 
 tated as a black mud (Weldon’s mud) : 
 
 MnCb -f 2CaO -f O = CaMnOg + CaCb,. 
 
 The calcium chloride solution is run off, and the residual calcium manganite employed 
 for the preparation of chlorine, when it conducts itself as a mixture of Mn02 -f CaO. 
 
 COMPOUNDS OF MANGANIC AND PERMANGANIC ACIDS. 
 
 When oxygen compounds of manganese are heated in the air in con- 
 tact with potassium hydroxide, or, better, with oxidizing substances, like 
 niter or potassium chlorate, a dark-green amorphous mass is produced, 
 which dissolves in cold water with a dark-green color. When this solu- 
 tion is evaporated under the air-pump, dark-green, metallic, rhombic 
 jirisms of potassium manganate, K^MnO^, crystallize out. This salt 
 is isomorphous with potassium sulphate and potassium chromate. It suf- 
 fers no change by solution in potassium or sodium hydroxide, but is de- 
 comj^osed by water, brown hydrated manganese dioxide separating, and 
 the green solution of the manganate changing into a dark-red solution of 
 the jiermanganate, KMnO^: 
 
 3K2Mnf)^ y 3H2O = 2KMnO^ -j- Mn02.H2C -f 4KOII. 
 
POTASSIUM PERMANGANATE. 
 
 391 
 
 A similar conversion of the green manganate into red permanganate 
 occurs more rapidly under the influence of acids or chlorine : 
 
 3K2MnO^ + 4HNO3 = 2KMn04 + MnO^ + 4KNO3 + 2H3O ; 
 
 K^MnO^ + Cl = KMnO^ -f KCl. 
 
 The red solution again assumes a green color on the addition of a con- 
 centrated solution of an alkaline hydroxide. 
 
 Owing to this ready alteration in color the solution of the manganate 
 is called chameleon 77 imeraL 
 
 Potassium Permanganate, KMnO^, is best prepared by conducting 
 carbon dioxide into the manganate solution until the green color has 
 passed into a red. Hydrated peroxide of manganese is i)recipitated. 
 When the solution is concentrated the salt crystallizes in dark-red 
 rhombic prisms isomorphous with potassium perchlorate, KCIO^. It is 
 soluble in twelve parts of water at ordinary temperatures. 
 
 The permanganate solution is a strong oxidizing agent. If the oxida- 
 tion takes place in the presence of a sufficient quantity of acid the per- 
 manganate is reduced to a faintly colored manganous salt : 
 
 2KMnO, + 3H2SO, = 2MnSO, + K^SO^ + 3H,0 + 5O. 
 
 When a permanganate solution is added to an acidulated ferrous solution, 
 the former is decolorized, and there results a faintly yellow-colored solu- 
 tion of ferric and manganous salts: 
 
 2KMn04 + loFeSO^ + SH^SO^ = 2MnSO^ + 5Fe2(SOd3 + SHp + K.^SO^. 
 
 Hence the solution of this salt serves for the volumetric estimation of 
 ferrous salts. 
 
 In neyfral ox alkaline the permanganate in oxidizing is decom- 
 posed with the separation of manganese oxides : 
 
 2KMn04 -f HjO = 2Mn02 -f- 2KOH -|- 3O. 
 
 In the same manner, the permanganate oxidizes and destroys many 
 organic substances, therefore its solution cannot be filtered through paper 
 without decomposition. It serves as a disinfectant. 
 
 The permanganate is also reduced by hydrogen peroxide in acid solu- 
 tion (p. 103). 
 
 The remaining permanganates are similar to and isomorphous with the 
 perchlorates. The sodium salt is very soluble in water, and does not 
 crystallize well. 
 
 Very cold sulphuric acid added to dry permanganate causes the 
 separation of Manganese Heptoxide, Mn^O^, an oily, dark-colored 
 liquid. By careful warming it is converted into dark-violet vapors, 
 which explode when heated rapidly. Manganese heptoxide has a 
 violent oxidizing action ; ])aper, alcohol and other organic matter are 
 inflamed by mere contact with it. 
 
392 
 
 inor(;anic cukmis'ikv. 
 
 METALS OF GROUF" VIII. 
 
 The last group of the periodic system comprises the elements: 
 
 Fe — 56 Ni = 58.7 Co = 59 
 
 Ru — 101.7 Kh = 103.0 Pel = 106 
 
 Os — 191 Ir 193-0 Pt = 194.8. 
 
 These elements are the middle members of the three great periods, and 
 they have no analogues in the two small jjeriods (]). 245). 
 
 As regards both atomic weights and i)hysical and chemical dei)ortment, 
 these elements constitute a transition from the ])receding members of the 
 great periods (Cr and Mn ; Mo; W) to the next following members (Cu, 
 Ag, An, and Zn, Cd, Hg). The elements standing side by side (heterol- 
 ogous) and belonging to the same periods are very similar in their 
 physical properties, and show, e. g., very close si)ecific gravities. They 
 are, therefore, usually arranged in groups, and are distinguished as (i) the 
 iron group (Fe, Ni, Co), with the specific gravity 7. 8-8. 6; (2) thegrouj) 
 of the light platinum metals (Ru, Rh, Pd), with the specific gravity 
 11.8-12. 1 ; and (3) the group of the heavy platinum metals (Os, Ir, Ptj, 
 with the specific gravity 21. 1-22.4. 
 
 On the other hand, the homologous elements (Fe, Ru, Os; Ni, Rh, 
 Ir ; and Co, Pd, Pt) show a like similaritv in their chemical i)roperties, 
 as do the other homologous elements. This re.semblance shows itself 
 chiefly in their combination forms, and, of course, too, in the properties 
 of the compounds (p. 329). We know that the metals of group VI 
 (chromium, molybdenum, tungsten) and of group VII (manganese) form 
 the highest oxides (MeOg and Me20^) having an acidic nature. In the 
 adjacent elements of group VIII (iron, ruthenium, and osmium) we find 
 salts : 
 
 KgFeO^, K2RUO4, K20 s0^, 
 
 derived from the unstable trioxides FeOg, RuOg and OsOg. This acid- 
 forming function disappears in the following members, Ni, Rh, Ir, and 
 Co, Pd, Pt ; their chemical valences diminish rapidly and they attach 
 themselves to Cu, Ag and Au. 
 
 Consequently the whole physical and chemical deportment of the nine 
 elements about to be considered is governed by their position in the 
 periodic system. 
 
 METALS OF THE IRON GROUP. 
 
 'I'he metals of this group, iron, nickel and cobalt, are distinguished 
 from all the other elements by their magnetic ])ro])erties. 
 
 Iron forms three series of compounds after the forms, FeOg, Fe20g and 
 lAO. In its highest combinations iron has an acidic character, and the 
 derivatives of ferric acid (H2FeOJ are perfectly similar to those of 
 
IRON. 
 
 393 
 
 chromic and manganic acids; they are, however, less stable than the 
 latter. Their analogues with cobalt and nickel are unknown. 
 
 The ferrzV compounds, FeXg, are much like the aluminium, chromic 
 and manganic derivatives. They are generally isomorphous with them. 
 They are characterized among iron salts by their relative stability. The 
 highest oxides of cobalt are far less stable, while the higher salts with 
 nickel are unknown. 
 
 Again, iron, nickel and cobalt form ous compounds (FeX^, NiX^, 
 C0X2) in which they appear to be dyads. They resemble the compounds 
 of chromium, manganese, and copper of the same form, and those of the 
 magnesium metals. The ferrous salts are not as stable as the ferric; they 
 are readily oxidized to the latter. 
 
 The cobaltous and nickelous compounds are quite stable, and in this 
 respect these metals ally themselves with co})per and zinc. 
 
 1. IRON. 
 
 Fe = 56.0. 
 
 This very important metal is widely distributed in nature. It is found 
 native on the earth’s surface almost exclusively in meteorites, which 
 generally contain nickel. Iron granules free from the last metal are said 
 to occur in the coal measures of Missouri. Iron is, however, present in 
 great masses in other worlds which (like the sun) are surrounded by an 
 atmosphere of glowing hydrogen, the heat of which is so intense that 
 only the free elements can exist side by side (p. 29). 
 
 The most important iron ores are : magnetite (FegOJ, hematite 
 (Fe203), brown iron ore or limonite (hydrated oxide) and siderite 
 (FeCOg). These ores constitute almost the sole material for the manu- 
 facture of iron ; the sulphide ores, like pyrite, are not so well adapted 
 for iron making. 
 
 When iron is mentioned a distinction must be made between the 
 metal iron, the pure element, and that iron which meets with a universal 
 application in the arts and industries. The latter is really not a metal 
 but an alloy of iron with carbon, and frequently also with silicon, phos- 
 phorus, sulphur, and manganese. Chemically pure iron, the real metal 
 iron, is not applied technically because its production would be too ex- 
 pensive. Hence, a distinction must be made between chemically pure 
 iron and the technical iron. 
 
 A. Chemically pure iron. 
 
 Chemically pure iron is obtained by heating the pure oxide or the oxa- 
 late in a current of hydrogen : 
 
 Fe^Og -f 3II2 = 2Fe -p ; 
 
 to have complete reduction the temperature finally should reach 600°. 
 If the reduction occurs at a red heat, the powder glows in the air, and 
 burns (pyrophoric iron). The strongly ignited powder is not inflam- 
 
394 
 
 inoR(;anic cukmistry. 
 
 nuible. Iron obtained by the ele'clrolysis of ferrous sulj)hate always ccni- 
 lains some hydrogen. 
 
 Chemically pure iron when hammered has a silver-white color, is toler- 
 ably soft, and is magnetizable. Its specific gravity is 7.78. It melts 
 transitionally at 1800° and vaporizes at higher temjieratures. In dry air 
 it alters slowly but rusts rapidly in moist air. Boiling water is decom- 
 posed by finely divided ])ure iron with evolution of hydrogen. It dis- 
 solves readily in acids, and in nitricacid with the liberation of nitric oxide. 
 
 B. Technical Irofi. d'his always contains carbon, d’he carlxm con- 
 tent increases as the temj^erature, at which the ore is fused with the coal, 
 rises. However, even at the most elevated temiieratures iron does not take 
 up more than 5 per cent, of carbon. The latter is either mechanically 
 mixed as graj)hite with the iron or it exists as iron carbide, he„C,„ (<?. g., 
 FegC) chemically combined — alloyed. 
 
 The quantity of carbon and the form in which it exists leads to the 
 varieties cast iron, steel and wrought iron, with their subdivisions. 
 
 I. Cast Iron contains 2.3 or more per cent, of carbon ; its melting 
 temperature rises as the carbon content falls (1050-1250°). It is brittle 
 and cannot be welded or forged. The varieties of cast iron are : 
 
 {ai) White Iron. All of its carbon is chemically combined. It is 
 hard, brittle, and white in color. It is only made for conversion into 
 wrought iron. White iron is distinguished as fibrous and “ spiegeleisen," 
 according to the appearance of its fracture. 
 
 {b) Gray Iron. Only a portion of its carbon is chemically combined 
 with the iron. The greater portion exists as graphite. It is softer, more 
 fusible than the preceding and of a bright gray to black color. It is 
 used in the manufacture of castings, and of malleable iron ; particularly 
 by the Bessemer process. 
 
 In the production of the different kinds of iron not only carbon but 
 also manganese and silicon play an important part. Their presence or 
 absence also determines the nature of the iron. Manganese favors the 
 chemical union of iron and carbon. Hence white iron is usually man- 
 ganiferous. Silicon decomposes the alloy of iron and carbon and favors 
 the production of grajihite. The color of gray cast iron is determined 
 by the size of the graphite plates. If they are very small the iron is fine- 
 grained and its color light; otherwise the aggregation is coarsely crystal- 
 line and the color dark — light and dark gray, and black pig iron. 
 
 H. Wrought Iron and Steel with 1.6 per cent, and less of carbon. 
 (Iron containing 1.6 to 2.3 ])er cent, of carbon is no longer manufactured 
 technically.) It fuses with difficulty, but is malleable and can be welded. 
 It melts at 1400° and higher. The method of preparation and the con- 
 dition in which the malleable iron exists at the conclusion of its manu- 
 facture distinguish it as 
 
 A. Fluid Iron. 
 
 B. Weld Iron. 
 
 'Fhe first has been melted during its manufacture and is therefore free 
 from slag, while the second varietv has only been reduced to a ]xisty form 
 and therefore contains slag. When a glowing piece of weld iron is 
 
IRON. 
 
 395 
 
 suddenly cooled {e. g., by immersion in water) its hardness is appreciably 
 increased, in which case it is called steel, otherwise malleable or wrought 
 iron. The following subdivisions are also known : 
 
 { tempered : Fluid steel, 
 not perceptibly 
 
 tempered : Fluid weld iron. 
 
 { tempered : Weld steel, 
 not perceptibly 
 
 tempered : Wrought iron. 
 
 The hardening is generally influenced by the carbon content ; steel 
 contains from 0.6-1. 6 per cent, and wrought iron less than 0.6 per cent, 
 of carbon. However, the formerly customary distinction between steel 
 and wrought iron on the basis of carbon content alone is no longer ten- 
 able, as at present tempered varieties of iron are produced, containing 
 very considerable quantities of manganese, nickel, silicon, tungsten, and 
 chromium, and very little carbon. 
 
 The classification just given of the technically available kinds of iron was prepared by 
 an international commission at the Centennial Exposition in Philadelphia (1876), but it 
 unfortunately has not been generally adopted. 
 
 Pig iron (crude iron) is always produced in blast furnaces by the reduc- 
 tion of its ores; while wrought iron is rarely made from ores directly, and 
 never in blast furnaces. As a rule, wrought iron is manufactured from pig 
 iron, as its name would imply. 
 
 Metallurgy of Iron. — The extraction of iron from its oxygen ores is 
 based upon the reduction of the same by carbon at a red heat. In the 
 oldest method, the ores were heated with carbon in forges; in this way 
 the excess of air consumed the greater portion of the carbon, and the prod- 
 uct was an iron poor in carbon, wrought iron, a spongy mass, which was 
 then worked under the hammer. The present methods have been de- 
 veloped since the beginning of the fifteenth century. According to 
 these pig iron is first prepared from the ores, and this afterwards con- 
 verted into steel or wrought iron. The smelting of the ores is executed 
 in large, walled blast furnaces, that permit the process to proceed without 
 interruption. The furnaces are filled from openings above, with alter- 
 nating layers of coal, broken ore and fluxes containing silica and lime; 
 the fluxes facilitate the melting together of the reduced iron. The air 
 necessary for the process is blown into the lower contracted portion of 
 the furnace by means of a blowing engine. The combustion of the coal 
 produces carbon monoxide, which reduces the iron oxides to metal : 
 
 FcjOg 3CO = 2Fe -f- 3CO2. 
 
 As the reduced iron sinks in the furnace it comes in contact with the 
 coal, takes up carbon and forms cast iron, which fuses as it sinks lower 
 and collects in the hearth of the furnace. Protracted and strong heating 
 converts the chemically combined carbon into tlie graphitic form, and thus 
 accelerates, with the assistance of silicon, the formation of the gray cast 
 
39 <> 
 
 INORGANIC CHEMISIRY. 
 
 iron. The eanhy impurities of tlie ores coml)inc willi the fiuxcs to a 
 readily fusible s/dg, winch euvel()i)S the fused iron aud protects it from 
 oxidation. 
 
 'I'o convert the cast iron thus produced into wrought iron, carl^on must 
 be withdrawn from it. lu making the wrought iron the cast iron is 
 fused in open hearths (refining process), or in reverberatory furnaces 
 with air access, aud the mass stirred thoroughly until it has become i)asty 
 (puddling process). In this way almost all the carbon is burned to car- 
 bon monoxide and the other admixtures, like silicon, sulphur, and phos- 
 phorus, present in small quantities, are oxidized. The wrought iron is 
 then worked up by rolling, or under the iron hammers (bar iron). If 
 the decarburization be not carried (piite as far the j)roduct will be steel 
 (puddle steel). Such products are not free from slag. Cast iron can be 
 changed into malleable iron by annealing. Castings of white iron are 
 imbedded in ferricoxide powder and ex])osed for some time to a red heat. 
 A portion of the carbon is lost and malleable iron results. 
 
 Steel can be manufactured from wrought-iron by cementation. The 
 iron bars, mixed with fine charcoal, are exposed to a red heat, when the 
 iron takes up carbon from the surface inward, d’he bars are then re- 
 forged, again heated with fine charcoal, and the proce.ss repeated until 
 the mass becomes as homogeneous as po.ssible (cementation steel). Cem- 
 entation is just the opposite of annealing. A more homogeneous steel is 
 obtained by refusion in crucibles (cast steel). 
 
 At present, steel is chiefly prejmred directly from pig iron, by the 
 method invented by Bessemer, in 1855. It consists in blowing air, under 
 high pressure, into the molten iron, until the necessary amount of carbon 
 has been burned out (Bessemer steel). 
 
 An iron rich in silicon (1.5-2 per cent, silicon) is well adapted for the 
 purpose, because by the simultaneous combustion of the silicon the tem- 
 perature is considerably increased. The operation is conducted in a 
 pear-shaped vessel known as the converter. The air is blown in through 
 openings in the bottom. The decarburization by the Swedish method 
 only continues to the point where steel is fused. This is better accom- 
 ]dished by the English method, which removes the carbon so as to con- 
 vert the mass into wrought iron, and then it is again carbonized by add- 
 ing molten spiegeleisen, or by the addition of coke. 
 
 d’he Bessemer process originally was only adapted to crude iron con- 
 taining as little sul]:)hur and phosphorus as possible (at the highest, 
 o 05 per cent, phosphorus), because in this process the phosphorus is not 
 consumed, but remains unaltered in the steel. By a slight, yet very 
 essential alteration, Snelus (1872) and Thomas and Gilchrist (1878) ren- 
 dered it suital)le for iron containing much phosphorus. Their process is 
 now known as the ‘‘ basic process,” and consists in lining the converter 
 with a basic lining material of burnt dolomite. By contact with this sub- 
 stance the iron is comjiletely de])hosphorized and the jfiiosphorus changed 
 to calcium and magnesium phospliates. 
 
 'I'lic basic (lopliospliorizing inctbod lias aUained very great importance because all the 
 jihosphorus contained in iron ores collects in the slag of the converter, d'he latter con- 
 tains as much as 15-20 per cent, of phosphoric acid (and may even be increased to 24 
 
FERROUS COMPOUNDS. 397 
 
 per cent.), existing as calcium phosphate, and this may be applied as a fertilizer in agri- 
 culture (Thomas slag) (p. 305). 
 
 A third method (Siemens-Martin) of making steel consists in puddling 
 the different varieties of iron together, or with iron ores. Siemens-Mar- 
 tin steel is obtained by fusing pig iron with wrought iron (old iron rails). 
 It is much used. Uchatius steel is prepared by fusing granulated pig 
 iron together with some iron ore and pyrolusite in graphite crucibles. 
 
 Ordinary iron, even the purest wire, always contains foreign ingredi- 
 ents, principally carbon and manganese, and minute quantities of silicon, 
 sulphur, phosphorus, nitrogen, nickel, cobalt, titanium and other metals. 
 The quantity of manganese is purposely increased (to 33 per cent.), as by 
 this means the iron acquires valuable technical properties; it becomes 
 more compact and solid. When iron, containing carbon, is dissolved in 
 hydrochloric acid the chemically combined carbon unites with hydrogen, 
 forming volatile hydrocarbons, while the mechanically admixed graphite 
 remains behind. The total carbon is determined by the solution of the 
 iron in bromine water or cupric chloride, when all the carbon remains 
 behind. 
 
 Ordinary iron rusts rapidly in moist air, or it becomes covered with a 
 layer of ferric hydroxide. When ignited in the air it is coated with a 
 layer of ferrous-ferric oxide (Fe^OJ which is readily detached. It burns 
 with an intense light in oxygen. 
 
 In contact with a magnet iron becomes magnetic ; steel alone retains' 
 the magnetism, while cast iron and wrought iron soon lose this property 
 after the removal of the magnet. 
 
 Iron decomposes water at a red heat, with the formation of ferrous- 
 ferric oxide, and the liberation of hydrogen. 
 
 The metal dissolves readily in hydrochloric and sulphuric acids, with 
 evolution of hydrogen ; the latter has a peculiar odor, due to hydrocar- 
 bons which are liberated at the same time. Iron dissolves in nitric acid 
 with evolution of nitrogen oxides. On dipping iron into concentrated 
 nitric acid, and then washing it with water, it is no longer soluble in 
 the acid (passive iron) ; this phenomenon is probabjy due to the produc- 
 tion of ferrous oxide upon its surface. 
 
 FERROUS COMPOUNDS. 
 
 These are produced by the solution of iron in acids, and may also be 
 obtained by the reduction of ferric salts : 
 
 Fe2Ck -h Zn = 2FeCl2 -f ZnCl2. 
 
 In the hydrous state they are usually of a green color; in the air they 
 oxidize to ferric salts: 
 
 2 FeO -|- O = FcjOj. 
 
398 
 
 INORGANIC CHP:MISTKY. 
 
 Ferrous Chloride, FcCl^^, crystallizes from aqueous solutions in green 
 monoclinic prisms, with four molecules of water. These delicjuesce in 
 the air and oxidize. Particularly beautiful crystals are obtained by 
 exposing an alcoholic solution of ferric chloride in a closed vessel to the 
 action of direct sunlight. The alcohol acts as a reducing agent and is 
 oxidized to aldehyde (Organic Chemistry). The anhydrous salt is 
 formed by conducting hydrogen chloride over heated iron. It is a white 
 mass, which fuses on application of heat and sublimes at a red heat in 
 white, six-sided leaflets. Its vapor density at 1300-1500° corresponds to 
 the formula FeCl.^, but it ajqiears that at lower temperatures it is also 
 possible for the molecules Fc2Cl^ to exist. 
 
 It forms double salts with the alkaline chlorides, e. g. : 
 
 YeC\.2KC\-{-2\\^0. 
 
 Ferrous Iodide, Fel2, is obtained by warming iron with iodine and 
 water. It crystallizes with four molecules of water. 
 
 Ferrous Oxide, FeO, is a black powder, resulting from the reduction 
 of ferric oxide by carbon monoxide. When warmed in the air it oxidizes 
 readily. Ferrous Hydroxide, Fe(OH)2, is thrown out of ferrous solu- 
 tions by the alkalies, as a greenish-white precipitate. Exposed to the 
 air, it oxidizes, becoming reddish-brown. It is somewhat soluble in 
 water, and has an alkaline reaction. 
 
 Ferrous Sulphate, FeSO^, crystallizes with seven molecules of water 
 in large, greenish, monoclinic prisms, and is generally called 
 The crystals effloresce somewhat in dry air. They oxidize in moist air, 
 and become coated with a brown layer of basic ferric sulphate. At 100° 
 they lose six molecules of water, and change to a white powder. The 
 last molecule of water escapes at 300°. Therefore, ferrous sulphate 
 behaves just like the sulphates of the metals of the magnesium group 
 (p. 315). Like them, it unites with alkaline sulphates to double sulphates, 
 which contain six molecules of water, e. g., FeSO^. K2SO4 -f- 6H2O. 
 These are more stable than ferrous sulphate, do not effloresce, and oxidize 
 very slowly in the air. 
 
 The salt FeSO^. (NHJ2SO4 -f- 6H2O — Mohr’ s salt — is particularly char- 
 acterized by its stability in the air, and is therefore employed to determine 
 the strength of the potassium permanganate used in titrations. 
 
 Ferrous sulphate is obtained by dissolving iron in dilute sulphuric acid. 
 Commercially, it may also be made from pyrite (FeS2), which by careful 
 roasting loses one atom of sulphur, and is converted into ferrous sul- 
 ])hide (FeS), which, in the presence of water, absorbs oxygen from the 
 air, and is converted into sulphate, which may then be extracted by water. 
 
 Iron vitriol has an extended practical application; among other uses, 
 it is emj)loyed in the ])reparation of ink, and in dyeing. 
 
 When heated it decomposes according to the following equation : 
 
 2FeS(), = Fe2^;t I L 
 
 On this is based the production of fuming Nordhausen sulphuric acid 
 (p. 194), and of colcothar. 
 
FERRIC COMPOUNDS. 
 
 399 
 
 Ferrous Carbonate, FeCOg, exists in nature as siderite, crystallized 
 in yellow-colored rhombohedra, isomorphous with calcite and smithsonite. 
 Sodium carbonate added to ferrous solutions precipitates a white volum- 
 inous carbonate, which rapidly oxidizes in the air to ferric hydroxide. 
 Ferrous carbonate is somewhat soluble in water containing carbon dioxide, 
 hence present in many natural waters. 
 
 Ferrous Phosphate, FegCPOJg + SH^O, occurs crystallized in bluish 
 monoclinic prisms as Vivianite. Precipitated by sodium phosphate from 
 ferrous solutions, it is a white amorphous powder, which oxidizes in the 
 air. 
 
 Ferrous Sulphide, FeS, is a dark-gray, metallic mass, obtained by 
 fusing together iron and sulphur. It is made use of in laboratories for 
 the preparation of hydrogen sulphide. If an intimate mixture of iron 
 filings and sulphur be moistened with water, the union will occur even at 
 ordinary temperatures. Black ferrous sulphide is precipitated from ferrous 
 solutions by alkaline sulphides. When the moist sulphide is exposed to 
 the air it oxidizes to ferrous sulphate. The alkaline sulphides also pre- 
 cipitate ferrous sulphide from ferric salts, but the latter first suffer reduc- 
 tion : 
 
 2Fea3 -f (NHd,S = 2FeCl2 + 2NHW'l -f S, 
 
 and 
 
 FeCb + (NHJgS = FeS + 2NH,C1. 
 
 FERRIC COMPOUNDS. 
 
 Ferric Sesquioxide of irofi^ exists in nature, in com- 
 
 pact, massive form, as red hematite, and as specular iron, in dark-gray 
 metallic rhombohedra. It may be prepared by heating the iron oxygen 
 compounds in the air, and is obtained on a large scale by the ignition 
 of green vitriol. It is then a dark-red powder {colcothar or caput 
 7 no 7 'tuu 77 i) used as a paint and for polishing glass. 
 
 Ferric Hydroxide, Fe(OH)3, is precipitated by alkalies from ferric 
 solutions as a voluminous, reddish-brown mass. On boiling, it becomes 
 more compact, gives up water, and is converted into the hydrate, 
 Fe,^0(0H)^. Many iron ores, like xanthosiderite, Fe20(0H)4, gdthite, 
 Fe202(0H)2 (isomorphous with diaspore, p. 350), and limonite, 
 Fe^Og^OH)^, are analogous compounds. 
 
 Freshly precipitated ferric hydroxide is soluble in a solution of ferric 
 chloride or acetate. When such a solution is subjected to dialysis, the 
 iron salt diffuses, and there remains a pure aqueous solution of ferric 
 hydroxide. All of the latter is precipitated as a jelly from such a solu- 
 tion upon the addition of a little alkali or acid. 
 
 Ferrous-ferric Oxide, FCgO^ = Fe0.Fe20g, occurs in nature crys- 
 tallized in black regular octahedra — magnetite. It is abundant in 
 Sweden and Norway, and in the Urals. It may be obtained artificially 
 by conducting steam over ignited iron (p. 397). Magnetite constitutes 
 the natural loadstone. 
 
400 
 
 INORGANIC CHEMISTRY. 
 
 Ferric hydroxide, like otlier sesciiiioxides, is a feelde liase, and does 
 not yield salts with weak acids, like carbonic or siilpliurous (p. 269). 
 
 Ferric salts arise by the solution of ferric oxide in acids, or by the 
 oxidation of ferrous salts in the ])resence of free acids (best by chloric 
 or nitric acids) : 
 
 2 FeS(), + 11,80, + O = Fe,(S0,)3 d H/). 
 
 Most of the ferric salts have a yellow-brown color, and are converted 
 by reduction into ferrous salts ; 
 
 2FeCl3 t- 11,8 = 2FeCl2 -f 21101 -f 8. 
 
 Ferric Chloride, FeCl,,. It is obtained in aqueous solution by con- 
 ducting chlorine into a solution of ferrous chloride: 
 
 2FeCl2 -f Cl, 2FeCl3. 
 
 The hydrate, 2FeCl3 -j- 3H2O, remains upon evaporation. It is a yel- 
 low crystalline mass, readily soluble in water, alcohol, and ether. When 
 heated, it is partly decomposed, hydrogen escapes, and a mixture of 
 chloride and oxide remains. 
 
 Anhydrous ferric chloride is produced by heating iron in a current of 
 chlorine gas; it sublimes in brownish-green, metallic, shining, six-sided 
 prisms and scales, which deliquesce in the air. Ferric chloride boils at 
 280-285°. Its vai)or density between 320-440° closely approximates the 
 formula Fe2Clg; with rising temperature it diminishes gradually, and from 
 750-1050° corresponds to the simple formula FeCl,. The vapor, however, 
 ])robably does not consist of these molecules, because at 440° a decompo- 
 sition into ferrous chloride, FeCl,, and chlorine commences (Friedel and 
 Crafts, Ber. (1888) 21, Ref. 580). Recently the molecular formula FeCl, 
 has been found by determining the rise in the boiling point of ethereal 
 and alcoholic solutions of ferric chloride. 
 
 Ferric chloride solutions can take up large quantities of ferric hydrate. 
 In the resulting, dark-colored solutions ferric hydroxychlorides are 
 present : 
 
 nFeClg -)- mFe203 -j- XH2O ^Liquor fcrri oxychlorati). 
 
 Ferric Sulphate, Fe2(SO,)3, is obtained by dissolving the oxide in 
 sulphuric acid. When its solution is evaporated, it remains as a white 
 mass, which gradually dissolves in water, with a reddish-brown color. It 
 forms alums (p. 352) with alkaline sulphates, e. g. : 
 
 Fe2(SO,)3. K, 80 , + 24H2O. 
 
 Potassium iron alum. 
 
 Ferric Phosphate, FePO,^ is a yellowish-white precipitate, thrown 
 out of ferric solutions by sodium phos])hate. It is insoluble in water and 
 acetic acid. 
 
 Iron Disulphide, FeS,, occurs in nature as iron pyrites, crystallized 
 in yellow, metallic, shining, (^ctahedra or pentagonal dodecahedra. It is 
 employed in the manufacture of sulphuric acid and green vitriol. The 
 
CYANOGEN DERIVATIVES OF IRON. 
 
 401 
 
 artificial sulphide can be prepared in many ways. It leaves ferrous sul- 
 phide when it is strongly heated in hydrogen. 
 
 COMPOUNDS OF FERRIC ACID. 
 
 On fusing iron filings with niter, or by conducting chlorine into potas- 
 sium hydroxide, in which ferric hydroxide is suspended, potassium fer- 
 rate, K2Fe04, is produced, and crystallizes from the alkaline solution in 
 dark-red prisms. This salt is isomorphous with potassium chromate and 
 sulphate. It dissolves quite easily in water ; but the dark-red liquid soon 
 decomposes with separation of ferric hydroxide and evolution of oxygen. 
 The free acid is not known, as it immediately breaks down when liberated 
 from its salts. 
 
 CYANOGEN DERIVATIVES OF IRON. 
 
 When potassium cyanide is added to aqueous solutions of the ferrous or 
 ferric salts, the cyanides, Fe(CN)2 and Fe(CN)3, are thrown down as yel- 
 lowish precipitates, which decompose rapidly in the air. They dissolve 
 in an excess of potassium cyanide to form the double cyanides, Fe(CN)2.- 
 4KCN and Fe(CN)3. 3KCN. When acids are added to strong solutions 
 of these salts the hydrogen compounds, H4Fe(CN)g, and H3Fe(CN)g sepa- 
 rate. Like the halogen hydrides they are acids and form salts by exchang- 
 ing their hydrogen for metals. The iron and the cyanogen group in these 
 salts and in the free acids cannot be detected by the usual reagents {e. g., 
 the iron is not precipitated by the alkalies). It is supposed that com- 
 pound groups of peculiar structure are present in these double cyanides, 
 and that they conduct themselves like the halogens. The group, FeCyg, 
 in the i^rrous compounds is called ferrocyanogen, that of FeCyg in the 
 ferr/V, ferricyanogen. (Nothing is known in regard to the molecular 
 magnitude of these compounds; Cy = CN). The ferro- behave toward 
 the ferri-compounds the same as the ferrous toward the ferric salts ; oxi- 
 dizing agents convert the former into the latter, and reducing agents 
 transform the latter into the former : 
 
 and 
 
 K,Fe(CN)g + Cl = K 3 Fe(CN)g + KCl 
 K3Fe(CN)g -f KOH + H = K4Fe(CN)e -f H^O. 
 
 Cobalt, manganese, chromium and the platinum metals afford similar 
 cyanides (pp. 377, 407). 
 
 Metallic acid radicals such as are known to exist in the chromates, manganates, per- 
 manganates and ferrates are assumed to be present in the ferro- and ferricyanides. Such 
 radicals, with a basic character, have been encountered: the radicals (WgCh)!! and 
 (MogCb)!! which j)lay the role of a bivalent metal towards the halogens in the com- 
 pounds — chlor-tungsten hydroxide and chlor-molybdenum hydroxide (p. 382). 
 
 All of these radicals have this in common — the metal present in the radical is no 
 34 
 
402 
 
 in()R(;anic chemistry. 
 
 longer (letccled l)y the reagents to wliieli it responds wlien existing as metal in its salts, 
 'This it will oidy do after the radical has been destroyed. If the radieals eonlain halogens 
 they too will behave like the metals. 'This failure to resi)ond to the usual reagents is an 
 indieation that the metal or halogen belongs to a compound radical and docs ncjt act ns an 
 independent part in the molecule. As mentioned on p. 270, salts, bases and acids are 
 regarded as compounds of two members when they are resolved in acjueous solutiem into 
 two kinds of ions. Numerous transpositions and particularly those of analytical value 
 occur between ions, hence they are briefly termed “ion reactions.” I'.very ion, be it 
 simple or compound, is characterized by reactions which belong to it alone and by which 
 it is detected. Hence in general we cannot speak of the reactions of the elements, e. jif., 
 of iron, when search is being made for its presence, Init we must rather consider the con- 
 dition, the combination form, the nature of tlie ion, for whicli the reaction has value. 
 There are no general reagents or reactions for iron, but only those for metallic iron, or 
 for ferrous, ferric, and ferrocyanide comi)ounds, etc. ; no general reaction for chlorine, 
 but only such as answer for free chlorine, hydrogen chloride, cldoric acid, percldoric 
 acid, etc. — reactions of the ions according to the theory of electrolytic dissociation, d'his 
 is expressed in the following formulas which are divided by lines into their ions : 
 
 FeiCk, K^KFeCyg) II Cl HKClOj) Mnl(SO,) K^^MnOJ 
 MoiCl, (Mo.p,)|Cl2 Cr2!(SOj3 K^KCrOJ. 
 
 Potassium Ferrocyanide, Yellow prussiate of potash, K^Fe(CN)g, 
 is produced by the action of potassium cyanide upon iron compounds, or 
 upon free iron (in which case the oxygen of the air or water takes part). 
 It is prepared commercially by igniting carbonized nitrogenous animal 
 matter (blood, horns, hoofs, leather offal, etc.) with potashes and iron. 
 
 In this operation, the carbon and nitrogen of the organic matter com- 
 bine with the potassium of the potashes to form potassium cyanide, while 
 the sulphur present forms iron sulphide with the iron. (By means of 
 alcohol, potassium cyanide can be extracted from the fusion.) Upon 
 treating the fusion with water, the potassium cyanide and iron sulphide 
 react upon one another, and potassium ferrocyanide results and is puri- 
 fied by crystallization : 
 
 FeS + 6KCN = K,Fe(CN)6 + K^S. 
 
 In Germany, at present, yellow prussiate of potash is manufactured 
 exclusively from the material used in purifying gas (p. 105), which con- 
 tains the greater part of the cyanogen of the crude gas in the form of 
 Prussian blue and ammonium sulphocyanide. The first is transformed 
 by lime into ferric oxide and calcium ferrocyanide, which in aqueous 
 solution is changed by potassium chloride to calcium chloride and very 
 sparingly soluble potassium calcium ferrocyanide : 
 
 Ca,Fe(CN)e + 2KCI =3 CaK.,Fe(CN )6 + CaCh. 
 
 Potashes are finally used to convert the insoluble salt into calcium carbon- 
 ate and potassium ferrocyanide. 
 
 It crystallizes from water in large, yellow, monoclinic prisms, having 
 three molecules of water, and soluble in 3-4 parts of water. The crystals 
 lose all their water at 100°, and are converted into a white })owder. At a 
 red heat the ferrocyanide breaks down into potassium cyanide, nitrogen, 
 and iron carbide (FcG^). When the salt is warmed with dilute sulphuric 
 
CYANOGEN DERIVATIVES OF IRON. 
 
 403 
 
 acid, half of the cyanogen escapes as hydrogen cyanide ; concentrated 
 sulphuric acid decomposes it, according to the following equation : 
 
 K,Fe(CN)6 + 6H2SO, + 6H,0 = FeSO, + aK^SO^ + 3(NH4),SO, + 6CO. 
 
 When strong hydrochloric acid is added to a concentrated solution of 
 potassium ferrocyanide hydrogefi ferrocyanide, H^FeCyg, separates as a 
 white crystalline powder, which soon turns blue in the air. It is an acid. 
 Its salts with the alkali and alkaline earth metals are very soluble in water. 
 The sodium salt crystallizes with difficulty. The salts of the heavy metals 
 are insoluble in water, and are obtained by double decomposition. When 
 potassium ferrocyanide is added to the solution of a ferric salt a dark- 
 blue cyanide 
 
 Fe^(Cy)ig = (Fe2)2(FeCyg)3, 
 
 called Prussian blue, is precipitated : 
 
 3K,FeCyg + 4Fe2Clg = (Fe2)2(FeCyg)3 + 12KCI. 
 
 This color was discovered accidentally by Diesbach, of Berlin, in 1704. 
 
 It is the ferric salt of hydroferrocyanic acid; and if potassium or 
 sodium hydroxide is poured over it, it is converted into ferrocyanide of 
 potassium and ferric hydroxide; 
 
 VI VI 
 
 (Fe2)2(FeCyg)3 + 12KOH = 3K^FeCyg -f- 4Fe(OH)3. 
 
 Potassium ferrocyanide produces a reddish-brown precipitate of cop- 
 per ferrocyanide, CiqFeCyg, in copper solutions ; and in ferrous solutions a 
 white precipitate, which on exposure to the air rapidly becomes blue in 
 color. 
 
 Oxidizing agents convert the ferro- into potassium ferricyanide, 
 K^FeCyg, red prussiate of potash. This conversion is most conveniently 
 effected by conducting chlorine into the solution of the yellow prussiate : 
 
 aK^FeCyg -j- Clg = 2K3FeCyg -j- 2KCI. 
 
 The quadrivalent ferrocyanogen group, FeCyg, is then changed to the 
 trivalent ferricyanogen group, FeCyg (p. 401). Cy = CN. 
 
 The red prussiate crystallizes from water in red rhombic prisms. The 
 free hydroferricyanic acid, H3Fe(CN)g, is precipitated upon the addition 
 of concentrated hydrochloric acid. It is rather unstable. 
 
 With ferrous solutions potassium ferricyanide yields a dark-blue pre- 
 cipitate, FegCyij = FegFcaCyii, very similar to Prussian blue, and called 
 Turnbuir s blue : 
 
 2K3FeCyg -f 3FeSO^ = FegCy,^ + 3^280,. 
 
 So far as its formation is concerned, 'rurnbull’s blue is the ferrous salt of hydroferri- 
 III II 
 
 cyanic acid, [Fe(CN)g] Fcg. Although liot alkalies decompose it into potassium ferro- 
 cyanide and ferrous-ferric hydroxide, this is most probably due to the fact that the potassium 
 
404 
 
 INORGANIC CIIKMIS'I’RY. 
 
 ferricyanide and ferrous liydroxidc, which arc i)roduccd at first, arc traiisjjoscd by the 
 alkali into potassium ferrocyanide and fcrrous-fcrric hydrate : 
 
 2Fe(CN),K3 -I- 2KOII 4 3l''e{OII), ^ 2K,Fc(CN), + 2Fc(OII)3 t Fc(()II),. 
 
 Very probably in the i)recipitation potassium ferricyanide and the ferrous salt transpose 
 into potassium ferrocyanide and a ferrous-ferric salt ; then 'I'urnbuirs blue would have to 
 
 III II IV 
 
 be regarded as the ferric-ferrous salt of hydroferrocyanic acid — J‘'e. 4 ‘'e[Fe((:N),.].^. 'I'urn- 
 bull’s blue is the principal constituent of commercial J’russian blue, which also contains, 
 because of its preparation from impure materials, ferriferrocyanide, Fe7(CN),y. '1 he 
 latter largely comprises Parisian blue. 
 
 Potassium ferricyanide does not cause precipitation in ferric solutions, 
 pjy these reactions, ferric salts may be readily distinguished from the 
 ferrous. Potassium sulphocyanide (KCNS) produces a dark- red colora- 
 tion in ferric solutions, while it leaves the ferrous unaltered. 
 
 Iron, like nickel (]). 233), combines with carbon monoxide to a vola- 
 tile, gaseous body — Iron Carbonyl [Mond and Quincke, Per. (1891) 
 24, 2248; Perthelot, Compt. rend. (1891) 112, 1343], which, from its 
 composition, is probably Iron Tetracarbonyl, Fe(CO)^. This substance 
 is produced by conducting carbon monoxide over very finely divided 
 iron at 40-80°, or under a pressure of eight atmospheres [Roscoe and 
 Scudder, Per. (1891) 24, 3843]. It is decomposed with the formation 
 of an iron mirror when it is passed through a glass tube heated to 200- 
 
 350°- 
 
 Iron pentacarbonyl, Fe(C0)5, hepiacarbonyl, Fe.2(CO)y, have been prepared. 
 
 The first is a yellow liquid, boiling at 103°, from which the second separates, on expos- 
 ure to the light, in golden crystals. Carbon monoxide is evolved at the same time. 
 
 2 . NICKEL. 
 
 Ni = 58.7.* 
 
 Nickel exists in native condition in meteorites; its most important 
 ores are niccolite, NiAs, and gersdorffite, NiSa . NiAsj (constituted like 
 cobaltite). The arsenical ores are now of little importance in the nickel 
 industry. The chief sources of the metal are at present nickeliferous 
 ])hyrrotite and the nickel silicates (Canada, Norway and New Caledonia). 
 Clarnierite — a New Caledonian mineral — a silicate, containing also iron, 
 calcium and magnesium, may be especially noticed in this connection. 
 Nickel is almost always accompanied in its ores by cobalt, and vice versa, 
 cobalt usually by nickel. The isolation of the latter from its ores and 
 from speiss-cobalt (p. 406) is very complicated. Nickel usually appears 
 in cfimmerce in cubical forms, which in addition to the chief ingredient 
 alvvays contain some copi)er, bismuth, and other metals. Chemically 
 ))ure nickel is procured by igniting the oxalate in a current of hydrogen. 
 Nickel is almost silver-white in color and is very lustrous, and very 
 
 * Sc(; 'I'll. \V. Richards and A. S. ('ushman, Z. f. anorg. Chem. 20 (1899), 352. 
 
NICKEL. 
 
 405 
 
 tenacious. Its specific gravity varies from 8.8 to 9.1. It fuses at a some- 
 what lower temperature than iron, and like it is attracted by the magnet. 
 It is not altered in the air; it dissolves with difficulty in hydrochloric and 
 sulphuric acids, but readily in nitric acid. 
 
 Its derivatives are almost exclusively of the ous form, NiX.^; nickelic 
 oxide behaves like a peroxide, and does not form corresponding salts. 
 
 Nickelous Hydroxide, Ni(OH).^, is a bright green precipitate pro- 
 duced by alkalies in nickelous solutions. It dissolves in ammonium 
 hydroxide, with a blue color. When heated it passes into gray nickelous 
 oxide, NiO. 
 
 Nickelous Chloride, NiCl^ -f bH.^O, consists of green, monoclinic 
 prisms. When heated they lose water and become yellow. 
 
 Nickelous Cyanide, Ni(CN)2, is preci})itated by potassium cyanide 
 as a green-colored mass from nickel solutions. It is soluble in excess of 
 the precipitant. The double cyanide, NiCyj. 2KCy -|- H2O, crystallizes 
 from the solution. This salt is readily decomposed by acids. Cyanogen 
 compounds of nickel, constituted like those of iron and cobalt, are not 
 known. 
 
 Nickelous Sulphate, NiSO^ -j- 7H2O, ■ appears in green, rhombic 
 prisms, isomorphous with the sulphates of the magnesium group, and 
 forms analogous double salts. 
 
 Nickelous Sulphide, NiS, is precipitated, black in color, by alkaline 
 sulphides from nickel solutions. 
 
 Nickelic Oxide, Ni203 and Hydroxide, Ni2(OH)g, are perfectly 
 similar to the corresponding cobalt salts ; when warmed with hydrochloric 
 acid they liberate chlorine. 
 
 Nickel Tetracarbonyl, Ni(CO)4, see pp. 234, 404. 
 
 Nickel is used for certain alloys Argentan consists, ordinarily, of 50 
 per cent, of copper, 25 per cent, of nickel and 25 per cent, of zinc. The 
 white color, the hardness and the power of receiving a higher polish in- 
 crease with the increase in the nickel content. The iron-nickel alloys 
 have great technical importance (nickel steel). The German nickel 
 coins in which we have evidence of the great coloring power of nickel 
 consist of 75 per cent, of copper and 25 per cent, of nickel. Nickel 
 alloys are used to make electrical resistances; their conductivity is slight 
 and little depends on this at the ordinary temperature {niaiiganin : 84 
 per cent, of copper, 4 per cent, of nickel, 12 per cent, of manganese; 
 consfaiiian : 60 per cent, of copper, 40 per cent, of nickel, etc.). At 
 present, cast-iron ware is coated with a layer of nickel to prevent it from 
 rusting and to imj)art to it a beautiful white surface. This is accomplished 
 by electroplating, or by boiling the iron ware in a solution of zinc 
 chloride and nickel sulphate. 
 
 In the electrolytic method a solution of the double sulphate of nickel 
 and ammonium is emjdoyed ; the positive electrode consists of a pure 
 nickel plate, while the object to be coated forms the negative electrode. 
 
4o6 
 
 INORGANIC CIIEMISIRY. 
 
 3. COBALT. 
 
 Co = 59 * 
 
 Cobalt occurs in nature as sinallile (CoAs^) and cobaltite (CoAs.^. CoSj). 
 The metal is obtained by the ignition of cobaltous oxide with carbon, or 
 in a current of hydrogen. It has a reddish-white color and strong luster, 
 is very tenacious, and fuses with difficulty. Its specific gravity is 8.5. It 
 is attracted by magnets, but to a less degree than iron. It is not altered 
 by the air or water. It is only slightly attacked by hydrochloric and 
 sulphuric acids ; nitric acid dissolves it readily, forming cobaltous nitrate. 
 
 The predominating compounds have the form CoX.^, and are called 
 cobaltous. They are very stable, and generally isomorphous with the 
 corresponding ferrous salts, 'bhe hydrous cobaltous compounds have a 
 reddish color, the anhydrous are blue. 
 
 COBALTOUS COMPOUNDS. 
 
 Cobaltous Chloride, C0CI2, is obtained by the solution of cobaltous 
 oxide in hydrochloric acid, and crystallizes with six molecules of water 
 in red monoclinic prisms. When heated, it loses water, and becomes 
 anhydrous and blue in color. Characters made with this solution upon 
 paper are almost invisible, but when warmed they become distinct and 
 blue (sympathetic ink). 
 
 Cobaltous Hydroxide, Co(OH) 2, is a reddish precipitate produced 
 by the alkalies in hot, cobaltous solutions. When exposed to the air, it 
 is colored brown by oxidation. Basic salts are precipitated from cold 
 solutions; these dissolve with a blue color in an excess of concentrated 
 alkalies. When heated out of air contact, the hydroxide passes into 
 green cobaltous oxide, CoO. 
 
 Cobaltous Sulphate, CoSO^ -j- yH^O, crystallizes in dark-red 
 monoclinic prisms; the hydrated suljffiate, CoSO^ -j- 6H2O, separates 
 from hot solutions. It is isomorphous with ferrous sulphate, and yields 
 double salts with alkaline sulphates (p. 398). 
 
 Cobaltous Nitrate, Co(N03)2 + 6H2O, forms red deliquescent 
 prisms. 
 
 Cobaltous Sulphide, CoS, is a black precipitate, produced in neu- 
 tral cobalt solutions by alkaline sulphides. It is insoluble in dilute 
 acids. 
 
 Cobalt Silicates. — When glass is fused with a cobalt compound it is 
 colored a dark blue, and when reduced to a powder is used as a pigment 
 under the name of smalt. 
 
 Smalt is j)repared comitiercially by fusing cobalt ores with potashes and quartz. The 
 cobalt forms a silicate (smalt) with the silica and potassium, while the other metals, such 
 as bismuth, arsenic, and csi^ecially nickel, accompanying it in its ores, are thrown out as 
 a metallic mass. 'I'his is called si)eiss-cobalt and serves for the preparation of nickel. 
 
 * See (d. Winkler, Z. f. anorg. Ch. 17 (1S98), 236. 
 
COBALTIC COMPOUNDS. 
 
 407 
 
 On igniting cobalt oxide, CO2O3, with alumina, a dark-blue mass is 
 produced — cobalt uli 7 -a 7 tia 7 -ine or The 7 ia 7 'd^ s blue. When zinc oxide and 
 cobalt oxide are ignited a green color — greeTt cmTtabar or Ri 7 i 7 naft 7 i' s 
 gree 7 i — is obtained. 
 
 COBALTIC COMPOUNDS. 
 
 Cobaltic Oxide, CO2O3, is left as a black powder on the ignition of 
 cobaltous nitrate. It becomes cobaltous-cobaltic oxide, COgO^, at a red 
 heat, and cobaltous oxide at a white heat. The hydroxide, Co2(OH)j., 
 separates as a dark-brown powder, if chlorine be passed through an alka- 
 line solution containing a cobaltous salt. 
 
 A cobaltous salt is produced and oxygen set free, when sulphuric acid 
 acts upon the oxide or the hydroxide. Chlorine is generated when it is 
 heated with hydrochloric acid : 
 
 C02O3 -f 6HC1 = 2C0CI2 + 3H2O + CI2. 
 
 The cobaltic hydroxide dissolves in dilute, cold hydrochloric acid, 
 with scarcely any liberation of chlorine ; the solution probably contains 
 cobaltic chloride, C0CI3, which decomposes into cobaltous chloride and 
 chlorine on evaporation. 
 
 Cobaltous-cobaltic Oxide, 0030^=00203.000, corresponding 
 to magnetite, FegO^, is formed upon the ignition of the oxygen cobalt 
 derivatives in the air, and is a black powder. 
 
 It is noteworthy that cobalt is capable of yielding complex derivatives 
 in which it appears to be in union with the groups NO2, ON or NH3, 
 forming peculiar radicals. One of the most interesting of these is 
 potassium cobaltic nitrite (Fischer’s salt). 
 
 When potassium nitrite, KNOg, is added to a cobaltous solution acidi- 
 fied with acetic acid, nitric oxide is set free, and in course of time 
 
 Co(N 02)3.3KN02+ nH20, 
 
 a double salt, separates as a yellow crystalline powder. This should be 
 viewed as the potassium salt of a hydronitroso-cobaltic acid : 
 
 H3Co(N02V 
 
 This reaction is very characteristic for cobalt, and serves to separate it 
 from nickel. 
 
 Ammonia-cobalt Compounds. — Numerous cobaltamines are known. On adding 
 ammonium hydroxide to a cobaltous chloride solution, the precipitate first formed dissolves 
 in the excess of the reagent, and when this liquid is permitted to stand exposed to the air, 
 the color, which is brown at first, gradually passes into red. On adding concentrated 
 hydrochloric acid to this solution, a brick-red, crystalline powder, of the composition 
 
 C02CI3.10NH3 -f 2II2O, 
 
4o8 
 
 INORGANIC CHEMISTRY, 
 
 called roseocobaltic chloride^ is precipitated. If, however, the red solution be boiled with 
 hydrochloric acid, a red j)owder, piirpureocobaltic chloride^ 
 
 COjClg. lONIIg, 
 
 separates out. If the ammoniacal red solution contain much ammonium chloride, hydro- 
 chloric acid will precipitate a yellowish-brown compound — liUeocobaltic chloride, 
 
 Co2Cl6.’i2XIl3. 
 
 These derivatives are supposed to contain compound basic radicals in which cobalt, halo- 
 gens and ammonia are all })resent. 
 
 'I'he other salts of cobalt, such as the sulphate and nitrate, yield similar compounds, : 
 
 Co 3 (N 03 )e.ioNIl 3 , 
 
 roseocobaltic nitrate. 
 
 Cyanogen Cobalt Compounds. — In solutions of cobaltous salts, potassium cyanide 
 produces a bright brown precipitate of cobaltocyanide, 
 
 Co(CN)3 , 
 
 soluble in an excess of the reagent. The solution absorbs oxygen from the air, and is 
 converted into potassium cobalticyanide, 
 
 K3Co(CN)„ 
 
 corresponding to potassium ferricyanide. When the solution is evaporated the cobalti- 
 cyanide crystallizes in colorless rhombic prisms, very soluble in water. Sulphuric acid 
 precipitates hydrogen cobalticyanide, 
 
 Il3Co(CN)„ 
 
 from the concentrated solution. This acid crystallizes in needles. 
 
 GROUP OF THE PLATINUM METALS. 
 
 Besides platinum, this group comprises palladium, rhodium, ruthenium, 
 osmium, and iridium — the constant companions of the first in its ores. 
 On p. 392 we observed that these metals are divided into two groups: 
 {hegrou/> of light platmiim tnetals, and group of heavy platinum metals. 
 The latter have higher atomic weights and specific gravities: 
 
 Ru, 101.7 Rh, 103.0 Pd, 106 Os, 191 Ir, 193.0 Pt, 194.8 
 
 Sp.gr. “ 12.3 “ 12. 1 “ II. 5 “ 22.48 “ 22.4 “ 21.5 
 
 At. vol. “ 8.3 “ 8.6 “ 9.1 “ 8.4 “ 8.6 “ 9.0. 
 
 The relations of the metals of these two groups to each other are per- 
 
 fectly similar to those of the iron group. Osmium and ruthenium, like 
 iron, have a gray color, fuse with difficulty, and are readily oxidized in 
 the air. Palladium and ])latinum, on the other hand, have an almost 
 silver-white color like nickel, are more easily fusible, and are not oxidized 
 by oxygen. In chemical respects osmium and ruthenium, like iron, also 
 show a metalloidal nature, inasmuch as their highest oxygen compounds 
 
GROUP OF THE PLATINUM METALS. 409 
 
 form acids. Their derivatives show a complete parallelism with those of 
 iron : 
 
 II 
 
 Ill 
 
 IV 
 
 VI 
 
 OsO 
 
 OSjOg 
 
 OSO2 
 
 (OSO3) 
 
 Osmous 
 
 Osmic 
 
 Osmium 
 
 Osmium 
 
 oxide. 
 
 oxide. 
 
 dioxide. 
 
 trioxide. 
 
 RuO 
 
 RUjOj 
 
 Ruthenic 
 
 RUO2 
 
 (RUO3) 
 
 Ruthenious 
 
 Ruthenium 
 
 Ruthenium 
 
 oxide. 
 
 oxide. 
 
 dioxide. 
 
 trioxide. 
 
 The acid oxides OsOg and RuO, are unknown, but the corresponding 
 acids, H20 sO^ (osmic acid) and H2RuO^ (ruthenic acid), and their salts 
 have been obtained. Besides the derivatives already mentioned we find 
 that osmium and ruthenium are capable of still higher oxidation, yielding 
 OsO^, per-osmic anhydride, and RuO^, per-ruthenic anhydride — which is 
 not the case with iron ; in these compounds the metals appear to be octads, 
 yet these oxides do not form corresponding acids or salts. 
 
 Rhodium and iridium, like cobalt, do not yield acid-like derivatives. 
 Their salts correspond to the forms : 
 
 II III IV 
 
 RhO Rh203 Rh02 
 
 Rhodous Rhodic Rhodium 
 
 oxide. oxide. dioxide. 
 
 The rhodic compounds are the more stable derivatives. 
 
 Palladium and platinum, finally, are relatively of more basic nature, as 
 their ous derivatives, PdX2 and PtX2, are proportionally more stable than 
 the ic forms, PdX^ and PtX^. Palladium also forms a lower oxide, palla- 
 dium suboxide, Pd20, in which it approaches silver. 
 
 The platinum metals are found in nature almost exclusively in the so- 
 called platinum ore, which usually occurs in small metallic grains in accu- 
 mulated sands of a few regions (in California, Australia, the island of 
 Sumatra, and especially in the Urals). The platinum ore, like that of 
 gold, is obtained by the elutriation of the platiniferous sand with water, 
 whereby the lighter particles are carried away. Platinum ore usually con- 
 tains 50-80 percent, of platinum, besides palladium (to 2 per cent.), irid- 
 ium (to 7 per cent.), osmium (i^ per cent.), and ruthenium (i)4 pei' 
 cent.), and different other metals, as gold, copper, and iron. 
 
 The separation of the platinum metals is generally executed in the fol- 
 lowing manner: The gold is first removed by dilute aqua regia. Then 
 the ore is treated with concentrated aqua regia, when platinum, palladium, 
 rhodium, ruthenium, and a portion of iridium are dissolved. Metallic 
 grains or leaflets, an alloy of osmium and iridium — platinum residues — 
 remain. Ammonium chloride is then added to the solution and platinum 
 and iridium precipitated as double salts. When the precipitate is ignited, 
 a spongy mass of iridium bearing platinum (platinum sponge) is obtained, 
 which is applied directly in the manufacture of platinum vessels. The 
 filtered solution from the insoluble chlorides contains palladium, rhodium, 
 35 
 
410 
 
 INORGANIC CHEMISTRY. 
 
 and ruthenium, which are thrown down as a metallic powder by iron; 
 their further sejiaration is then effected in various ways. 
 
 Formerly spongy })latinum was employed almost exclusively for the 
 manufacture of })latinum objects ; itwas pressed into moulds, then ignited 
 and hammered out. Now the fusibility of jdatinum in the oxyhydrogen 
 flame is resorted to, and the fused metal run into moulds. 
 
 Platinum containing both iridium and rhodium may be fused directly 
 out of the platinum ore by means of the oxyhydrogen blowpi])e. The 
 greater ])ortion of the osmium and ruthenium is consumed in this opera- 
 tion. The jiresence of iridium and rhodium makes platinum harder and 
 less readily attacked by many reagents. 
 
 RUTHENIUM AND OSMIUM. 
 
 Ru 101.7. Os = 191. 
 
 Ruthenium, discovered by Claus in 1845, has a steel-gray color ; it is very hard, brittle, 
 and difficultly fusible (at about 1800°). When pulverized and ignited in the air it oxidizes 
 to RuO, Ru^ 03 and RuO.^. It is insoluble in acids, and only slowly dissolved by aqua 
 regia. When fused with potassium hydroxide and nitrate, it forms potassium ruthenate, 
 K2l^^i04- 
 
 Ruthenium heated in chlorine gas yields ruthenium dichloride^ RuClj, a black pow- 
 der, insoluble in acids. The trichloride, RuCl,, is obtained by the solution of Ru2(0ing 
 in hydrochloric acid, and is a yellow, crystalline mass, which deliquesces in the air. It 
 yields crystalline double chlorides with potassium and ammonium chlorides, e. g., 
 RuClg.KCl. The tetrachloride, RuCh, is only known in double salts. Ruthenious 
 oxide, RuO, the sesquioxide, RU2O3, and dioxide, RUO2, are black powders, insoluble in 
 acids, and are obtained when ruthenium is roasted in the air. 
 
 The hydroxides, RujlOH)^ and Ru(0H)4, are produced by the action of the alkalies 
 upon the corresponding chlorides, and are very readily soluble in acids. Ruthenic acid, 
 H2RUO4, is not known in a free condition. Its potassium salt, K2RuO^, is formed by 
 fusing the metal with potassium hydroxide and niter or by digesting the tetroxide with 
 dilute alkali at 60° until the evolution of oxygen ceases : 
 
 RuO^ -f- 2KOH = KgRuO^ HgO O. 
 
 It crystallizes in black prisms with a green luster. It contains one molecule of water of 
 crystallization. It absorbs moisture and carbon dioxide with avidity from the air. Its 
 deep orange-red, dilute, aqueous .solution gradually, on exposure to the air, becomes 
 green and a perruthenate results. At the same time a black oxide of the formula RU2O5 
 separates. This decomposition is more rapid when the solution is acted upon with car- 
 bon dioxide, dilute acids, chlorine or bromine. Potassium ruthenate therefore behaves 
 like potassium manganate (p. 390). When chlorine is conducted through a concentrated 
 .solution of ])otassium ruthenate, or when ruthenium is roasted at 1000° in a current of 
 oxygen, ruthenium tetroxide, RuO^, and dioxide are formed. By the chlorine method 
 it volatilizes and may be collected in a well-cooled receiver. It consists of a gold-yellow 
 crystalline mass. It fu.ses at 25.5°, therefore in the hand, to a deep orange-red colored 
 li(|uid. It readily sublimes in large, yellow transparent ciystals with an orange-yellow 
 reflex. At the ordinary pressure it does not boil at 106°, but it decomposes with ex- 
 plosion at 108°. It can be gasified without decomposition under diminished pressure. 
 Its vapors then correspond to the formula RuO^. 'bhe tetroxide dissolves slowly in 
 water with a golden-yellow color, but does not yield a hydrate. Black compounds of 
 varying comi)()sition se[)arate from this solution. Potassium perruthenate, KRuO^, is 
 furni'al along with the more iiolublc julhcnalc on adding the tdtroxidc to caustic potash. 
 
RHODIUM AND IRIDIUM. 
 
 41I 
 
 It crystallizes in black octahedra with metallic luster; they dissolve in water with a deep- 
 green color [see Debray and Joly, Jahresber. der Chemie (1888), 669, 672]. 
 
 Osmium is very much like ruthenium. It is not even fusible in the oxyhydrogen 
 flame ; it only sinters together. According to Violle it fuses at 2500°. Of all the plati- 
 num metals it is the most difficult to fuse, but is the most easily oxidized. Covipact 
 osinhivi is the heaviest metal specifically (specific gravity 22.48) and is insoluble in aqua 
 regia. Reduced to a fine powder it will burn when ignited in the air to OsO^. Nitric 
 acid and aqua regia convert it into the same oxide. The compounds, OsClj and OsO, 
 OsClg and OsgO,, OsOg and OsCl^, are very similar to the corresponding ruthenium deriva- 
 tives. By fusion with potassium hydroxide and niter we get potassium osmate, K.^OsO^, 
 which crystallizes from aqueous solution with two molecules of water in dark-violet octa- 
 hedra. The most stable and a very characteristic derivative of osmium is the tetroxidcy 
 OSO4, which is produced by igniting the metal in the air, or by the action of chlorine on 
 osmium in the presence of water. It crystallizes in large colorless prisms, which fuse 
 below 100° and distil at a somewhat higher temperature. It has a very sharp, piercing 
 odor {bcfjL'^y odor), similar to that of sulphur chloride. It dissolves slowly but copiously 
 in water ; the solution is not acid. Reducing and organic substances precipitate pulveru- 
 lent osmium from it. This is the basis of its application in microscopy. 
 
 OsO^ and RUO4 do not afford corresponding salts. 
 
 RHODIUM AND IRIDIUM. 
 
 Rh = 103.0. 
 
 Ir 193.0. 
 
 These metals are lighter in color and are more easily fusible than ruthenium and 
 osmium. (Iridium fuses at 1950°.) When pure they are not attacked by acids or aqua 
 regia ; but dissolve in the latter when alloyed with platinum. 
 
 Rhodiufn, discovered by Wollaston in 1803, forms three oxides : 
 
 RhO, RhgOg and RhOj, 
 
 of which the second forms salts with acids. RhOj results when rhodium is heated with 
 niter. 
 
 Of the chlorides only RhCl, is known. It results when the metal is heated in chlorine 
 gas. It is a brownish-red mass. It forms readily crystallizing, red-colored (hence its 
 name, from podd£/f, rose-red) double salts with alkaline chlorides. 
 
 Iridium y discovered in 1804 by Tennant, has perfectly analogous derivatives : 
 
 IrO, IrjOg, Ir02, IrClg, IrCl^. 
 
 The sesquichloride, IrjClg, formed by heating iridium in chlorine, is an olive-green, 
 crystalline mass, insoluble in water and acids. It affbr&s double salts with the alkaline 
 chlorides, e.g. : 
 
 KsIrClg - 3H2O, 
 
 which crystallizes from water in green or brown crystals. They are also produced by 
 the action of sulphur dioxide upon the double salts of hydrochloridic acid. 
 
 Iridium Tetrachloride, IrCl^, is produced in the solution of iridium or its oxide in 
 aqua regia, and remains, on evaporation, as a black mass, readily soluble in water (with 
 red color) to hydrochloriridic acid : 
 
 Il^IrClg. 
 
 When alkaline chlorides are added to the .solution double chlorides are precipitated, e.g, : 
 
 (NH,)2lrClg, K^IrClg, 
 
412 
 
 INORGANIC CHEMISTRY. 
 
 isomorphous with the corresponding double chlorides of platinum. When a solution of 
 IrCl^ is boiled with caustic potash, Ir(OII)^ will be precipitated, d'he name iridium 
 alludes to the differently colored compounds yielded by the metal. 
 
 PALLADIUM. 
 
 I’d = io6. 
 
 Palladium, in addition to occurring in platinum ores, is found alloyed 
 with gold (Brazil), and in some selenium ores (Hartz) ; it has a silver- 
 white color, and is somewhat more fusible (at about 1400°) than platinum. 
 When finely divided it dissolves in boiling concentrated hydrochloric, 
 sulphuric, and nitric acids. When ignited in the air it at first becomes 
 dull by oxidation, but at a higher temperature the surface again assumes 
 a metallic appearance. It was discovered in 1803 by Wollaston and 
 named after the planet Pallas, which had been found shortly before. 
 
 Palladium absorbs hydrogen gas (occlusion) to a much greater extent 
 than platinum or silver. Freshly ignited palladium leaf absorbs upwards 
 of 370 volumes of hydrogen at ordinary temperatures, and about 650 
 volumes at 90-100°. A greater absorption may be effected at ordinary 
 temperatures in the following manner: 
 
 Water is decomposed by the electric current, palladium foil being used 
 as a negative electrode. The liberated hydrogen is then taken up by the 
 palladium (to 960 volumes) ; the metal expands ( jlg- its volume), becomes 
 specifically lighter, but retains its metallic appearance entire. Accord- 
 ing to the investigations of Debray, the compound Pd^H2 is produced, 
 which contains dissolved hydrogen, and deports itself similarly to an 
 alloy. Recent investigations show that palladium hydride is a so-called 
 “solid solution” of hydrogen in palladium (pp. 46, 256). Palladium 
 charged with hydrogen usually remains unaltered in the air, and in a 
 vacuum ; it, however, sometimes becomes heated in the air, as the hydro- 
 gen is oxidized to water. The same occurs when palladium hydride is 
 heated to 100° ; in vacuo, all the hydrogen escapes as gas. Palladium 
 hydride is a strong reducing agent, like nascent hydrogen. Ferric salts 
 are reduced to the ferrous state; chlorine and iodine in aqueous solution 
 are converted into hydrochloric and hydriodic acids. 
 
 Palladium black absorbs hydrogen more energetically than the compact 
 variety (at 100° upwards of 980 volumes). This substance is obtained by 
 the reduction or electrolysis of palladic chloride. If palladium sponge 
 be heated in the air until the white metallic color becomes black, in con- 
 sequence of the superficial oxidation, it will absorb hydrogen very ener- 
 getically at ordinary temperatures, and partly oxidize it to water. 
 
 When ])alladium-sheet or sponge is introduced into the flame of a 
 spirit-lamp, it is covered with soot; this is due to the fact that the 
 metal witlidraws the hydrogen of the hydrocarbons, and carbon is set 
 free. 
 
 'I'here are two series of ])alladium compounds: the palladious, PdX2, 
 and palladic, PdX.^. The first are well characterized and are distinguished 
 by their stability. 
 
PLATINUM. 
 
 413 
 
 Palladious Chloride, PdClj, remains as a brown, deliquescent mass, 
 on evaporating the solution of j)alladium in aqua regia. It yields easily 
 soluble crystalline double salts, with alkaline chlorides, e.g., PdCl2. 2KCI 
 = K^PdCl,. 
 
 Palladious Iodide, Pdl2, is precipitated from palladious solutions by 
 potassium iodide as a black mass, insoluble in water. As palladium chlo- 
 ride and bromide are very soluble in water, palladium salts can be used to 
 detect iodine in the presence of hydrobromic or hydrochloric acid. 
 
 Palladious Oxide, PdO, is a black residue left upon careful ignition 
 of the nitrate. It is difficultly soluble in acids. When heated, it loses 
 oxygen, and forms palladium suboxide, Pd20. 
 
 When palladium is dissolved in sulphuric or nitric acids, the corre- 
 sponding salts are produced. 
 
 The sulphate, PdSO^ -f 2H2O, is composed of brown crystals, readily 
 soluble in water. Much of the latter decomposes it. 
 
 Palladio Chloride, PdCl^, Hydrochlorpalladic Acid, H2PdClg, is 
 formed when the metal is dissolved in aqua regia. They decompose, on 
 evaporation, into PdCl2 and CI2. When potassium or ammonium chlo- 
 ride is added to their solutions, red-colored difficultly soluble double 
 chlorides crystallize out ; they are analogous to the corresponding salts 
 of platinum. 
 
 PLATINUM. 
 
 pt = 194.8. 
 
 The separation of platinum from the ore was described on p. 409. 
 (See Myliusand Foerster, Ber. 25 (1893), 665.) The metal has a grayish- 
 white color, and a specific gravity of 21.4. It is very tough and malleable, 
 and may be drawn out into very fine wire and rolled into foil. At a high 
 heat it softens without melting, and may be easily welded. It fuses in 
 the oxy hydrogen flame (at about 1770° — Violle), and is somewhat volatile. 
 On fusion, it absorbs oxygen, which it gives up again on cooling (like 
 silver). At ordinary temperatures, it also condenses hydrogen and 
 oxygen upon its surface; as foil and sponge, but very few volumes ; as 
 platinum black, about 100 times its volume of oxygen and 310 times its 
 volume of hydrogen. Two hundred volumes of the latter form water 
 with the oxygen always present in the sponge (Z. f. anorg. Ch. 10 (1895), 
 178). These gases are fully expelled at a red heat. Platinum sponge is 
 obtained as a fine black powder, if reducing substances, like zinc, be added 
 to solutions of platinic chloride or upon boiling the solution with sugar 
 and sodium carbonate; it remains on the ignition of PtCl^. 2NH^C1. The 
 production of various reactions is due to the power of platinum to con- 
 dense oxygen ; thus hydrogen will inflame in the air, if it be conducted 
 upon platinum sponge ; sulphur dioxide combines with oxygen at 100° 
 to form the trioxide. At a red heat platinum permits free passage to 
 hydrogen, while it is not permeable by oxygen and other gases (j). 93). 
 
 Platinum is not attacked by acids; it is only soluble in liquids gener- 
 ating free chlorine, e. g., aqua regia. In consequence of this resistance 
 
414 
 
 INOROANIC CHEMISTRY. 
 
 to acids, and its nnalterability upon ignition, tliis metal answers as an 
 imdecomposable material for the production of chemical crucibles, dishes, 
 wire, etc. The usual presence of iridium in ordinary platinum increases 
 its durability. 
 
 The alkaline hydroxides, sulphides, and cyanides attack it strongly at 
 a red heat. It forms readily fusible alloys with ])hosphorus, arsenic, and 
 many heavy metals, especially lead, and many heavy metals are reduced 
 from their salts by platinum. Therefore such substances must not be 
 ignited in platinum crucibles, etc. 
 
 Platinum, like palladium, forms platifious, PtX.^, and platinic, PtX^, 
 derivatives ; in the first it is more basic, in the latter more acidic. 
 
 Platinic Chloride, PtCl^, is obtained by the solution of platinum in 
 aqua regia, and when the liquid is evaporated with an excess of hydro- 
 chloric acid, hydrochlorplati ri ate , 
 
 PtCh.2HCl ^ HjPtCle, 
 
 crystallizes with six molecules of water in brownish-red, deliquescent 
 prisms. It forms characteristic double chlorides, 
 
 PtCh.2KCl, 
 
 with ammonium and potassium chlorides. They are the chlorplatinaies. 
 Formerly they were considered double salts. They are difficultly soluble 
 in water; hence, on mixing the solutions, they immediately separate out 
 as a crystalline yellow powder. Ignition completely decomposes the 
 ammonium salt, leaving spongy platinum. 
 
 Platinum chloride yields similar insoluble salts with the chlorides of 
 rubidium, caesium, and thallium, while that with sodium, 
 
 NajPtCle + 6H.,0, 
 
 is very soluble in water. Platinum may be very readily separated from 
 other metals by recrystallizing the sodium salt from hot water rendered 
 alkaline by soda (Finkener). 
 
 On adding sodium hydroxide to platinic chloride and then supersatu- 
 rating with acetic acid, there separates a reddish-brown precipitate of 
 platinic hydroxide, Pt(OH)^. This dissolves readily in acids (excepting 
 acetic), with formation of salts. The oxygen salts, as Pt(S 04 ") 2 , are very 
 unstable. The hydroxide has also an acidic character (^platinic acid), 
 and dissolves in alkalies, yielding salts with them. These, also, result 
 on fusing platinum with potassium and sodium hydroxide. The barium 
 salt, 
 
 is ])recipitated from platinic chloride, by barium hydroxide, as a yellow, 
 crystalline compound. The acidic nature of its hydroxide allies platinum 
 to gold. If hydrogen sulphide be conducted through platinic solutions, 
 black platinum disulphide, PtSj, is precipitated; it is soluble in alkaline 
 siil|)hides, with formation of sulpho-salts. 
 
 Platinous Chloride, PtCl.^, is a green powder, insoluble in water. 
 
SPECTRUM ANALYSIS. 415 
 
 remaining after heating platinic chloride to 200°. It forms double salts 
 with alkaline chlorides, e. g. : 
 
 PtCl, . 2NaCl. 
 
 When digested with potassium hydroxide it yields the hydroxide, 
 Pt(OH),. 
 
 Cyanogen Compounds. — Like cobalt, platinum forms double cy- 
 anides corresponding to the ferrocyanides. When platinous chloride is 
 dissolved in potassium cyanide platinum potassium cyanide, K.2Pt(CN)^ -|- 
 4H2O, crystallizes on evaporation in large prisms exhibiting magnificent 
 dichroism ; in transmitted light they are yellow and in reflected light 
 blue. This salt must be viewed as the potassium compound of hydro- 
 pi atino- cyanic acid, H2Pt(CN)4. When separated from its salts it crys- 
 tallizes in gold-yellow needles. Its salts with heavy metals are obtained 
 by double decomposition, and all show a beautiful play of colors. The 
 barium salt, barium platino-cyanide, BaPt(CN)4 4H2O, is used to 
 render the Rontgen rays visible. 
 
 Platinum-ammonium Compounds. — There is a whole series of 
 these, which must be viewed as platinum bases and their salts. They are 
 constituted according to the following empirical formulas : 
 
 {a) Platosamines (platinous series) : 
 
 PtR2(NH3)2; PtR2(NH3)3 ; PtR2(NH3b. 
 
 [{h) Platinamines (platinic series) : 
 
 PtR,(NH3); PtR,(NH3)2; PtR,(NH3)3 ; PtR,(NH3),. 
 
 (R = OH, Cl, Br, I, NO3, etc.). 
 
 Diplatosamine derivatives, <?. g., (NH3)^Pt2R2, are also known. They 
 evidently contain radicals or ions, consisting of platinum, the group NH3, 
 and halogens (also NO3, etc.). They result from the action of ammonia 
 upon platinous chloride. When the acid residues are replaced by 
 hydroxyl groups platinum bases are formed, e. g., Pt(NH3)^(OH)2, which 
 resemble alkaline hydroxides in their chemical properties. The other 
 platinum metals form similar amine derivatives. The nature and chemical 
 constitution of these interesting compounds is, however, not fully ex- 
 plained. 
 
 SPECTRUM ANALYSIS.* 
 
 We observed that many substances, when introduced into a non-lumi- 
 nous flame, imparted to it a characteristic coloration. Thus, the sodium 
 
 * A more exhaustive, concise, and distinct presentation of the s])ectrum plienomena may 
 be found in Herman VV. Vogel’s “ Prakti.sche .Spektralanalyse irdi.scher Stoffe,” 1889, and 
 in “ Lehrbuch der .Spektralanalyse,” by H. Kayser, 1883. 
 
4i6 
 
 INORGANIC ClIKMIS'I'kV. 
 
 compounds color it yellow; the iiotassiiim, violet; thallium, green, etc. 
 The decomj)osition of the light thus obtained, and, indeed, of every liglu, 
 by means of a prism, and the study of the resulting spectrum, form the 
 basis of spectrum analysis, established in 1859 by Kirchhoff and bunsen. 
 
 Every substance, solid or liquid, heated to white heat {e. g., molten 
 platinum; lime heated in the oxyhydrogen flame; the ordinary flame 
 containing glowing particles of carbon) emits rays of every refrangi- 
 bility ; and hence furnishes a cotitinuous speci 7 'uin, which firings to view 
 all the colors of the rainbow, from red to violet, if the light be conducted 
 through a prism. Glowing gases and vajiors, on the contrary, whose 
 molecules can execute unobstructed oscillations, emit light of definite 
 refrangibility, and, therefore, give spectra, consisting of single, bright 
 lines. Thus, the spectrum of the yellow sodium flame is recognized as 
 
 composed of one very bright yellow line, which by increased magnifying 
 power is shown to consist of two lines lying very near each other. The 
 violet potassium light gives a spectrum, consisting of a red and a blue 
 line. The crimson strontium light shows in the spectrum several dis- 
 tinct red lines and a blue line. (See the Spectrum Plate.) Each of these 
 lines corresjionds to rays of definite wave-length, which are differently 
 refracted from all the others; hence the lines appear in the same spec- 
 trum at the same point, when rays of their wave-length fall on the slit. 
 
 If substances affording different colors be introduced into a flame, the 
 most intense color generally obscures the others ; in the spectrum, how- 
 ever, each individual substance shows its peculiar bright lines, which 
 ai)pcar simultaneously or succeed each other, according to the volatility 
 of the various substances. 
 
SPECTRUM ANALYSIS. 
 
 417 
 
 The spectrum apparatus or spectroscope, shown in Fig. 69, serves for 
 the observation of the spectra. 
 
 In the middle of the apparatus is a flint-glass prism P. At the further 
 end of the tube A is a movable vertical slit, in front of which is placed 
 the light to be investigated. The entering light rays are directed by a 
 collecting lens into the tube A, upon the prism, and the refracted rays 
 (the spectrum) are observed by the telescope B. The tube C is employed 
 to ascertain the relative position of the spectrum lines. This is provided 
 at the outer end of the tube S with a transparent horizontal scale. When 
 a luminous flame is placed before the scale its divisions are reflected from 
 the prism surface and thrown into the telescope B. We then see the 
 spectrum to be studied and the scale divisions in B at the same time, and 
 can readily determine the relative position of the lines of the spectrum. 
 To study two spectra at the same time, and compare them, a three-sided, 
 right-angled glass prism is attached in front of one-half (the lower or 
 upper) of the slit of the tube A ; this directs the rays of a light placed 
 at the side (/, Fig. 69) through A upon the prism P. By means of B, 
 two horizontal spectra will be observed, one above the other, and be- 
 tween are the bright divisions of the scale. 
 
 Adjustment of the Spectroscope. — To observe the spectra in the apparatus described, it 
 is necessary first to adjust the same correctly. The tube A (objective or collimator tube) 
 contains, besides the slit, also a lens (collimator lens), which serves to render parallel the 
 bunch of rays proceeding from the slit ; hence, the latter must be accurately placed in 
 the focus of the lens. This is best accomplished as follows : The telescope {B) is drawn 
 out and adjusted for some distant object, that it may be adapted for the reception of 
 parallel rays ; observe the slit illuminated by the sodium chloride flame, and change its 
 position with reference to the collimator lens until its image appears sharply defined in 
 the telescope B. The scale in tube .S' is similarly adjusted. To have the spectrum lines 
 sharply defined, the slit must be made quite narrow ; for feebly luminous lines it must, 
 however, be widened. 
 
 The determination of the position of the lines of the spectra is usually effected by 
 means of a scale. Since the refraction and dispersion of the rays and therefore the ex- 
 pansion of the spectrum and the length-relations of its variously colored parts are influ- 
 enced by the angle and the quality of the glass of the prism, the scale indications of 
 different forms of apparatus are not directly comparable. To attain this it is ordinarily 
 sufficient to indicate how the most important Fraunhofer lines (see below) are divided 
 upon the scale. 
 
 The spectra represented on the plate are then produced by the fre- 
 quently mentioned colors shown by compounds of the alkali and alkaline 
 earth metals when heated in the non-luminous flame of the Bunsen burner. 
 
 The other metals, however, require a much higher temperature than 
 that of the gas flame for their conversion into gases. To vaporize them 
 and observe their spectra, the electric spark is made to pass from elec- 
 trodes constructed of them. A very convenient and universally appli- 
 cable procedure consists in passing the sparks between the surface of the 
 solution of the salt, into which the negative pole dips, and a platinum 
 wure (positive pole) placed before the slit of the spectroscope [R. Bun- 
 sen, Pogg. Ann. (1873) I55» 230, 336; Z. f. anal. Ch. (1876) 15, 68, 
 and M. Lecoq de Boisbaudran, S])ectres lumineux, Paris, 1874]. Many 
 of the metal spectra, obtained in this way, are remarkably rich in lines. 
 Thus, over 450 lines have been established for iron. 
 
4i8 
 
 ORGANIC CHRMISIRY. 
 
 Tlie speclra of the gases may l)e obtained l)y means of llie ('.cissler 
 tubes. Tlie latter contain tlie gas in a rarefied condition and they give 
 their own i)eculiar colors on jiassing induction sparks through tliem. 
 Hydrogen gives a red light ; its spectrum consists of a bright red, a bine, 
 and a green line. Nitrogen gives a violet light and affords a s[)ectrum 
 of many lines and bands. 
 
 These methods give ns a means for readily distinguishing the individual 
 chemical elements, and e\‘en detecting them in traces. Since the year 
 i860 various new elements, <f. , cmsinm, rubidium, thallium, indium, 
 
 scandium, gallium, germanium, and several others, not accurately studied, 
 have been discovered by their aid. Argon, helium and the other sub- 
 stances discovered recently in the air are esi)ecially distinguished by their 
 spectra. 
 
 In addition to the bright emissioti spectra just described there are yet 
 dark absorption spectra. If a white light giving an uninterru|)ted spec- 
 trum be allowed to pass through different transparent bodies, the latter 
 will absorb rays of definite refrangibility, allowing all others to ])ass. 
 Therefore we observe the sun spectrum interrupted by dark lines or bands 
 in the spectroscope. The solutions of didymium and erbium absorb cer- 
 tain rays, and show corresponding dark lines in the spectrum. The 
 gases deport themselves similarly. White light that has traversed a broad 
 layer of air shows several dark lines in the spectrum; the air absorbed 
 the corresponding rays. This power of absorption is possessed to a much 
 higher degree by all incandescent gases or vapors. If a white light, like 
 the Drummond calcium light, be conducted through the yellow sodium 
 flame (through glowing sodium vapors), a dark line will appear in the 
 spectrum of the white light, and its position will correspond exactly to 
 that of the yellow sodium line; the latter thus appears converted into a 
 dark line. Such spectra are designated the inverted spectra of the corre- 
 sponding metals. The inverted spectra of all elements may be obtained 
 in this way, and they correspond accurately to the direct bright spectra. 
 The cause of these phenomena lies in the proposition deduced by Kirch- 
 hoff from the undulatory theory of light, that the ratio betiveen the emis- 
 sive a 7 id absorptive potver is the same for almost all bodies at like tempera- 
 tures. According to this incandescent gases only absorb rays of just the 
 same refrangibility as those which they emit. For example, when bright 
 white light is passed through the yellow sodium flame the yellow rays of 
 the former are absorbed and retained, while all others pass on almost 
 entirely unaltered. Therefore, in the rainbow spectrum of white light 
 the yellow rays of definite refrangibility will be absent ; and if the other 
 refracted rays of the white light are brighter than the yellow rays emitted 
 from the sodium flame, the latter will be relatively darker ; a dark line 
 will therefore make its appearance. 
 
 These jihenomena have presented a new and wide province to spectral 
 analysis, inasmuch as they open up avenues for the investigation of the 
 chemical nature of the sun and other bodies. 
 
 It is knf)wn that the bright rainbow sun spectrum is intersected by a 
 number of dark lines which have been called the Fraunhofer lines from 
 their discoverer. Kirchhoff has shown that these lines can be easily ac- 
 
SPECTRUM ANALYSIS. 
 
 419 
 
 counted for, after what has already been said, by the following hypothesis 
 upon the nature of the sun ; The latter consists of a solid or liquid lumi- 
 nous nucleus, surrounded by an atmosphere of incandescent gases and 
 vapors. Then the uninterrupted spectrum of the glowing nucleus must 
 be intersected by the dark lines of the inverted spectra of those gases and 
 vapors which occur in the sun’s atmosphere. An accurate comparison of 
 the Fraunhofer lines with the spectrum lines of the various elements has 
 revealed the fact that iron, sodium, magnesium, calcium, chromium, 
 nickel, barium, copper, zinc, and hydrogen are present in the sun’s 
 atmosphere. Tims dark lines have been found in the sun’s spectrum 
 corresponding to all of the 450 lines of the iron spectrum. According 
 to Rowland 34 of the elements known to us are present in the sun ; 15 
 are absent, and much doubt prevails as to the presence of the remainder. 
 Helium, first discovered in the earth in 1895, was known to exist in the 
 sun’s photosphere since 1868. The inferences upon the chemical consti- 
 tution of the sun possess as much and, indeed, a higher degree of proba- 
 bility than falls to many other deductions. 
 
 All the fixed stars thus far investigated possess a constitution like that 
 of the sun. They give spectra intersected by dark lines, and therefore 
 consist of incandescent nuclei surrounded by gaseous atmospheres. The 
 spectra of nebulae, however, only show bright lines ; hence these consist 
 of uncondensed, incandescent masses of vapor. 
 
INDEX 
 
 A. 
 
 Absorptiometer, 120 
 Acetylene, 15 1 , 154 
 Acids, 39, 60, 84 
 haloid, 60 
 meta-, 201 
 normal, 201 
 ortho-, 201 
 oxygen, 6 1 
 radicals, 1 74 
 Active oxygen, 84 
 chlorine, 304 
 Affinity, chemical, 23 
 Air, composition of, 118, 119, 1 20, I21, 
 122, 123 
 destroyed, 118 
 fire, 1 18 
 
 liquefaction of, 1 17 
 liquid, 1 17 
 Algaroth, 1 47 
 Alkali metals, 272 
 Alkalies, 272 
 Alkaline earths, 300 
 recognition of, 310 
 Allotropy, 87 
 Alloys, 255 
 
 Alum, ammonium, 353 
 burnt, 353 
 cubical, 353 
 ordinary, 352 
 Alums, 352 
 Aluminates, 350 
 Aluminium, 346 
 
 bronze, 336, 347 
 chloride, 348 
 hydroxides, 350 
 oxide, 349 
 phosphate, 353 
 silicates, 353 
 sulphate, 352 
 Amalgams, 257 
 Ammonia, 126 
 
 chemical properties of, 128 
 physical properties of, 127 
 quantitative composition of, 129 
 thermo-chemical deportment of, 1 29 
 
 Ammonia-cobalt compounds, 407 
 Ammonium, 129 
 amalgam, 296 
 carbonate, 298 
 chloride, 297 
 compounds, 129, 296 
 hyponitrite, 297 
 nitrate, 297 
 nitride, 298 
 nitrite, 297 
 persulphate, 297 
 phosphates, 298 
 sulphate, 297 
 sulphides, 299 
 Ampere, law of, 75 
 Anglesite, 371 
 Anhydrite, 304 
 Anhydrides of acids, 172 
 Animal charcoal, 15 1 
 Anion, 92 
 Anode, 92 
 Anthracite, 15 1 
 Antichlor, 289 
 Antimony, 146 
 acids, 224 
 bromide, 148 
 butter of, 147 
 chlorides, 147, 148 
 cinnabar, 225 
 fluorides, 148 
 iodides, 148 
 mirror, 147 
 oxides, 223, 224 
 oxychloride, 147 
 spots, 147 
 sulphides, 225 
 Antimonyl, 224 
 Antozone, 88 
 Apatite, 304 
 Apparatus of Carr^, 127 
 Marsh, 144 
 Aqua regia, 203 
 Aragonite, 305 
 Argentan, 336, 405 
 I Argon, 1 16, 124 
 I Argyroditc, 362 
 
 421 
 
422 
 
 INDEX. 
 
 Arsenic, 143 
 
 acids, 221, 222 
 bloom, 220 
 bromide, 146 
 chloride, 146 
 fluoride, 146 
 iodides, 146 
 Marsh’s test for, 145 
 mirror, 144 
 pentoxide, 222 
 spots, 144 
 sulphides, 222, 223 
 sulpho-salts, 223 
 trioxide, 220 
 Arsenious acid, 221 
 Arsenites, 221 
 Arsine, 144, I45 
 Artificial cinnabar, 327 
 Atmosphere, 1 16 
 
 gases discovered in, 123, 124 
 Atomic compounds, 170 
 heat, 253 
 hypothesis, 69, 72 
 -molecular theory, 73 
 volume, 252 
 weight, 167 
 
 choice of, 69 
 relative, 25, 72 
 table of, 26 
 thermal, 254 
 Atomicity, 165 
 Atoms, 24, 75, 167 
 Auric acid, 344 
 
 compounds, 344, 345 
 Auripigment, 223 
 Aurous compounds, 343 
 Avogadro, law of, 75, 98 
 Azoimide, 132 
 Azote, 118 
 Azurite, 335 
 
 B. 
 
 Barium, 308 
 
 carbonate, 310 
 nitrate, 309 
 oxide, 308 
 peroxide, 309 
 persulphate, 309 • 
 sulphate, 309 
 sulphide, 310 
 T^aryta water, 309 
 Bases, 39, 60, 84 
 basicity of acids, 1 72, 214 
 Bauxite, 353 
 Beryllium, 317, 318 
 l>csscmer steel, 396 
 Bismuth, 148, 372, 373 
 alloys, 373 
 miiuic, 373 
 
 Bismuth, oxide, 373 
 oxychloride, 372 
 subnitrate, 373 
 Bismuthic acid, 373 
 Bleaching, 53 
 lime, 303 
 Borax, 243, 294 
 Boric acid, 242 
 
 meta, 242 
 
 Boro- fluoride, hydrogen, 242 
 l)oron, 241, 294 
 carbide, 243 
 chloride, 241 
 fluoride, 242 
 hydride, 241 
 nitride, 243 
 trioxide, 242 
 Brass, 335 
 Britannia metal, 364 
 Bromic acid, 179 
 Bromine, 53 
 solid, 54 
 Bronzes, 336 
 
 C. 
 
 Cadmium, 320 
 oxide, 320 
 sulphate, 320 
 sulphide, 321 
 Cadmous oxide, 320 
 Caesium, 284 
 Calamine, 319 
 Calcite, 305 
 Calcium, 300 
 carbide, 307 
 carbonate, 305 
 chloride, 302 
 fluoride, 302 
 hydride, 301 
 hydroxide, 301 
 hypochlorite, 303 
 nitrate, 304 
 nitride, 301 
 oxide, 301 
 peroxide, 302 
 phosphates, 304, 305 
 silicate, 306 
 sulphate, 304 
 sulphides, 307 
 Calomel, 323 
 Calorie, 22, 66, 253 
 Caput mortuum, 194, 399 
 Carbides, metallic, 257 
 Carbon, 149 
 
 amorjflious, 151 
 chlorides, 159 
 
 compounds with hydnigen, 151 
 dioxide, 227, 228, 229, 230 
 liquid, 228 
 
INDEX. 
 
 423 
 
 Carbon bisulphide, 234 
 gas, 15 1 
 group, 149 
 
 monoxide, 231, 232, 233 
 oxysulphide, 235 
 Carbonates, 23 1 
 Carbonic acid, 230, 231 
 
 amido-derivatives of, 234 
 Carbonyl chloride, 234 
 Carborundum, 163 
 Carry’s ice machine, 127 
 Carnallite, 276, 315 
 Cassiterite, 363 
 Catalytic reactions, 103 
 Caustic potash, 275 
 soda, 286 
 Cement, 302 
 Cementation steel, 396 
 Cerium, 355 
 Cerrusite, 367, 371 
 Chalk, 305 
 Chamber acid, 19 1 
 Chameleon minerale, 39 1 
 Charcoal, 15 1 
 animal, 151 
 Chemical affinity, 23 
 combination, 19 
 constitution, 168 
 elements, 19 
 energy, 23 
 equations, 27 
 formulas, 25, 27 
 industries, importance of, 1 8 
 phenomena, 19 
 structure, 168 
 symbols, 25 
 tension, 23 
 
 Chlor-detonating gas, 57 
 Chloric acid, 177 
 Chloride of lime, 303 
 Chlorine, 49 
 
 dioxide, 176 
 hydrate, 52 
 preparation, 49 
 properties, 51, 52 
 tetroxide, 176 
 thionyl, 185 
 trioxide, 176 
 Chlorites, 176 
 Chlorous acid, 176 
 Chlorsulphonic acid, 195 
 Chromates of potassium, 379 
 Chrome yellow, 380 
 Chromic acid, 377 
 
 anhydride, 378 
 chloranhydrides, 380 
 compounds, 375 
 oxide, 376 
 Chromium, 374 
 
 Chromium group, 373, 374 
 Chromous compounds, 375 
 Chromyl chloride, 380 
 Cinnabar, 322, 327 
 9^y, 353 
 
 Coal, anthracite, 151 
 bituminous, 15 1 
 Cobalt, 406 
 
 ammonia-compounds, 407, 408 
 cyanides, 408 
 Cobaltic compounds, 407 
 CobaltoLis compounds, 406 
 ultramarine, 407 
 Coke, 15 1 
 Colcothar, 399 
 Colloids, 238 
 Combining weight, 70 
 weights, law of, 72 
 Condensation of gases, 47 
 Conservation of energy, 2I 
 Constant proportions, law of, 69, 72 
 Constantan, 405 
 Copper, 330, 331, 332 
 alloys, 335 
 carbonates, 335 
 hydride, 333 
 sulphate, 334 
 Corrosive sublimate, 325 
 Corundum, 346, 349 
 Critical condition, 47 
 pressure, 47, 229 
 temperature, 47 
 volume, 47 
 Crystallography, 31 
 Crystalloids, 238 
 Cupric compounds, 334 
 Cuprite, 332 
 
 Cuprous compounds, 332 
 Cyanogen, 235 
 
 compounds, 235 
 
 derivatives of iron, 401, 402, 403 
 
 D. 
 
 Davy’s lamp, 59 
 Decipium, 27, 356 
 Density, gas, 74, 79, 80 
 Determined compounds, c2 
 Dialysis, 237 
 Diamide, 131 
 
 Diammonium compounds, 310, 311, 312 
 Diamond, 1 50 
 Didymium, 355 
 Dissociation, 94 
 Dithionic acid, 198 
 Doctrine of linking of atoms, 172 
 I )oberciner’s lamp, 46 
 Drummond light, 83 
 ‘ Dysprosium, 27, 356 
 
424 
 
 INDEX. 
 
 E. 
 
 Earth metals, 345, 346, 354 
 Eau de Javelle, 278 
 
 Labarraque, 278 
 
 Electro-chemical equivalents, 265 
 Electrolysis of salts, 262 
 Electrolytic dissociation theory, 269 
 I^lements, chemical, 19 
 Endothermic compounds, 30 
 reactions, 30, 67 
 Energy, 22, 29 
 chemical, 22 
 degradation of, 28 
 electric, 22 
 mechanical, 22 
 radiant, 22- 
 Entropy, 23 
 Equivalence, 167 
 Equivalent weight, 167 
 Erbium, 356 
 Ethane, 153 
 Ethylene, 153 
 
 tetrachloride, 159 
 Eudiometer, 120 
 Eudiometric analysis, 120 
 Exothermic compounds, 30 
 reactions, 30, 67 
 
 E. 
 
 Faraday’s law, 263 
 Ferric acid, 401 
 chloride, 400 
 compounds, 399, 400 
 Ferrous compounds, 397, 398 
 sulphate, 398 
 sulphide, 399 
 Fire-damp, 152 
 Fischer’s salt, 407 
 
 Flame, nature of, 155, 156, 157, 158 
 oxidizing, 158 
 reducing, 158 
 Fluorine, 56 
 Fluorite, 30 1 
 Formulas, structural, 168 
 Fulminating gold, 348 
 silver, 338 
 
 Furnace, electric, 253 
 
 G. 
 
 Gadolinite, 354 
 Gadolinium, 27, 356 
 Gahnite, 351 
 Galenite, 367 
 Gallium, 358 
 
 Gallium compounds, 358 
 group, 355 
 (ias, laughing, 21 1 
 
 condensation of, 47 
 critical condition of, 47 
 pressure of, 47 
 volume of, 47 
 density of, 79 
 Gases, diffusion of, 123 
 
 drying and purifying, 42 
 general properties of, 73 
 measuring of, 120 
 General sciences, 17 
 Germanium, 164, 238, 362 
 Germanic compounds, 363 
 Germanous compounds, 362 
 Glass, 306 
 Glauber’s salt, 287 
 Glover tower, 190 
 Glucinum, 317 
 Gold, 342, 343 
 
 chlorides, 343, 344 
 fulminating, 345 
 oxides, 344 
 Graphite, 150 
 Graphitite, 150 
 Green cinnabar, 407 
 Greenockite, 321 
 Groups, 167 
 Guignet’s green, 376 
 Gunpowder, 280 
 Gypsum, 304 
 
 H. 
 
 Halogen compounds of metals, 257 
 hydrides, 57 
 Halogenides, 92 
 Halogens, 49 
 
 compounds with one another, 68 
 . oxygen compounds of, 173 
 thermo-chemical deportment of, 66 
 Haloid acids, 57, 61 
 salts, 61 
 
 Heat, atomic, 253 
 latent, 90 
 modulus, 31, 66 
 specific, 253 
 unit of, 66, 89 
 Helium, 116, 125 
 Hepar, 283 
 Ilolmium, 27, 356 
 Horn-silver, 339 
 Hydrates, 258 
 Hydraulic cement, 302 
 Hydrazine, 131 
 hydrates, 31 1 
 * Hydrazoic acid, 132 
 
INDEX. 
 
 425 
 
 Hydriodic acid, 63 
 Hydrobroraic acid, 62 
 Hydrocarbons, saturated, 153 
 unsaturated, 153 
 Hydrochloric acid, 59 
 Hydrochlorpalladic aoid, 413 
 Hydrocyanic acid, 236 
 Hydrofluoric acid, 65 
 Hydrosulphuric acid, 109 
 Hydrosulphurous acid, 86 
 Hydrogen, 40 
 
 antimonide, 147 
 bromide, 61 
 chloride, 57 
 
 chemical properties of, 60 
 cyanide, 236 
 fluoride, 64 
 iodide, 63 
 pentasulphide, iio 
 peroxide, 99 
 
 detection of, 103 
 properties of, loi 
 thermo-chemistry of, 103 
 persulphide, iio 
 preparation, 40 
 properties, 43, 45 
 selenide, 1 12 
 silicofluoride, 163 
 sulphide, 107, 108 
 
 molecular formula of, 109 
 thermo-chemical deportment of, 
 109 
 
 telluride, 113 
 Hydroxides, 84 
 Hydroxyl, 104 
 Hydroxylamine, 130, 131 
 Hypobromous acid, 179 
 Hypochlorous acid, 1 74 
 oxide, 174 
 Hyponitric acid, 206 
 Hyponitrous acid, 212 
 Hypophosphoric acid, 216 
 Hypophosphorous acid, 214 
 Hypophosphites, 215 
 Hyposulphites, 197 
 Hyposulphurous acid, 197 
 
 I. 
 
 Illuminating gas, 155 
 
 Indestructibility of matter, principle of, 20 
 Indium, 358 
 
 compounds, 359 
 Iodic acid, 180 
 
 anhydride, l8l 
 Iodine, 54, 55 
 bromide, 69 
 chloride, 68 
 
 36 
 
 Iodine fluoride, 69 
 trichloride, 69 
 lodonium hydroxide, 1 80 
 Iridium, 41 1 
 Iron, 393, 394 
 alums, 400 
 carbide, 394, 402 
 carbonyl, 404 
 cast, 394 
 disulphide, 400 
 fluid, 394 
 group, 392 
 metallurgy, 395, 396 
 pentacarbonyl, 404 
 technical, 394 
 tetracarbonyl, 404 
 vitriol, 398 
 weld, 394 
 wrought, 394 
 Isomerism, 87 
 Isomorphism, 255 
 
 — V K. 
 
 Kaolin, 353 
 Kathode, 92 
 Kation, 92 
 Kelp, 54 
 
 Kermes mineral, 225 
 Kieserite, 287, 315 
 Krypton, 1 16 
 
 L. 
 
 Lanthanum, 355 
 Lapis lazuli, 354 
 Latent heat, 90 
 Laughing gas, 21 1 
 Law of Boyle, 74 
 
 Dalton, 74 
 
 definite proportions, 69 
 Faraday, 263 
 Gay-Lussac, 74 
 Mariotte, 74 
 Lead, 164, 238, 364 
 alloys, 368 
 carbonate, 37 1 
 chamber crystals, 191 
 chloride, 370 
 chromate, 380 
 iodide, 370 
 nitrate, 37 1 
 oxides, 369 
 peroxide, 369 
 red, 369 
 sulphate, 371 
 sulphide, 371 
 tetrabromide, 371 
 tetrachloride, 369, 370 
 
INDEX. 
 
 426 
 
 Lead telra-iodide, 37 1 
 tree, 368 
 wliite, 371 
 Lignite, 151 
 Lime, 301 
 
 cliloride of, 303 
 light, 83 
 Litharge, 369 
 Lithium, 295 
 
 carbonate, 296 
 chloride, 296 
 hydride, 295 
 phosphate, 296 
 Lunar caustic, 341 
 
 M. 
 
 Magnesia, 314 
 usta, 314 
 Magnesium, 313 
 
 group, metals of, 312 
 carbonate, 316 
 chloride, 314 
 nitride, 316 
 oxide, 314 
 phosphates, 315 
 sulphate, 315 
 Magnetite, 399 
 Malachite, 335 
 Manganese, 386 
 alums, 384 
 bronze, 336 
 heptoxide, 39 1 
 peroxide, 389 
 Manganic acid, 390 
 compounds, 388 
 Manganin, 405 
 Manganites, 88, 389, 390 
 Manganous compounds, 387 
 sulphate, 388 
 Marsh gas, 1 52 
 test, 145 
 Mass action, 271 
 Massicot, 369 
 Matter, 18 
 
 Maximum valence, 170 
 Mechanical equivalent of heat, 22 
 mixture, 18 
 
 Mercuric compounds, 325, 326 
 Mercurous compounds, 323, 324, 325 
 Mercury, 321, 323, 323 
 Meta-acids, 201 
 
 j)hos])lioric acid, 217 
 stannic acid, 367 
 Metallic ions, 402 
 Metalloids (non-metals), 19, 39 
 oxygen compounds of, 172 
 Mcrtals, 19, 39, 251 
 
 Metals, heavy, 252 
 light, 252 
 
 properties of, 251, 255 
 rare, 354 
 Metargon, 116 
 Methane, 152, 153 
 Mineral water, 91 
 Mohr’s salt, 398 
 Mol, 98 
 
 Mol-volume, 98 
 Molecular compounds, 170 
 Molecules, 24, 25, 75, 167 
 Molybdenum, 382 
 
 compounds, 382, 383 
 Molybdic acid, 383 
 Mosaic gold, 367 
 Mosandrium, 26 
 Multiple proportions, 71 
 law of, 72 
 Muriatic acid, 59 
 
 N. 
 
 Neodymium, 355 
 Neon, 1 16 
 Nickel, 404 
 alloys, 405 
 carbonyl, 233 
 oxides, 405 
 plating, 405 
 sulphate, 405 
 tetracarbonyl, 405 
 Niobium, 226 
 Nitramide, 204 
 Nitrates, 202 
 Nitric acid, 202, 203 
 fuming, 203 
 oxide, 209, 210 
 Nitrites, 202, 205 
 Nitrogen, 115, I16 
 carbonyl, 234 
 chloride, 133 
 group, 1 14 
 
 oxygen derivatives of, 201 
 iodides, 134 
 pentasulphide, 213 
 pentoxide, 204 
 tetroxide, 206 
 trioxide, 205 
 sulphide, 213 
 Nitro-metals, 207 
 Nitroso-acid, 208 
 Nitrous acid, 205 
 oxide, 21 1 
 
 Nitrosulphonic acid, 207 
 Nitrosyl chloride, 204 
 sulphuric acid, 207 
 Nitroxyl chloride, 204 
 
INDEX. 
 
 427 
 
 Non-metals, 19, 39 
 Nordhausen sulphuric acid, 194 
 Normal acids, 201 
 
 O. 
 
 Olefiant gas, 154 
 Organic chemistry, 15 1 
 Ortho-acids, 201 
 Osmium, 410, 41 1 
 Osmotic pressure, 267 
 Oxidation, 83 
 
 theory of Traube, loi 
 Oxides, 84, 258 
 indifferent, 84 
 Oxygen, 80 
 
 atomic weight of, 98 
 group, 80 
 liquid, 82 
 oxidized, 87 
 preparation of, 80 
 properties of, 82 
 reduction of, loi 
 salts, 61 
 Ozone, 84 
 
 constitution of, 86 
 preparation of, 85 
 tests for, 86 
 
 thermo-chemical deportment of, 87 
 
 P. 
 
 Palladium, 412 
 
 hydride, 46, 412 
 Palladic compounds, 412, 413 
 Palladious compounds, 412, 413 
 Passive iron, 397 
 Pattinson’s method, 337 
 Pearl ash, 281 
 Pentathionic acid, 199 
 Perbromic acid, 179 
 Percarbonic acid, 231 
 Perchloric acid, 178 
 
 hydrate of, 179 
 Perchromic acid, 374, ^75 
 anhydride, 381 
 Periodates, 279 
 Periodic acid, 181 
 
 hydrate of, 181 
 •sy.stem of, 164, 243 
 Periodicity of chemical valence, 248 
 Permanent white, 309 
 Permanganic acid, 390 
 Peroxides, 259 
 Persulphates, 188 
 Persulphuric acid, 188 
 Philippium, 26 
 Phlogiston, 1 18 
 Phosgene gas, 234 
 
 Phospham, 220 
 Phosphates, 214 
 Phosphine, 138, 139 
 
 molecular formula of, 140 
 Phosphites, 21 3 
 Phosphonium, 140 
 Phosphonium iodide, 140 
 Phosphoric acid, 216 
 
 anhydride, 217 
 hypo-, 216 
 meta-, 217 
 pyro-, 216 
 Phosphorite, 301 
 Pho.sphorous acid, 215 
 oxide, 215 
 
 Phosphorus, 135, 137 
 acids, 214 
 bromides, 142 
 bronze, 336 
 burns, 137 
 chloranhydrides, ,218 
 fluoride, 142 
 iodides, 142 
 metallic, 136 
 oxides, 213 
 oxychloride, 218 
 pentachloride, 141 
 pentoxide, 217 
 red, 136 
 salt of, 298 
 sulphochloride, 219 
 sulphur derivatives, 220 
 trichloride, 141 
 yellow, 136 
 Photography, 340 
 Physical phenomena, 18 
 Pink salt, 366 
 Platinum, 413, 414 
 
 ammonium-compounds, 415 
 black, 413 
 
 cyanogen compounds of, 415 
 metals, 408, 409, 410 
 heavy, 408 
 light, 408 
 separation of, 409 
 sponge, 413 
 
 Platinic compounds, 414 
 Platinous compounds. 414 
 Plumbates, 369, 370 
 Iflumbic acid, 370 
 Polychromates, 378 
 Polysulphides, 283 
 Polythionic acids, 197 
 Porcelain, 354 
 Potashes, 281 
 Potassium, 273 
 borates, 280 
 bromate, 278 
 bromide, 276 
 
INDEX. 
 
 428 
 
 Potassium carbonate, 280 
 chlorate, 276 
 cliloricle, 276 
 chromates, 379 
 
 compounds, recognition of, 283 
 cyanide, 276 
 ferrocyanide, 402 
 ferricyanide, 403 
 fluoride, 276 
 hydride, 274 
 hydroxide, 275 
 hypochlorite, 278 
 iodate, 278 
 iodide, 276 
 manganate, 390 
 nitrate, 279 
 nitrite, 280 
 oxides, 275 
 percarbonate, 282 
 perchlorate, 278 
 permanganate, 391 
 phosphates, 280 
 silicate, 282 
 sulphates, 279 
 sulphides, 282, 283 
 sulphites, 279 
 Praseodymium, 355 
 Preparing salts, 367 
 Prussian blue, 403 
 Prussic acid, 236 
 Purple of Cassius, 345 
 Pyrites, 400 
 Pyroarsenic acid, 222 
 Pyromorphite, 367 
 Pyrophosphoric acid, 216 
 Pyrosulphuric acid, 193 
 Pyrosulphuryl chloride, 196 
 
 Q- 
 
 Quadrant oxides, 259 
 Quartz, 236 
 (Quicksilver, 322 
 
 R. 
 
 Radiant heat, 21 
 Radical, 167 
 Realgar, 223 
 Reduction, 84 
 Residue, acid, 174 
 Reversible reaction, 91 
 Rhodium, 4II 
 Rimnann’s green, 407 
 Roentgen rays, 21, 415 
 Rose’s metal, 373 
 Rubidium, 284 
 Ruby, 350 
 Ruthenium, 410 
 
 S. 
 
 Safety lamp, 159 
 Salt of phosphorus, 298 
 producers, 49 
 Saltpeter, 279, 296 
 Salts, 60, 259 
 
 ammonium, 129 
 basic, 260 
 double, 260 
 haloid, 61 
 oxygen, 61 
 transposition of, 270 
 Samarium, 356 
 Sapphire, 350 
 Scandium, 355 
 Scheele’s green, 335 
 Schlippe’s salt, 225 
 Schweizer’s reagent, 334 
 Selenium, 1 12 
 acids, 200 
 chlorides, 112 
 hydride, 1 1 2 
 oxide, 199 
 Siderite, 399 
 
 Siemens-Martin method, 397 
 Silica, 236 
 Silicates, 238 
 Silicic acid, 237, 238 
 Silicon, 160 
 
 bromide, 162 
 bronze, 336 
 carbide, 163 
 chloride, i6i 
 chloroform, 162 
 dioxide, 236 
 disulphide, 238 
 
 fluoride, 162 — _ 
 
 hydride, i6i 
 iodide, 162 
 Silver, 336, 337, 338 
 
 allotropic forms of, 338 
 bromide, 340 
 chloride, 339 
 coins, 331 
 cyanide, 342 
 fineness of, 338 
 iodide, 340 
 nitrate, 34I 
 nitride, 341 
 nitrite, 341 
 oxides, 338, 339 
 plating, 342 
 sulphate, 34I 
 sulphide, 341 
 Silvering, 342 
 Slags, 396, 397 
 Slaked lime, 301 
 Smalt, 406 
 Soda, caustic, 286 
 
INDEX. 
 
 429 
 
 Soda, residue, 290 
 Solvay, 291 
 Sodium, 284 
 
 azoimide, 295 
 borate, 294 
 bromide, 287 
 
 carbonate, 289, 290, 291, 292 
 chlorate, 287 
 chloride, 286 
 
 compounds, recognition of, 295 
 Sodium hydroxide, 286 
 hyposulphite, 289 
 iodate, 287 
 iodide, 287 
 nitrate, 293 
 nitride, 295 
 nitrite, 293 
 oxides, 285 
 perchlorate, 287 
 periodate, 287 
 peroxide, 285 
 phosphates, 293, 294 
 silicate, 295 
 sulphate, 287, 288 
 sulphite, 288 
 thiosulphate, 289 
 Soft solder, 368 
 Solutions, 265 
 
 supersaturated, 288 
 theory of dilute, 267 
 thermo-chemistry, 92 
 Special sciences, 17 
 Specific gravity of gases, 73 
 heat, 253 
 volume, 45, 252 
 
 Spectrum analysis, 415, 416, 417, 418, 
 419 
 
 Speiss-cobalt, 406 
 Spinels, 351 
 Stannic acid, 66, 367 
 
 compounds, 366, 367 
 Stannous compounds, 365 
 Status nascens, 53, 77, 88, 102, 130, 175, 
 203 
 
 Steel, 394 
 Stibine, 147 
 Stoechiometric laws, 71 
 S trass, 306 
 Strontium, 307 
 
 carbonate, 308 
 nitrate, 308 
 oxide, 307 
 
 Structure, chemical, 168 
 Suboxides, 259 
 Substitution, 159 
 Sulphamide, 196 
 Sulphates, 193 
 Sulphides, 108 
 Sulphimide, 196 
 
 Sulphites, 186 
 Sulphocarbonic acid, 235 
 Sulpho-group, 198 
 Sulpho-stannates, 367 
 Sulphur, 104, 105, 106 
 bromides, iii 
 chemical properties of, 106 
 chlorides, no, iii 
 dioxide, 183, 184 
 heptoxide, 188 
 iodides, iii 
 
 oxygen compounds of, 183 
 sesquioxide, 187 
 trioxide, 187 
 
 Sulphuric acid, 189, 190, 19 1, 192, 193 
 amides, 196 
 chloranhydrides, 195 
 Sulphuric acid, di-, 193 
 fuming, 194 
 hydrates, 192 
 Nordhausen, 194 
 Sulphurous acid, 185 
 Sulphuryl, 195 
 chloride, 185 
 Superphosphate, 305 
 Sylvite, 276 
 Sympathetic ink, 406 
 
 T. 
 
 Tantalum, 226 
 Tellurium, 113 
 acids, 200 
 bromides, 113 
 chlorides, 113 
 hydride, 1 13 
 Tension of vapors, 90 
 Terbium, 26, 356 
 Tetrathionic acid, 199 
 Thallic acid, 360 
 
 compounds, 360 
 Thallium, 359 
 alum, 360 
 
 Thallous compounds, 360 
 Thenard’s blue, 407 
 
 Thermo-chemistry of the elements, 28, 66, 
 92, 103, 109, 1 14, 129, 133, 134, 137, 
 142, 160, 164, 186, 187, 212, 236, 327, 
 328, 330, 340, 348, 387 
 Thionic acids, I97 
 Thionyl chloride, 185 
 Thiosulphuric acid, 197 
 Thomas slag, 397 
 Thorium, 238, 240 
 Thulium, 27, 356 
 Thyroidin, 55 
 Tin, 164, 238, 363 
 dichloride, 364 
 dioxide, 366 
 
INDEX. 
 
 430 
 
 'l in (lisulpliide, 367 
 monoxide, 365 
 oxysulphide, 367 
 salt, 364 
 stone, 363 
 tetrachloride, 366 
 Titanium, 238, 239 
 Trithionic acid, 199 
 d'ornbac, 336 
 Tungsten, 384 
 
 compounds, 384 
 Turpeth mineral, 327 
 Turnbull’s blue, 403 
 Type metal, 368 
 
 U. 
 
 Uchatius steel, 397 
 Ultramarine, 354 
 Uranium, 384 
 Uranic compounds, 385 
 Uranous compounds, 385 
 Uranyl, 385 
 
 V. 
 
 Valence, 165, 166, 167, 169 
 periodicity of, 248 
 variable, 170 
 Vanadium, 226 
 Varec, 54 
 
 Vitriol, copper, 334 
 green, 398 
 
 oil of (see Sulphuric acid) 
 Volume, atomic, 252 
 specific, 252 
 
 W. 
 
 Water, 88 
 
 chemical properties of, 9 1 
 crystallization, 266 
 constitution. 315 
 dissociation of, 93 
 distilled, 91 
 
 Water, electrolysis of, 92 
 gas, 232 
 glass, 295 
 hard, 91 
 mineral, 91 
 
 molecular formula of, 97 
 natural, 91 
 oxidation of, lOO 
 physical i)roperties of, 88 
 ({uantitative composition of, 95 
 soft, 91 
 
 thermo-chemical deportment of, 92 
 Weight proportions in the union of the ele- 
 ments, 69 
 
 Welsbach mantle, 240 
 White precipitate, 325 
 lead, 371 
 
 Wood’s metal. 373 
 Wulfenite, 3^7 
 
 X. 
 
 Xenon, 1 16 
 
 V. 
 
 Yellow prussiate of potash, 402 
 Ytterbium, 356 
 Yttrium, 355 
 
 Z. 
 
 Zinc, 318 
 
 blende, 319 
 chloride, 319 
 dust, 318 
 oxide, 318 
 silicate, 319 
 sulphate, 319 
 sulphide, 319 
 w'hite, 318 
 Zircon, 239 
 Zirconium, 238, 239 
 light, 83 
 
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 Spirits, Wines, Liquors, etc., etc. A Handbook for all interested 
 in the manufacture and sale of Alcohol and Its Compounds. Edited 
 by John Gardner, f.c.s.. Editor of '‘Cooley’s Cyclopedia” and 
 “ Beasley’s Druggists’ Receipt Book.” Illustrated. Cloth, ^1.50 
 
 “Trustworthy and valuable.” — German and America}! Brezvers' Journal. 
 
 “A very complete handbook.” — Boston Journal of Chemistry. 
 
 Bleaching, Dyeing, and Calico Printing. With P'ormulae ; 
 a Chapter on Dye Stuffs. Illustrated. 121110. Cloth, $1.50 
 
 “A serviceable manual.” — Inventors' and Manufacture}^ Gazette. 
 
 The prices of these books are net. No discount allowed retail purchasers. 
 
10 BOOKS FOR CHEMISTS AND MANUFACTURERS. 
 
 CAMERON. Oils and Varnishes. A Practical Handbook, 
 by^ James Cameron, f.i.c. With Illustrations, Formiiloe, d’ables, 
 etc. i2mo. Cloth, $2.25 
 
 Soap and Candles. A New Handbook for Manufacturers, 
 Chemists, Analysts, etc. 54 Illustrations. 121110. Cloth, $2.00 
 
 EMERGENCIES. 
 
 DULLES. Accidents and Emergencies. A Manual for 
 the Treatment of Surgical and other Injuries, Poisoning and various 
 Domestic Emergencies, in the absence of the Ph}'sician. By 
 Charles W. Dulles, m.d., Surgeon to the Out-Door Department 
 of the University and Presbyterian Hospitals, Philadelphia. Fifth 
 Edition, Enlarged. New Illustrations. 121110. Cloth, ^i.oo 
 
 “This is a revised and enlarged edition, with new illustrations, of the 
 manual, explaining the treatment of surgical and other injuries in the absence of 
 the physician. The simple and practical suggestions of this little book should be 
 known to every one. Accidents are constantly occurring, and a knowledge of 
 what should be done in an emergency is very valuable. Such a handbook should 
 be in every home, placed where it can always be found readily,” — Boston Joiirnal 
 of Education. 
 
 WATER AND MILK ANALYSES. 
 
 LEFFMANN. Examination of Water for Sanitary and 
 Technical Purposes. Presenting those Processes that are Most 
 Trustworthy and Practical. By Henry Leffmann, m.d.. Professor 
 of Chemistry and Metallurgy, Pennsylvania College of Dental 
 Surgery; Professor of Chemistry, Woman’s Medical College of 
 Pennsylvania, etc. Fourth Edition. Revised and Enlarged. Illus- 
 trated. i2mo. Cloth, $1.25 
 
 “ This is a well-compiled and useful little treatise.” — London Lancet. 
 
 “ An admirable digest of our present knowledge.”— of Analytical 
 Chemistry. 
 
 “ Especially valuable is the section on interpretation of results.” — Railroad 
 and Ejigineering Neivs. 
 
 BY THE SAME AUTHOR. 
 
 Analysis of Milk and Milk Products. Arranged to suit 
 the needs of Analytical Chemists, Dairymen, and Milk Inspectors. 
 Second Edition. Enlarged. 121110. Cloth, $1.25 
 
 “ 'fhe book is one which will be useful in the hand of the dairyman, as well 
 as in the hands of those whose duty it is to see that he deals fairly with his cus- 
 tomers.” — London Sanitary Record. 
 
 The prices of these books are net. No discount allowed retail purchasers. 
 
BOOKS FOR SCIENTISTS. 
 
 11 
 
 CARPENTER. The Microscope and Its Revelations. By 
 
 W. B. Carpenter, m d , f r s. Eighth Edition. By Rev. Dr. 
 Dai.ltnger, f r s. Revised and Enlarged, with 800 Illustrations 
 and 21 Lithographs. Octavo. 1100 pages. Pi'eparing. 
 
 “ The book, therefore, cannot fail to be of value to chemists and others at 
 iron works intrusted with the microscopical examination of the metals with which 
 they deal.” — The Americaii Manufacturer, Pittsburgh. 
 
 “ Every one who has a microscope will need also Carpenter’s book to get the 
 most out of his instrument, and every one who has the book will be certain to want 
 a microscope.” — Popular Science Monthly. 
 
 “It is without a rival in its particular field, and is beyond question the best 
 single work on the subject, not only in English but in any other language. . . . 
 
 A splendid specimen of the book-maker’s art.” 
 
 “ The book is more than ever a standard, unrivaled in its kind, and is a neces- 
 sity to every one who pretends to any scientific use of the microscope.” — New York 
 Evening Post. 
 
 WETHERED. Medical Microscopy. By Frank J. Weth- 
 ered, M.D., M.R.c.p. With 100 Illustrations. 121110. 
 
 Cloth, $2.00 
 
 REEVES. Medical Microscopy. Including chapters on 
 Bacteriology, Neoplasms, Urinary Examination, etc. By James E. 
 Reeves, m.d., Ex-President American Public Health Association, 
 etc. Illustrated. 121110. Cloth, $2.50 
 
 THE BEST DICTIONARY. 
 
 GOULD. Illustrated Dictionary of Medicine and Allied 
 Sciences, including Chemistry, Biology (Zoology and 
 Botany;, Hygiene, etc. Large, Square Octavo. 1633 pages. 
 
 Fulfsheq;?,^^^^ } Thumb Index, $11.00 
 
 Fifth Edition. Half Russia, Thumb Index, $12.00 
 
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 it was thought eminently proper to include both these sciences in this book. They 
 are closely related to medicine, and each is largely dependent upon the others, 
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 the general scientist. Bacteriology and Parasitology are particularly well pre- 
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 mains, Resins, Soaps, Stains, Starches, Sugars, Tests, Theories, Wave Lengths, 
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 sold. 
 
 The prices of these books are net. No discount allowed retail purchasers. 
 
12 
 
 HYGIENE AND SANITARY SCIENCE. 
 
 HOTTER and FIRTH. The Theory and Practice of 
 Hygiene. A Complete Treatise by J. Lane Notter, m.a., m d,, 
 F.C.S., Fellow and Member of Council of the Sanitary Institute of 
 Great Britain; Jh'ofessor of Hygiene, Army Medical School; 
 Examiner in Hygiene, University of Cambridge, etc , and R. H. 
 Firth, f.r.c.s., Assistant Professor of Hygiene, Army Medical 
 School, Netly. Illustrated by lo Lithographic Plates and other 
 Illustrations, and including many Useful Tables. Second Edition, 
 Carefully Revised. Octavo. In Press. 
 
 The standard authority in tlie United States and british Army and Navy. 
 
 » 
 
 I 
 
 PARKES. Hygiene and Public Health. A Practical 
 Manual. By Louis C. Parkes, m.d., d.p.h., London Hospital ; 
 Assistant Professor of Hygiene and Public Health at University 
 College, etc. Fifth Edition, Enlarged and Revised! 8o Illustra- 
 tions. 121110. Cloth, $2.50 
 
 “ Dr. Parkes’ experience as a teacher of public hygiene has enabled him to 
 deal with the whole subject in a practical and intelligible manner.” — The Boston 
 Medical and Surgical Journal. 
 
 STEVENSON AND MURPHY. A Treatise on Hy- 
 giene. Illustrated. Edited by Thomas Stevenson, m.d., f.r.c.p., 
 Lecturer on Chemistry and on Medical Jurisprudence at Guy’s Hos- 
 pital, Official Analyst to the Home Office ; and Shirley F. Mur- 
 phy, Medical Officer of Health to the County of London. 
 
 Vol. I. 9 Plates. 186 Illustrations. 1013 pages. Cloth, 56.00 
 Vol. II. 45 Plates. 31 Illustrations. 847 pages. Cloth, 56.00 
 Vol. III. Sanitary Law. 459 pages. Cloth, 55.00 
 
 EACH VOLUME SOLD SEPARATELY. 
 
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 plates, diagrams, and wood-cuts, and pains appear to have been taken to bring the 
 information upon each topic up to date.” — The Boston Medical and Surgical 
 Journal. 
 
 “ All the topics are treated with a thoroughness of detail leaving nothing to 
 be desired. The contents are valuable alike to the physician, the municipal health 
 officer, and the sanitary engineer.” — Medical Record., New York. 
 
 KENWOOD. Public Health Laboratory Work. By H. 
 
 R. Kenwood, m.il, d.p.h., f.c.s.. Instructor in Hygienic Labora- 
 tory, University College; late Assistant Examiner in Hygiene, 
 Science and Art Department, South Kensington, London, etc. 
 With 1 16 Illustrations and 3 Plates. Cloth, 52.00 
 
 ***For complete list of books on Hygiene, send for our Special Catalogue. 
 
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