INORGANIC CHEMISTRY. 
RICHTER. In Preparation. 
A TEXT-BOOK. 
OF 
ORGANIC CHEMISTRY. 
BY 
AUTHORIZED TRANSLATION FROM LAST GERMAN EDITION. 
PROFESSOR VICTOR VON RICHTER. 
12mo. A Companion Volume to Inorganic Chemistry. A 
TEXT-BOOK 
OF 
INORGANIC CHEMISTRY 
BY 
PROF. VICTOR von RICHTER, 
UNIVERSITY OF BRESLAU. 
AUTHORIZED TRANSLATION OF THIRD GERMAN EDITION, 
BY 
PROF. OF CHEMISTRY IN WITTENBERG COLLEGE, SPRINGFIELD, OHIO; FORMERLY IN THE 
LABORATORIES OF THE UNIVERSITY OF PENNSYLVANIA AND MUHLENBERG COLLEGE: 
MEMBER OF THE CHEMICAL SOCIETIES OF BERLIN AND PARIS, OF THE ACA- 
DEMY OF NATURAL SCIENCES OF PHILADELPHIA, ETC. 
EDGAR F. SMITH, A.M., Ph.D., 
EIGHTY-NINE ILLUSTRATIONS ON WOOD 
COLORED LITHOGRAPHIC PLATE OF SPECTRA 
AND 
PHILADELPHIA: 
P. BLAKISTON, SON & CO., 
No. 1012 Walnut Street, 
1883. Entered according to Act of Congress, in the year 1883, by 
In the Office of the Librarian of Congress, at Washington, D. C. 
P. BLAKISTON, SON & CO., PREFACE 
The success of Prof, vox 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 stu- 
dent, 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 performed, with a view of drawing 
conclusions from the same and proving the intimate connec- 
tion 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 vox 
Richter’s work, occasioning the rapid appearance of three 
editions, the common verdict would seem to be unanimously 
in favor of its inductive methods. 
Iii the third edition, of which the present is a translation, 
the Periodic System of the Elements, as announced by Mex- 
delejeff and Lothar Meyer, is somewhat different, in the 
manner of development and presentation, from that appearing 
in the previous editions. This was done to give more promi- 
nence to and make more general the interesting relations dis- 
closed 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 IV 
PREFACE 
arriving at new relations 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 present volume. In ac- 
cordance 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 periodicity. “Hereby more importance is 
attributed to the principle of the greatest heat development 
than at present appears 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 opportunity 
of deductively obtaining the majority of facts from the ther- 
mal numbers, on the basis of a simple principle. To facilitate 
matters, there is appended to the volume a table containing 
the heat of formation of the most important 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 OE CONTENTS. 
INTRODUCTION. 
Definition of Chemistry, 9. Chemical Elements, 12. Principle of 
Conservation of Substance, 12. Principle of Conservation ot 
Force, 13. Chemical Energy, 15. Conditions ot Chemical 
Action, 15. Chemical Symbols and Formulas, 17. Constitution 
of Matter, 19. Crystallography, 20. 
SPECIAL PART 
Classification of the Elements, 30. 
Hydrogen, 31. Purifying and Drying of Gases, 33. Apparatus 
for Collection of Gases, 33. Condensation of Gases, 35. Chem- 
ical Properties of Hydrogen, 37. 
Group of Halogens, 39. 
Chlorine, 39. Bromine, 43. Iodine, 44. Fluorine, 46. Gen- 
eral Characteristics of the Halogens, 47. Compounds of the Hal- 
ogens with Hydrogen, 47. Hydrogen Chloride, 47. Acids, 
Bases, Salts, 50. Hydrogen Bromide, 52. Hydrogen Iodide, 53. 
Hydrogen Fluoride, 55. Characteristics of the Hydrogen-Halo- 
gen Compounds, 56. Heat of Combination of the latter, 57. 
Compounds of the Halogens with each other, 58. 
Weight Proportions in the Union of the Elements, Laic of Constant 
Proportions, Atomic Theory, 59. Gas Density of Compounds, 
Volume Proportions in the Union of Gases, Atomic Molecular 
Theory, 62. Law of Avogadro, 68. Status nascens, 69. 
Oxygen Group, 71. 
Oxygen, 71. Oxidation and Reduction, 75. Ozone, 76. Iso- 
merism and Allotropy, 79. Compounds of Oxygen with Hydro- 
gen, 80. Water, 80. Solutions, 82. Dissociation, 85. Quanti- 
tative Composition of Water, 86. Molecular Formula of Water, 
Atomic Weight of Oxygen, 87. Hydrogen Peroxide. 90. Chemi- 
cal Molecules, 92. Sulphur, 94. Molecules of the Elements, 96. 
Compounds of Sulphur with Hydrogen, 97. Hydrogen Sulphide, 
97. Hydrogen Persulphide, 100. Compounds of Sulphur with 
Halogens, 101. Selenium, 102. Tellurium, 103. Characteris- 
tics of the Elements of the Oxygen Group, 104. Heat of Combi- 
nation of their Hydrogen Compounds, 105. VI 
TABLE OF CONTENTS 
Nitrogen Group, 106. 
Nitrogen, 106. Atmosphere, 108. Eudiometry, 111. Measur- 
ing Gases, 112. Diffusion of Gases, 113. Compounds of Nitro- 
gen with Hydrogen, 115. Ammonia, 115. Ammonium Salts, 
118. Atomic Weight of Nitrogen, 120. Hydroxylamine, 121. 
Compounds of Nitrogen with the Halogens, 122. Phosphorus, 
123. Compounds of Phosphorus with Hydrogen, 126. Com- 
pounds of Phosphorus with the Halogens, 129. Arsenic, 131. 
Arsine, 132. Compounds of Arsenic with the Halogens, 134. 
Antimony, 135. Stibine, 136. Compounds of Antimony with 
the Halogens, 136. Vanadium, Niobium, Tantalum, 137. Char- 
acteristics of the Elements of the Nitrogen Group, 139. Thermal 
Deportment, 139. 
Carbon Group, 140. 
Carbon, 140. Compounds of Carbon with Hydrogen, 142. 
Methyl, 143. Atomic Weight of Carbon, 143. Ethyl, 144. 
Ethylene, 144. Acetylene, 145. Nature of Flame, 145. Com- 
pounds of Carbon with the Halogens, 151. Silicon, 152. Hy- 
drogen Silicide, 152. Compounds of Silicon with the Halogens, 
153. 
Atom and Molecule, 157. Determination of Molecular Size from 
Chemical Reactions, 159. Quantivalence of the Elements, 160. 
Chemical Structure, 163. 
Oxygen Compounds of the Metalloids, 168. 
Oxygen Compounds of the Halogens, 169. Oxides of Chlorine, 
170. Chloric Oxide, 172. Hypochlorous Oxide, 170. Chlorine 
Trioxide, 171. Chlorous Acid, 171. Chlorine Tetroxide, 171. 
Chloric Acid, 172. Perchloric Acid, 172. Oxides of Bromine 
and Iodine, 173. Hydrates of the Acids. 174. Heat of Combi- 
nation of the Oxygen Compounds of the Halogens, 176. 
Oxygen Compounds of the Elements of the Sulphur Group, 177. 
Oxides' of Sulphur, 177. Sulphur Dioxide, 177. Sulphurous 
Acid, 179. Hydrosulphurous Acid, 180. Sulphur Trioxide, 181. 
Sulphuric Ariel, 181. Disulphuric Acid, 186. Chloranhydride 
of Sulphuric Acid, 187. Polythionic Acids, 188. Thiosuiphuric 
Acid, 189. Oxides of Selenium and Tellurium, 191. Heat of 
Combination of the Oxides and Acids of the Elements of the Sul- 
phur Group, 192. 
Oxygen Compounds of the Elements of the Nitrogen Group, 193. 
Oxides of Nitrogen, 194. Nitric Acid, 194. Nitrogen Pentox- 
ide, 196. Nitrogen Trioxide, 196. Nitrous Acid, 197. Nitro- 
gen Tetroxide, 197. Nitrosylsulphuric Acid, 199. Nitrogen 
Oxide, 201. Hyponitrous Oxide, 203. Hyponitrous Acid, 204. 
Heat of Combination of Oxides and Acids of Nitrogen, 204. 
Oxides of Phosphorus, 205. Hypophosphorous Acid, 206. Phos- 
phorous Acid, 207. Phosphoric Acids, 208. Phosphorus Pentox- 
ide, 209. Cliloranhydrides of Phosphoric Acid, 210. Phospho- 
rus Sulphides, 210. Oxides of Arsenic, 211. Arsenic Trioxido, TABLE OF CONTENTS. 
211. Arsenic Acid, 212. Compounds of Arsenic with Sulphur, 
213. Sulpho-Salts 214. Oxides of Antimony, 214. Antimonic 
Oxide, 215. Antimonic Acid, 216. Antimony Sulphides, 216. 
Heat of Combination of Acids of the Nitrogen Group, 217. 
Oxygen Compounds of the Elements of the Carbon Group, 218. 
Oxides of Carbon, 219. Carbon Dioxide, 219. Carbon Monox- 
ide, 222. Sulphides of Carbon, 223. Cyanogen Compounds, 224. 
Heat of Formation of Carbon Compounds, 225. Oxides of Sili- 
con, 226. Dialysis, 227. Crystalloids and Colloids, 227. Sili- 
cates, 227. Titanium, Zirconium, Thorium, 228. Boron, 230. 
Boron Chloride, 231. Boron Fluoride, 232. Boric Acid, 232. 
The Periodic System of the Elements, 233. 
Periodicity of Chemical Valence, 240. Periodicity of Thermo- 
Chemical Phenomena, 242. 
METALS. 
Physical Properties of the Metals, 244. Atomic Volumes, 245. Spe- 
cific Heat, Atomic Heat, 247. Isomorphism, 251. Chemical 
Properties of Metals, 252. Alloys, 252. Halogen Compounds, 
253. Oxides and Hydrates, 254. Peroxides, Salts, 256. Action 
of Metals upon Salts, 258. Electrolysis of Salts, 260. Double 
Transposition of Salts, 263. 
Group of Alkalies, 265. 
Heat of Combination of the Alkali Metals, 266. Potassium, 267. 
Potassium Hydrate, 268. Potassium Chloride, 269. Potassium 
Chlorate, 270. Potassium Hypochlorite, 271. Potassium Sul- 
phate, 272. Potassium Nitrate, 272. Gunpowder, 273. Potas- 
sium Carbonate, 273. Potassium Silicate, 274. Potassium Sul- 
phides, 275. Recognition of Potassium Compounds, 276. Ru 
bidium, 276. Caesium, 2/6. Sodium, 277. Sodium Hydrate, 277. 
Sodium Chloride, 278. Sodium Sulphate, 279. Supersaturated 
Solutions, 280. Sodium Hyposulphite, 281. Sodium Carbonate, 
282. Sodium Nitrate, 284. Sodium Phosphates, 284. Borax, 
285. Recognition of Sodium Compounds, 286. Lithium, 286. 
Ammonium Compounds, 287. Ammonium Chloride, 288. Am- 
monium Carbonate, 289. Ammonium Sulphide, 290. 
Metals of Group II, 290. 
Group of Metals of the Alkaline Earths, 291. Calcium, 292. Cal- 
cium Oxide, 292. Cement, 293. Calcium Chloride, 294. Calci- 
um Fluoride, 294. Chloride of Lime, 295. Calcium Sulphate, 296. 
Calcium Phosphates, 297. Calcium Carbonate, 298. Glass, 296. 
Calcium Sulphides, 299. Strontium, 300. Barium, 301. Barium 
Oxide, 301. Barium Peroxide, 302. Recognition of the Alka- 
line Earths, 303. TABLE OF CONTEXTS. 
Maqnesium Group, 308. Magnesium, 305. Magnesia, 306. Mag- 
nesium Chloride, 306. Magnesium Phosphates, 307. Magnesium 
Carbonates, 308. Recognition of Magnesium Compounds, 309. 
Beryllium, 309. Zinc, 310. Zinc Oxide, 311. Zinc Sulphate, 
311. Zinc Sulphide, 312. Cadmium, 312. Heat of formation 
of the Metals of Group II, 314. 
Mercury, 316. Amalgams, 317. Mercurous Compounds, 318. 
Mercuric Compounds, 319. . . 
Copper, Silver, Gold, 322. General Characteristics, 324. forms 
of Combination, 325. Copper, 325; Metallurgy of Copper, 325. 
Cuprous Compounds, 327. Cupric Compounds, 328. Copper 
Sulphate, 329. Alloys of Copper, 330. Silver, 331. Metallurgy, 
331. Silver Oxide, 333. Silver Chloride, 334. Photography, 
335; Silver Nitrate, 335. Silvering, 336. Gold, 337. Aurous 
Compounds, 338. Auric Compounds, 339. 
Metals of Group III, 340. . . 
Group of Earth Metals, 342. Aluminium, 342- Aluminium Chlo- 
ride, 343. Aluminium Oxide, 344. Aluminates, 345. Alum. 
347. Aluminium Silicates, 349. Porcelain, 349. Ultramarine, 
349 Rare Earths, 350. Scandium, 350. Yttrium, 350. _ Lan- 
thanum, 350. Cerium, 351. Didymium, 351. Erbium, 351. 
Gallium Group, 351. Gallium, 352. Indium, 353. Thallium, 354. 
Thallous Compounds, 355. Thallic Compounds, 356. 
Tin and Lead, 356. 
Tin, 357. Stannous Compounds, 358. Stannic Compounds, 359. 
Lead, 361. Lead Peroxide, 362. 
Bismuth, 365. 
Chromium Group, 367. 
Chromium, 367. Cliromous Compounds, 369. Chromic Com- 
pounds, 369. Chromic Acid Compounds, 371. Chromium Oxy- 
chloride, 374. Molybdenum, 375. Tungsten, 377. Uranium, 
378. 
Manganese, 379. 
Forms of Combination, 380. Manganous Compounds, 381. Man- 
ganic Compounds, 381. Manganic Acid, 383. 
Metals of Group VIII, 385. 
Iron Group, 386. Iron, 387. Metallurgy of Iron, 388. Ferrous 
Compounds, 391. Ferric Compounds, 392. Ferric Acid Com- 
pounds, 394. Cyanogen Compounds, 394. Cobalt, 397. Cobalt- 
amine Compounds, 399. Cobalt-cyanogen Compounds, 399. 
Nickel, 400. 
Platinum Metals, 401. Ruthenium and Osmium, 403. Rhodium 
and Iridium, 404. Palladium, 405. Platinum, 406. 
Spectrum Analysis, 409. 
Periodicity of Spectra Lines, 415. A TEXT-BOOK 
<JF 
INORGANIC CHEMISTRY. 
INTRODUCTION. 
Upon contemplating Nature, surrounding us on all sides, 
we at once discover an endless multitude of objects or bodies. 
That which forms the basis of the latter, strongly character- 
ized by extent and weight, we designate substance or matter. 
lhe investigation of the iuternal and external structure of 
bodies, their classification according to conformable or distin- 
guishing characteristics, constitute the task of the descriptive 
sciences; of mineralogy, of geology, of descriptive botany and 
zoology, of anatomy, etc., etc. 
A closer scrutiny of natural objects discloses the fact that 
they in time succumb to many more or less serious alterations 
or changes. We observe that minerals are formed, crystallize, 
then disintegrate and crumble to pieces ; that plants and 
animals spring up, grow, and then fall into decay and 
decomposition. Such changes in the condition of bodies 
occurring with time are entitled phenomena. The investiga- 
tion of these during their progression, the determination of 
the laws according to which they occur, the explanation of the 
causes underlying them, form the task of the speculative 
sciences, physics and chemistry—depending upon the nature 
of the phenomena.. 
Although nature does not present any abrupt boun- 
daries, but throughout gradual transitions and intermedi- 
ate steps, we yet can distinguish with tolerable distinct- 
ness two different classes of phenomena. Some changes 
in the condition of bodies are only superficial (external), and 
are not accompanied by material alteration in substance. 
Thus heat converts water into steam, which upon subsequent 10 
INORGANIC CHEMISTRY. 
cooling is again condensed to water, and at lower temperatures 
becomes ice. In these three conditions, the solid, liquid, and 
gaseous, the substance or the matter of water or ice is un- 
changed ; only the separation and the motion of the smallest 
particles—their state of aggregation—is a different one. 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 sus- 
pended vertically 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 inspects, in their external 
and internal form or condition, they have suffered no percep- 
tible alteration; the glass is glass, and the iron remains iron. 
All such changes in the condition of bodies, unaccompanied 
by any real alteration in substance, are known as physical phe- 
nomena. 
Let us turn our attention now to the consideration of another 
class of phenomena. It is well known that ordinary iron under- 
goes a change, which we term rusting-; i. e., it is transformed into 
a brown substance which is entirely different from iron. On 
mixing finely divided copper filings with flowers of sulphur (pul- 
verulent sulphur) there results an apparently uniform, grayish- 
green powder. If this be examined, however, under a mag- 
nifying glass, we can very plainly distinguish the red metallic 
copper particles in it from the yellow of sulphur; by treat- 
ment with water, the specifically lighter sulphur particles can 
easily be separated from those of the copper. From all this 
the powder above will represent nothing more than a mechan- 
ical mixture. If, however, this mixture be heated, e. g., in a 
glass test tube, it will commence to glow, and on cooling, a 
black, fused mass remains, which in all respects differs from 
copper and sulphur, and even under the strongest microscope 
does not reveal the slightest trace of the latter, and elutria- 
tion with water fails to effect a separation of the ingredients. 
By the mutual action of sulphur and copper in presence of 
heat, a new body with entirely different properties has arisen, 
and is named copper sulphide. A mixture of sulphur and 
iron or other metals deports itself in a similar manner; the 
resulting bodies are termed sulphides. 
Such mutual action of different bodies occurs not only under 
the influence of heat, but frequently at ordinary temperatures. INTRODUCTION. 
11 
If, e. g., mercury and sulphur are rubbed continuously 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 
autimony 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 complete, entire alteration 
of bodies entering the reaction follow, are termed chemical 
phenomena. Chemistry, then, is that department of natural 
science ivhich occupies itself with the study of those phenomena 
in which an alteration of substance has occurred. 
In the previously described experiments we observed the 
phenomena of chemical combination; from two different bodies 
arose new homogeneous ones. The opposite exist, consisting 
in the decomposition of compound bodies into two or more 
dissimilar 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 mer- 
cury. From this we observe that on heating solid red mer- 
curic oxide two different bodies arise: gaseous oxygen and 
liquid mercury. We conclude, then, that mercuric oxide 
holds in it, or consists of, two constituents—oxygen and mer- 
cury. This conclusion, arrived at by decomposition, by 
analysis, may readily be verified inversely by combination, 
by synthesis. It is only necessary to. heat mercury at a some- 
what lowrer temperature than in the preceding experiment, for 
some time, in an atmosphere 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 another body is required, which will com- 
bine 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, 12 
INORGANIC CHEMISTRY. 
bodies or substances are finally reached which have withstood 
all attempts to bring about their division into further con- 
stituents, and which cannot be formed by the union of others. 
Such substances are the chemical elements, which also constitute 
the limit of chemical change. Their number, at present, is 
about 67; some have been only recently discovered. To them 
belong all the metals, like iron, copper, lead, silver and gold. 
Other elements do not possess a metallic appearance, and are 
known as metalloids; i. e., they are substances resembling (as 
regards their further indivisibility) the metals. It would be 
more correct to term them non-metals. To these belong 
sulphur, carbon, phosphorus, oxygen, etc. The line between 
metals and non-metals is not very marked. 
By the union of the elements with each other in smaller or 
larger number and quantity, result the compound bodies 
known to us. Water is a compound of two giiseous elements 
—hydrogen and oxygen ; ordinary cooking salt consists of 
the metal sodium and the gaseous chlorine. These elements 
make up not only our own earth, but the heavenly bodies are 
composed of them, at least, so far as has been proven by spec- 
trum analysis. 
If the quantities by weight of substances entering into a 
chemical change be determined, we will discover that in all 
transpositions, in the decomposition of a compound into its 
constituents, and in the union of the elements to form com- 
pound bodies, loss in weight never occurs. The weight of the 
resulting compounds is invariably equal to the sum of the weights 
of the bodies entering the reaction. Well known, general phe- 
nomena apparently contradict this scientific conclusion. We 
observe plants springing from a small germ and constantly 
acquiring weight and volume. This spontaneous increase of 
substance, however, is only seeming. Closer inspection proves 
conclusively that the growth of plants occurs only in conse- 
quence of the absorption of substance from the earth and 
atmosphere. The opposite 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 
transformation of visible solid or liquid bodies into non-visible 
gases. Carbon and hydrogen, the usual constituents ot 
The Principle.of the Indestructibility of Matter. INTRODUCTION. 
combustible substances, e. g., a candle, combine in their 
combustion with the oxygen of the air and yield gaseous 
products—the so-called carbon dioxide and water, which 
diffuse themselves in the atmosphere. If these products be 
collected, 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. 
From what has been remarked we can conclude that in 
chemical transpositions loss in matter does not occur, nor is 
there a new creation of the same observed. Compounds are 
formed and disappear, because they are converted into new 
forms, but their substance (matter), their weight, does not 
disappear, and is not produced anew. This fundamental truth 
is called the principle of the indestructibility of substance (mat- 
ter). Lavoisier, in the eighteenth century, first determined it 
by convincing experiments, and in combination with the 
principle of the conservation of energy, it constitutes the firm 
foundation of all scientific knowledge. 
The Principle of the Conservation of Energy—Chemical Energy. 
Causes underlie and influence all material phenomena. 
The final cause of phenomena we term force, accepting for the 
various sorts of phenomena a variety of forces. Some of these 
are attraction and repulsion, light, heat, electricity, cohesion, 
chemical affinity, and others. These names, however, only 
represent kinds of phenomena, without explaining their true 
nature. The positive knowledge we possess of the nature of 
some forces is that they consist of modes of motion. Thus, the 
phenomena of light are explained by the vibrations of ethereal 
particles which penetrate all substances; and those of heat are 
dependent upon the motions of the smallest particles of mat- 
ter. Accurate physical investigations have established that 
the different forces or causes of motion can never be destroyed, 
but merely changed from one kind to another. The move- 
ments or vibrations of one variety pass into those of another. 
For example, a discharged bullet is heated by coming in con- 
tact with any obstruction in its course; the visible motion of 
the entire mass in this instance is transformed into the invisi- 
ble motions of the smallest particles, and appears as heat. The 
heat motions can, on the other hand, be again changed into 14 
INORGANIC CnEMISTRY. 
mechanical force (molecular motion), or into light, magnetism 
or electricity. 
In all these transformations of the different forces into each 
other, we observe a perfect equivalence of their quantity ; 
if a mechanical force can produce a certain degree of heat, so 
can, vice versa, the latter perform the same mechanical work 
(the mechanical equivalent of heat, light, electricity). Upon 
this equivalence o f force rests the principle of the conservation 
of energy, according to which the various forces or motions of 
matter can neither be annihilated nor produced anew. This 
principle, forming one of the most important corner stones 
of natural science, was first sharply defined by the specula- 
tive observations of Dr. J. R. Mayer, of Heiibronn, in 
1842, and since then has been repeatedly confirmed experi- 
mentally. 
The most recent advance in physics has led to the negation of the 
objective existence of all abstract physical forces. Not considering 
the phenomena of electricity and those of chemical affinity—the re- 
duction of which to forms of motion is clearly foreseen, and not to be 
doubted—the ouly remaining, as yet, enigmatical force is that of attrac- 
tion or gravity. To affirm the existence of gravity is nothing more 
than to give expression to the fact that bodies in space tend to 
approach each other. The supposition that the active cause of 
gravity existed within the bodies themselves was long ago discarded 
by Newton as “ absurdum; ” it is merely a mathematical fiction. 
The action of a body in a place where it does not exist, without the 
aid of a medium, is inconceivable. The transference of the gravita- 
tion into material bodies, further, contradicts the principle of conser- 
vation of energy, as gravity is neither transferred nor exhausted—be 
it through the approach of bodies, whereby the force always increases; 
be it by planet movement, in whieh the centrifugal component is con- 
stantly overcome. Therefore, the active cause of gravity is not to be 
sought after in bodies themselves, but without them, and, indeed, in a 
substantial medium—ether—without the acceptance of which natural 
investigation cannot proceed. 
If we desire to make a preliminary presentation upon these rela- 
tions, the following would be the simplest and most probable: Space 
is filled by the smallest possible material particles, but as they are all 
alike, they do not possess gravity, and are found in constant transfer- 
able motion—ether substance. Bv the congress of the smallest ether 
particles to mass aggregates arise the chemical elementary atoms, 
which constitute material bodies—substance or matter. If, now, in 
addition to this one mass aggregate, a second appear in space, an 
effort to approach each other, produced by the action (collision) of the 
disturbed ether surrounding them, will appear; they possess gravity. 
By these suppositions the obscure ideas upon potential energy and INTRODUCTION. 
15 
energy of place are removed. A clearer and more thoroughly distinct 
presentation and confirmation of these representations, and especially 
upon the nature of forces, are found in A. Secchi’s “ Die Einheit der 
Naturkrafte.” 
In the chemical union of bodies heat is almost invariably 
disengaged, and as it is a variety of motion, and as motion of 
one kind can only be derived from another, we must conclude 
that bodies acting chemically, especially the elements, do pos- 
sess a peculiar sort of motion, which, in chemical union, is 
partially converted into heat motion (also into light and elec- 
tricity). This special motion of matter is designated chemical 
energy or chemical tension. And in the chemical decomposi- 
tion of a compound body into its constituents, heat is 
absorbed, disappears as such, and is transformed into chemical 
energy. Thus, for instance, in the union of 1 kilogram of 
hydrogen with 8 kilograms of oxygen a quantity of heat is 
liberated which can perform a mechanical work equal to 
34,462 x 423.5 = 14,629,000 kilos; in the decomposition, on 
the other hand, of 9 kilos of water into hydrogen and oxygen, 
the same force or quantity of heat is necessary; therefore, in 
the liberated hydrogen and oxygen, the same quantity of force 
or motion must be contained in the form of chemical energy. 
Chemical energy is not only a quantitative phenomenon ; it 
also presents qualitative differences. Although all bodies, 
and particularly the elements, possess it, they do not disclose it 
in the same way in their action upon each other. Some 
unite or react readily with each other; others, on the contrary, 
only with difficulty, or not at all. The reason of this deport- 
ment is to us entirely unknown. We designate it with the 
phrase chemical affinity, and add that bodies capable of union 
have affinity for each other (are related), and that by union 
they satisfy their affinity. This expression is incorrectly 
chosen, because, generally, the bodies least alike chemically 
unite with each other most readily. 
Conditions of Chemical Action. 
The main condition in the chemical action of bodies con- 
sists in their immediate contact. Whilst, in physical changes, 
the bodies act at smaller or greater distances upon each other 
(by aid of the ethereal fluid), in chemical alterations they act 
at immeasurably small distances. The indispensable intimate 16 
INORGANIC CHEMISTRY. 
contact of solid bodies is difficult to attain by mechanical 
mixture. Thus, no action results when we rub dry tartaric 
acid and soda together in a mortar, but as soon as some water 
is added to the mixture chemical action at once sets in, 
accompanied with effervescence. As shown in this case, the 
requisite intimate contact is generally obtained by bringing 
both or one of the constituents into liquid form, either by 
solution in a solvent or by fusion. The early chemists expressed 
these conditions by the sentence: Corpora non agunt nisi 
fluido. Liquid and gaseous bodies are, therefore, eo ipso, 
adapted to chemical action. 
Besides intimate contact, a definite temperature and other 
conditions are necessary. Thus in the preparation of gun- 
powder, a mixture of carbon, sulphur and nitre, moistened 
with water, is strongly rubbed and worked through rollers 
without any chemical action occurring; but if a portion of 
the mixture be heated to 300° C., or an electric spark be 
conducted through it, chemical action with explosion immedi- 
ately takes place. The explanation of these phenomena is based 
on the fact that the very powerful cohesion of similar particles 
in solid bodies opposes the chemical affinity of the dissimilar 
particles mixed in with each other. As heat works expansively 
and diminishes the cohesion of the bodies, it is easily compre- 
hended that with the increase of temperature the chemical 
attraction of the various substances is in condition to overcome 
the cohesive force. Light and electricity behave like heat. 
Again, on the opposite hand, the chemical attraction between 
the different constituents of a compound body is sometimes so 
weak that the heat vibrations increasing with the temperature 
outweigh the chemical attraction, so that at a definite temper- 
ature decomposition or alteration of the compound body will 
result. At very high temperatures almost all compounds 
decompose into their constituents (compare Dissociation of 
Water). Heat is, therefore, an important agent in chemical 
changes, because it induces both the phenomena of combination 
and decomposition. Light and electricity do the same. 
In addition to these circumstances we should remember that 
both the compound bodies and elements consist of several 
particles (atoms) combined with each other, and that this 
union must first be broken to render possible the action of 
the other particles. INTRODUCTION. 
17 
Consequent upon the mutual action of all these influences 
follow the complicated phenomena observed in chemical 
changes. The qualitative phenomena of chemical energy are 
extremely different, and the determination of the laws con- 
trolling them remains for the future; in special instances 
reference will be made to individual generalizations. The 
quantitative relations, on the other hand, in the action of 
bodies on each other, in the combination of the elements, are 
thoroughly investigated, and lead to very important laws, 
speculations and theories, which constitute the main portion 
of the present scientific chemistry. 
As briefly represented in the preceding, two different factors are to 
be considered in every chemical phenomenon. First is the material 
side, which finds expression in the solid weight proportions of the 
reacting and resulting bodies ; secondly, a dynamical event presents 
itself. When hydrogen and oxygen chemically unite to produce 
water, there occurs, in addition, a considerable development of heat, by 
which the chemical energy and affinity become evident (p. 57). Although 
the physical cause of the latter, which formerly was identified with 
electrical differences, is as yet unexplained, the investigations of recent 
date have disclosed a close relation between chemical affinity and the 
thermal phenomena: the greater the liberation of heat in the union of 
two elements or bodies, the greater also is their affinity, so that the 
former may serve as an approximate measure of the latter. The 
general proposition follows, from this, that the qualitative alterations 
occurring in the reaction of bodies invariably pursue the direction in 
which the most heat is disengaged. This fundamental proposition of 
the greatest development of heat (Bertholet) is a special case of 
Clausius’ proposition of the entropy of energy, according to which 
every form of energy tends to pass into heat motion. A closer deduc- 
tion of this, for chemical changes, so important a proposition, will be 
given later in the most important special cases. 
For simplicity and convenience, the elements are repre- 
sented by the first letters of their names, derived either from 
the Latin or Greek. Hydrogen is represented hy the letter 
H, from the word hydrogenium ; nitrogen by N, from nitro- 
genium. When several elements happen to have the same 
letter there is added to the capital a second, small letter ; 
thus Na represents natrium; Ni, nickel; Hg, mercury 
(hydrargyrum), etc., etc. The subjoined table comprises all 
the elements at present known (67), together with their 
chemical symbols and atomic weights :— 
Chemical Symbols and Formulas 18 
INORGANIC CHEMISTRY. 
ELEMENTS. 
Symbol. 
Atomic 
Weight. 
ELEMENTS. 
| Symbol. 
£ 
IS 
Aluminium 
A1 
27.4 
Hg 
200 
Antimony (Stibium) 
Sb 
122 
Molybdenum. A.t.V:. 
Mo 
96 
Arsenic 
75 
N 
14 
Barium 
137 
Ni 
59 
Beryllium 
Be 
9.4 
Nb 
94 
Bismuth 
Bi 
210 
Os 
199 
Boron 
B 
11 
0 
16 
Bromine 
Br 
80 
Palladium R.T.v'f. 
Pd 
106.6 
Cadmium 
Cd 
112 
Phosphorus..'S.t.A. 
P 
31 
Caesium 
Cs 
133 
Pt 
197 
Calcium 
Ca 
40 
Rh 
101.4 
Carbon. - 
c 
12 
Rb 
85.4 
Cerium 
Ce 
92 (?) 
Sc 
44 
Chromium 
Cr 
52.5 
Sulphur.. 
s 
32 
Copper 
Cu 
63.5 
Se 
79.4 
Chlorine 
Cl 
35.4 
Silver (Argentum) I 
Ag 
108 
Cobalt 
Co 
59 
Si 
28 
Didymium 
Di 
95 (?) 
Sodium (Natrium) / 
Na 
23 
Erbium 
Er 
112.6 (?) 
Sr 
87 
Fluorine 
FI 
19 
Tantal .vT... 
Ta 
182 
Ga 
69.8 
Te 
128 
Au 
197 
Thl 
204 
Hydrogen 
II 
1 
Thorium.... 
Th 
231 
Indium 
In 
113.4 
Tin (Stannum )?-.7.7r 
Sn 
118 
Iodine 
I 
127 
Ti 
50 
Iridium 
Ir 
198 
Tungsten( Wolfram) 
4W 
184 
Iron (Ferrum) 
Fe 
56 
Uranium ?n... 
Ur 
240 
Kalium (Potassium) 
K 
39 
Vanadium 
Vd 
51.3 
Lanthanum 
92 (?) 
Yb 
113 
Lithium 
Li 
7 
Yttrium .'5... 
Y 
61.7 (?) 
Lead (Plumbum)... 
Pb 
210 
Zinc ?r.. 
Zn 
65.2 
Magnesium 
Mg 
24 
Zr 
89.6 
Manganese. 
Mn 
55 
The existence of some other elements, as Dianium, Jargonium, 
Terbium, Philippium, Wasinm, Vesbium, Norwegium and Thulium, is 
yet doubtful. 
Compounds produced by the union of the elements are rep- 
resented by placing their corresponding symbols together and 
designating these chemical formulas. Cooking salt, a compound 
of sodium and chlorine, is represented by the formula NaCI; 
mercury oxide, a compound of mercury and oxygen, by HgO ; 
iron sulphide by FeS; hypochlorous acid, a compound of 
chlorine, hydrogen and oxygen, by CIOH. INTRODUCTION. 
19 
Chemical formulas not only express the nature of the ele- 
ments, but also contain relative proportions of their weights, 
according to which they unite, and which are compared with 
hydrogen as unity. Thus H represents 1 pt. by weight of 
hydrogen; Cl, 35.4 pts. by weight of chlorine; Na, 23 pts. by 
weight of sodium. Then the formula NaCI would represent 
the union of 25 parts of sodium with 35.4 parts of chlorine, 
Na, 23 parts by weight of sodium. 
These numbers indicate the relative weights of the atoms 
constituting the elements. 
If we seek to obtain a representation of the constitution of 
the elements and substances in general, but two possibilities 
appear to be foremost. Either the substance continuously 
fills space, or it consists of very small, separated particles 
filling space, chemical individuals, which are termed atoms. 
The latter idea alone corresponds to the present view of 
physical and chemical investigation, so that the atomic consti- 
tution of substance alone has at present any value. The in- 
ductive derivation and establishment of the atomic theory will 
be given subsequently (see page 59) ; here only the following 
propositions are laid down. The elements consist of similar 
atoms, of like size and similar weight, while atoms of different 
elements possess a different weight. The absolute atomic 
weights are, at present, not determined with sufficient ac- 
curacy; the relative weights are referred to the hydrogen 
atom, which has the smallest weight, hence is made equal to 
1 (H = 1). By the chemical union of the atoms are formed 
the smallest particles of compound bodies, termed molecules, 
physical individuals; these are chemically divisible. By 
these premises chemical formulas acquire a very precise and 
intuitive importance. The formula NaCl designates the union 
of 1 atom of sodium (Na) with 1 atom of chlorine, and indi- 
cates that in it 23 parts, by weight, of sodium are combined 
with 35.4 parts of chlorine. If in a compound several atoms 
of an element are present, this is denoted by numbers attached 
to the symbol of the atom :— 
HC1 
Hydrochloric acid. 
H20 
Water. 
h3n 
Ammonia. 
CH4 
Methane. 
The formula of water (H20) means that its molecule consists 
of two atoms of hydrogen (2 parts by weight) and 1 atom of 
oxygen (O = 16 parts by weight.) The formula of sulphuric 20 
INORGANIC CHEMISTRY. 
acid (H2S04) indicates it as a compound consisting of 1 atom 
of sulphur (32 parts), 4 atoms oxygen (4 x 16 — 64 parts), and 
2 atoms of hydrogen (2X1 = 2 parts), from which the com- 
position of the acid may be at once calculated into per cent., 
or in any desirable quantity by weight. 
Atomic Composition. 
In per cent. 
Sulphur, S — 32 
82.65 
Oxygen, 0* — 64 
Hydrogen, H2 — 2 
65.30 
2.05 
H2S04=98 
100.00 
When the elements enter into combination in several pro- 
portions by weight, this is indicated by placing figures to the 
right and below the symbol, e. g.:— 
S04 H2. Cl 03 H. 
The chemical union of bodies is shown by the sign and 
the resulting products are placed to the right, following the = 
sign :— 
By this equation of chemical transposition is meant that by 
the union of mercury sulphide (HgS) and iron (Fe), iron sul- 
phide (FeS) and tree mercury (Hg) are formed. At the 
same time such equations show the proportions by weight of 
the substances entering into and resulting from the reac- 
tion ; the weight of the acting substances is equal to that of 
those resulting. Therefore every chemical equation is at 
once an expression of the principle of the indestructibility of 
matter (substance). (See p. 13.) 
HgS + Fe = FeS + Hg. 
Crystallography. 
Chemistry only occupies itself with the study of the 
chemical alterations of bodies. Its subject is not the latter 
for themselves, in their external properties, but only according 
to their material composition, and their genetic relations to 
other substances. The investigation of the physical properties 
of the non-organized bodies constitutes the province of 
Mineralogy, or, if the same is not limited to naturally occur- 
ring bodies, but also includes the numberless artificially pre- 
pared substances, it becomes the proviuce of Inorganography. 
Pure chemistry considers the physical properties only in so far INTRODUCTION. 
21 
as they serve for the external characterization and eventual 
recognition of the given substances and for the deduction of 
chemical regularities. The most important physical proper- 
ties—state of aggregation, the temperature of fusion and 
boiling, the specific gravity, capacity for heat, etc.—are partly 
obtained from Physics, and, in part, will be considered later, 
in special cases. 
Here, therefore, only the morphological characters of the 
solid bodies will receiVe a brief consideration. 
The homogeneous solids exhibit either similar properties in 
all their parts, are amorphous, or show differences in certain 
definite directions, which permit them to appear as crystalline. 
The cause of this deportment lies in the arrangement of the 
smallest substance particles, of the molecules, which in the first 
instance is irregular, hence cannot cause differences in any 
direction, while in the crystalline structure the molecules are 
regularly grouped according to directions of varying density 
and coherence, which find expression in the cleavage and the 
optical and thermal behavior of bodies. A consequence of 
this regular arrangement is, in the case of undisturbed forma- 
tion, the external limitation of bodies by planes, edges and 
angles, which represent the crystal form. The number and 
forms of these crystal elements are very much multiplied, since 
several thousand forms are known. It is, however, possible 
to reduce the numberless varieties to a few classes or systems, 
by comparing the same in their manner of formation, and 
by referring their principal elements—the planes—to definite 
axes, i. e., to directions or lines, which are imagined so placed 
through the middle point of the crystals that their planes lie 
symmetrically with them. In this manner, by distinguishing 
the various axis intersection, crystallography arrives at the 
following six systems of crystallization:— 
1. The Regular or Tesseral System. 
2. Quadratic or Tetragonal System. 
3. Rhombic System. 
4. Hexagonal or Rhombohedral System. 
5. Monoclinic System. 
6. Triclinic System. 
The position of any even plane in space is determined by 
three points of a system of coordinates, and also the position 
of a crystal plane by its points of intersection with the 
three axes, or by the distances from the centre of the axis 22 
INORGANIC CHEMISTRY. 
at which the plane (by suitable expansion) cuts or inter- 
sects the three axes. These distances are termed the parame- 
ters of the plane. In the regular system all three axes are of 
equal length and equal value, for which reason they are 
designated by the same letter a (Fig. 1). If a plane intersect 
all three axes at equal distances (as in the octahedron) the 
parameter ratio is a : a : a; if at unequal distances the ratio is 
a : a : ma,or a : ma : na; in which the parameter upon the first 
axis is also made = l. If a plane lie parallel to an axis (in- 
tersecting it at infinity) its parameter with reference to this 
axis = infinity (a : a : ao a); if it be parallel to two axes, two 
parameters are = infinity (a : ao a : oo a). Hence these para- 
meter ratios designate the position of a plane; and as all the 
planes in simple crystal forms are similar, these marks repre- 
sent also the entire simple form, i. e., the whole of all planes, 
which are deduced from the same parameter ratio. 
In the regular system, when all these axes are of equal 
length, the parameters m, n (the parameter of the first axis 
made = 1) are, in accordance with experience in all crystals, 
simple rational numbers (1, f, 2, f, 3, 4, etc.) In other 
systems of crystallization the axes are unequal and do not 
stand in any rational ratio to their lengths (a:b:c), and, 
therefore, also to the parameters of the planes (a : mb : nc) ; 
but the ratio of the parameter coefficients 1. m. n., in them, 
as in the regular system, is, as learned from experience, a 
simple, rational one. 
According to the number of axes, five of the crystallographic systems 
are triaxial or trimetric; for practical reasons we accept four axes in 
the hexagonal system—tetrametric system. The axes of the five 
trimetric systems are either all at right angles to each other, and are 
distinguished only by varying lengths—the orthometric systems, the 
regular, the quadratic and the rhombic ; or they include also oblique 
angles—in the two dinometric systems, the monoclinic and tri- 
clinic. In the regular system the three directions of development 
are at right angles and perfectly similar, for which reason the bodies 
belonging to it, like amorphous substances, refract light simply— 
isometric system. The remaining five systems with unlike axes are, 
on the contrary, doubly refracting, and, indeed, they either only 
exhibit simple refraction in one direction (optically uniaxial systems 
—the quadratic and hexagonal), or truly in two directions (optically, 
diaxial systems—the rhombic, mono- and triclinic.) A necessary, 
geometrical derivation of the various systems of crystallization, based 
upon the molecular constitution of matter, has been recently proposed 
by Bravais and Sohncke, in that they proceed from the study of pos- 
sible normal positions of points or molecules in space ; i. e., of such INTRODUCTION. 
23 
positions in which about each point the grouping of the remaining 
points is the same. Thus, it is possible to embrace the point position 
possible in space into seven groups, which correspond to the six 
crystallographic systems; in this manner the rhombohedral forms, 
which crystallography views as the hemihedral of the hexagonal, appear 
as an independent system. 
1. Regular System. The crystals of this system are 
similarly developed according to three directions, and therefore 
have three similar axes, a, of the same length, which intersect 
each other at right angles (Fig. 1). If we imagine the planes 
arranged in such manner that they intersect the three axes at 
Fig. 1 
Fig. 2. 
Fig. 3. 
Fig. 4. 
Fig. 5. 
Fig. 6. 
similar or dissimilar distances, then, for the various possible 
parameter ratios (page 22) we have the following seven 
simple forms :— 
1. The octahedron (Fig. 2), with the parameter ratio a: a: 
a; for which the abbreviated symbol O may be written. 2. 
The cube or hexahedron (Fier. 3), with the symbol a: aoa: 24 
INORGANIC CHEMISTRY. 
co a or briefly ooOoo. 3. The rhombic dodecahedron, twelve 
sides, a: a: oo a or oo O (Fig. 4). 4. The trigonal tris-octahe- 
dron (Fig. 5), a : a: ma or mO; there are several of these, 
when m = §, 2 or 3. 5. The trapezohedron (Fig. 6) a : ma : 
ma or mOm, in which m = 2 (in garnet) or m = 3 (in 
ammonium chloride). 6. The tetrahexahedron (Fig. 7) a: 
ma: oo a or mO cc, in which m = f, 2 or 3. 7. The hexocta- 
hedron (Fig. 8) a: ma: na or mOn, when, e. g., in fluorite, m 
= 2 and n = 3. 
Fig. 7. 
Fig. 8. 
Fig. 9. 
These seven simple forms may appear combined on one 
crystal, and thus crystal combinations arise. Thus Fig. 9 
represents a combination of octahedron and cube (on alum) ; 
Fig. 10, a combination of octahedron, cube and dodecahedron 
(galenite). 
Fig. 10. 
Fig. 11. 
Fig. 12. 
In addition to these simple forms, appearing with their full 
number of planes, hence termed holohedral, others appear 
having only the half or fourth of the possible forms, 
hemihedral or tetartohedral forms. We may suppose them 
produced by the enlargement of the symmetrically distributed INTRODUCTION. 
25 
half number of faces of the holohehral form, and therefore 
we designate them with the corresponding holohedral symbol 
divided by 2. Such hemihedral forms are : The tetrahedron 
-2°- (Fig. 11), resulting from the octahedron, and the pen- 
tagonal dodecahedron m °2-x- (Fig. 12) derived from the 
tetrahexahedron. 
2. Quadratic System. Crystals of this system are developed 
at right angles in three directions, of which two are alike, 
the third unlike the others. Therefore 
they possess two similar secondary axes, 
a, and an unlike, longer or shorter, prin- 
cipal axis, c, (Fig. 13). The axis ratio, 
therefore, is a: a: c ; the ratio of a: c is, 
for every body of the system, definite but 
not rational (see page 22), e. g. for tin 1: 
0.3857. The principal (ground) form* 
of this system is the quadratic pyramid 
a: a: c, with the symbol P, an obtuse 
pyramid (Fig. 14) if e be smaller, or an 
acute pyramid if c is greater than a. 
From these pyramids of the first order, in which the axes pass 
through the angles, we distinguish pyramids of the second 
order, in which the two secondary axes pass through the 
Fig. 13. 
Fig. 14. 
Fig. 15. 
Fig. 16. 
middle of the edges. The planes of the latter are parallel to 
one secondary axis, therefore its symbol is a : ooa: c or P oo. 
Fig. 15 represents a combination of a pyramid of the first and 
second orders. 
* The (ground) principal form is that to which the remaining forms 
of the same mineral may be most readily referred. 26 
INORGANIC CHEMISTRY. 
In addition to the above pyramids, others may occur in the same 
crystal, which intersect the principal axis at the distance me. In this 
case m is also a simple rational number, e. g., 3, ‘2, etc. The 
symbol of such secondary pyramids of first order, is a : a : me — mP ; 
that of the second order, a : cc a : me = mP oo. The coefficient of the 
secondary axis is always written before P. 
If the pyramid planes intersect the principal axis at infinity, 
the quadratic prism results, and, indeed, a prism of the first 
order, with the symbol a : a : co c = P oo; and a prism of 
second order, a : go a : go c =oo P oo. Fig. 16 represents a com- 
bination of the quadratic prism with the quadratic pyramid, as 
seen on zircon and potassium phosphate. With the parameter 
ratio of the planes, a : ma : nc, the ditetragonal pyramid nPm 
results. Corresponding to it is the ditetragonal prism, a : ma : 
co c = oo Pm. 
Different hemihedral forms are possible in this system, of 
which the tetragonal sphenoid,> corresponding to the tetra- 
hedron, may be mentioned. Chalcopyrite, tin, prussiate of 
potassium and potassium phosphate crystallize in this system. 
3. Hexagonal System. The forms of this system, like those 
of the preceding, exhibit one peculiarly striking direction of 
development, and hence this is chosen as the direction of the 
principal axis c. They are distinguished from the four-sided 
forms of the quadratic system by their sixfold symmetry, which 
finds expression in their similar secondary axes a (Fig. 17), in- 
Fig. 17. 
Fig. 18. 
Fig. 19. 
tersecting each other at 60°. The principal axis is at right 
angles to these. The ratio of the basal axis is a : a (: a) c, and 
the ratio of a: c for every substance is definite, but not rational; 
e. g., in quartz, 1 : 1.100; calcite, 1 : 0.8543, etc. INTRODUCTION. 
27 
The (ground) fundamental form of the system is the hexa- 
gonal pyramid a : a : (qo a) : c = P, from which is derived the 
hexagonal prism a : a : (qo a): oo c = qo P ; and, indeed, as in 
the quadratic system, there are pyramids and prisms of first 
and second order—latter with the symbol a : 2a : (2a) : c = 
P2, and a : 2a : (2a) : oo c = 00 P2. Further, other pyramids 
intersecting the principal axis, at the distance, me, occur; their 
symbol is mP and mP2. Fig. 18 represents the combination 
of pyramid and prism found on apatite. With the common 
parameter ratio a : 11a : (ra) : me (in which the parameter of 
the third secondary axis, ra, as in all other hexagonal forms, 
is always given by the parameters of the two first secondary 
axes), results the dihexagonal pyramid, mPn, and the dihexa- 
gonal prism, qo Pn. 
From the pyramids mP are derived, as hemihedral forms, 
by enlargement of the alternating planes, the rhombohedra 
± mR (Fig. 19). Another important hemihedral form 
is the scalenohedron ± derived from the dihexagonal pyra- 
mid. Worthy of note is the fact that the hemihedral forms 
of the hexagonal system occur much more frequently in 
nature in numberless combinations (especially on calcite), 
whence sometimes they are treated as a separate system. 
4. Rhombic System. Three axes of unequal length a. b. c. 
Any one, as c, is chosen as principal axis, and of the secondary 
axes the shorter a is designated as the brachydiagonal, the 
Fig. 20. 
Fig. 21. 
Fig. 22. 
longer b as macrodiagonal (Fig. 20). The axis ratio a : b : c 
is definite for every substance, e. g., for sulphur, 0.811 : 
1.0899. 28 
INORGANIC CHEMISTRY. 
The principal forms of the system are the rhombic pyra- 
mid a : b : c = P (Fig 21), the rhombic prism a : b : oo c = 
oo P, and the domes—brachydiagonal co a : b : c = P oo and 
the macrodiagonal a:oob:c = Poo = consisting of two pairs 
of planes. Fig. 22 shows a combination of two pyramids with 
the brachydiagonal dome. Sulphur (native and that crystal- 
lized from carbon disulphide), potassium nitrate, aragonite 
and barite belong in this system. 
5. Monoclinic System. Three unequal 
axes, two at right angles to each other, 
the third, however, at right angles with 
one and oblique to the other (Fig. 23) ; 
a to b and a to c at right angles, c to b 
oblique with the angle /?. The crystals 
of this system are principally developed 
according to an axis oblique to another 
(c to b), and hence one of these is chosen 
as principal axis c. Of the two second- 
aries, one, b, is called the clinodiagonal, 
and a the orthodiagonal. The axis ratio 
is always definite for every substance; e. g., for ferrous 
sulphate 1.1704 : 1 : 1.5312 with the angle /? = 76° 33'. 
The monoclinic pyramid is principal form a : b : c = P, con- 
sisting of two hemipyramids, -f P and — P, (Fig. 24). It 
corresponds to the monoclinic prism oo P. Sulphur (fused), 
soda, borax, disodium phosphate (HNa2 PO* + 12 H20), 
Glauber’s salt, and orthoclase, crystallize in this system. 
Fig. 23. 
Fig. 24. 
Fig. 25. 
Fig. 26. 
6. Triclinic System. Three unequal axes all oblique to 
each other. The axis ratio a : b : c, and the three angles are 
definite for each substance. The forms of this system are very 
complicated, as the triclinic pyramid P must be considered as INTRODUCTION. 
consisting of four quarter pyramids, and the triclinic prism 
oo P (Fig. 25), of two hemiprisms. Potassium dichromate, 
albite, boracic acid, copper sulphate, crystallize in this system. 
Fig. 26 represents one of the common forms of the latter. 
The crystals in nature, however, because they have grown, 
rarely occur so regularly developed as represented in the 
preceding drawings. Ordinarily they are developed more or 
less in the direction of single axes, whereby the faces of the 
same form are unlike and the entire crystal appears distorted. 
But the position of the planes with reference to the axes and 
the angles which they form with each other always remain 
unchanged. Therefore the measurement of the angle of the 
planes by means of the goniometer serves as the only accurate 
means of determining complicated crystalline forms ; ffom the 
angles we calculate the axis ratio. 
Substances crystallizing in two or three different systems are 
said to be dimorphous, or trimorphous (see sulphur). Various 
substances which crystallize in the same or very similar forms 
are termed isomorphous (compare Isomorphism). SPECIAL PART. 
CLASSIFICATION OF THE ELEMENTS. 
Ordinarily we are accustomed to divide the elements into 
two groups: metals and metalloids (see p. 12). The former 
possess metallic appearance, are good conductors of heat and 
electricity; the latter or non-metals do not have these proper- 
ties, or at least in less degree. In chemical respects the metal- 
loids have the tendency to combine with hydrogen, forming 
volatile, generally gaseous, compounds: the oxygen derivatives 
form acids with water. The metals, on the contrary, rarely 
unite with hydrogen, and their oxygen derivatives yield the 
so-called bases with water. Further, the binary compounds of 
metals with the metalloids are so decomposed by the electric 
current that the metal separates at the electro-negative, and 
the metalloid 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 which in their external appear- 
ance appear as metals, while in a chemical respect they deport 
themselves throughout as metalloids, and vice versa. Thus 
hydrogen, a gaseous element, approaches the metals through- 
out in its chemical character, while metallic antimony arranges 
itself with the metalloids. 
It is therefore advisable to divide the elements, according to 
their chemical analogies, in separate, natural groups. The 
best and only correct classification 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 will treat 
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 and some metals, are the fol- 
lowing :— HYDROGEN 
Fluorine. 
Chlorine. 
Bromine. 
Iodine. 
Oxygen. 
Sulphur. 
Selenium. 
Tellurium, 
Nitrogen. 
Phosphorus. 
Arsenic. 
Antimony. 
Bismuth. 
Carbon. 
Silicon. 
Tin. 
Hydrogen does not belong to either of the above; uniting 
the metal and non-metallic character in itself, it represents, as 
it were, the type of all elements, and therefore it will receive 
first attention. Boron, which has also been classed with the 
metalloids, although differing somewhat from them in chemical 
deportment, occupies an isolated position. 
HYDROGEN. 
H =1 
II2 = 2. 
Hydrogen (Hydrogenium), a gaseous body, occurs only ex- 
ceptionally in a free condition upon the earth’s surface, as it com- 
bines readily with the oxygen of the air. Found in considerable 
quantity in the photosphere of the sun and fixed stars. In 
combination, we find it principally as water and in substances 
of vegetable and animal origin. Paracelsus first discovered 
this element in the sixteenth century, and called it inflamma- 
ble air. In 1781 
Watts and Caven- 
dish showed that 
water resulted 
from the combus- 
tion of hydrogen 
in the air. 
Preparation. It 
may be readily 
obtained from 
water, a com- 
pound of hydro- 
gen and oxygen. 
The decomposi- 
tion of the same 
by the removal 
of oxygen can be 
effected by some 
metals, like Na 
and K, at the 
Fig. 27. 32 
INORGANIC CHEMISTRY. 
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. 27). Bubbles of 
hydrogen are at once disengaged, which displace the water 
and collect in the cylinder. The reaction occurring between 
the sodium and water is expressed by the following chemical 
equation:— 
H20 + Na = NaOH + H. 
Water. 
Sodium. 
Hydrogen 
The compound NaOH, known as sodium hydrate, remains 
dissolved in the excess of water. 
Other metals decompose water in a similar manner, at an 
elevated temperature. To effect this with iron allow steam to 
pass through a tube filled with iron filings, which are exposed 
to a red heat in a combustion furnace. The iron withdraws 
oxygen from the water, combining with it, while the hydrogen 
set free is collected. 
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 + H2S04 = ZnS04 + 2H. 
Sulphuric 
acid. 
Zinc 
sulphate. 
Fig. 28. HYDROGEN. 
Place granulated zinc (obtained by dropping molten zinc 
into water) in a double-necked tlask (Fig. 28), and introduce 
sulphuric acid through the funnel tube, b (diluted with about 
3 vols. H20). The liberation of gas begins immediately, and 
the hydrogen, escaping through the exit tube,/, is collected 
as previously described. 
The hydrogen thus formed, in consequence of a slight 
admixture of foreign substances, has a faint odor. Pure 
hydrogen may be obtained by heating potassium formate with 
potassium hydrate:— 
CH02K + KOH = K2C03 + 2H. 
Purifying and Drying of Gases. To free gases of the substances 
which, during their disengagement, might have mechanically been 
carried along, it is best to conduct them through variously constructed 
wash bottles, tilled with water or such liquids which will absorb the 
impurities. Ordinarily the so-called Woulff’s three-necked bottles are 
employed (compare Figs 36 and 42). The open tube, placed in the 
middle tubulure, is called the safety tube. It serves to equalize the 
inner pressure with that of the external atmosphere. 
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 hygroscopic sub- 
stances (see Fig. 34). For this pur- 
pose calcium chloride, potassium 
hydrate, sulphuric acid and other 
reagents may be used. 
Apparatus for the Generation 
and Collection of Gases. -In the 
apparatus pictured in Fig. 28, the 
liberation of hydrogen continues 
uninterruptedly as long as zinc and 
sulphuric acid are present. To be 
able to control the generation of the 
gas at pleasure, different forms of 
apparatus are resorted to. One of 
the most practicable of these is that 
of Kipp. It consists of two glass 
spheres d and b, Fig. 29, in the upper 
opening of which, c, a third sphere, 
with an elongated tube, fits air-tight 
and serves as a funnel. In the mid- 
dle sphere there is placed granu- 
lated zinc, through the tubulure e, 
and diluted sulphuric acid is poured 
in the sphericaj funnel, which first 
fills d, then ascends to b, where it 
comes in contact with the metal; 
at once the evolution of hydrogen 
Fig. 29. 34 
commences and the gas escapes through e. Upon closing the stop- 
cock of the tube fixed in e, the hydrogen that is being set free presses 
the sulphuric acid out of 6, and consequently the liberation of gas 
ceases. On again opening the cock, the acid rises in b to the zinc, and 
the evolution of gas commences anew. The vessel a contains water to 
wash the escaping hydrogen 
The somewhat complicated Kipp apparatus may be advantageously 
replaced by the following simple contrivance (Fig. 30). Two bottles 
provided with openings near their bottom, in which are glass tubes, 
INORGANIC CHEMISTRY. 
Fig. 30. 
are connected by a gum tube. The bottle A is filled with granulated 
zinc, and B with dilute sulphuric acid. The cock It closes A. When 
this is opened, the sulphuric acid flows from B to A, to the zinc, and 
the evolution of gas commences. On closing the stop-cock the hydro- 
gen presses the acid back, from A to B; the evolution of gas ceases. By 
elevating and sinking the flasks the regulation can be hastened. 
Gasometers of various construction serve to collect and preserve 
gases. In (Fig. 31) we have the ordinary gasometer of Pepys. It is 
constructed from 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 b. The tubes c and c' are only 
supports. To collect gases in this apparatus it must first be filled with 
water. To this end, pour water into the upper cylinder, and open a 
and e; the water then flows through a, nearly reaching the bottom of 
the lower cylinder, while the air escapes through e. The side glass 
tubes, b, allow the operator to observe the height of fhe water level. 
When the lower cylinder is filled with water, close e and a (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 place in the HYDROGEN. 
same 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 a, after filling the upper cylin- 
der, and then, if desired, open a, and the gas can be set free, either by 
e or b. 
In addition to the gaso- 
meter described, various 
other forms are employed ; 
gas bags are very well 
adapted for preserving 
gases. 
Physical Properties. 
Hydrogen is a colorless, 
odorless and tasteless 
gas. Its ability to re- 
fract light and conduct 
heat is, in accordance 
with the metallic nature 
of hydrogen, greater 
than that of all other 
gases. By cooling 
(—140° C.) and power- 
ful pressure (600 atmos- 
pheres), it is condensed 
to a steel-blue, non- 
transparent liquid, 
which upon evaporation 
even becomes solid; 
consequently liquid 
hydrogen resembles a 
molten metal, or, at 
ordinary temperatures, 
liquid mercury. 
Fig. 31. 
Condensation of Gases. Hydrogen and several other gases (oxygen, 
nitrogen, carbon oxide, nitrogen oxide, methyl), until recently, were 
considered non-coercible gases (permanent gases), because they could 
not even be condensed by the strongest pressure (over 2000 atmos- 
pheres). This supposed non-condensability was due to the fact that a 
general property of gases, called by Andrews their critical temperature, 
had been overlooked. All gases show, indeed, as first observed with 
carbon dioxide (see this), a definite temperature beyond which they 
cannot be condensed by any pressure. Conversely, all liquids, by the 
same corresponding temperatures, even under the greatest pressure, 
are transformed into gases (absolute boiling temperature of Mendele- 
jeff). Consequently, in the condensation of gases, in addition to 
pressure, the corresponding lower temperature must be applied. In 36 
INORGANIC CHEMISTRY. 
fact, by this means, of late years, all the so called permanent gases 
have been condensed (Pictet and Cailletet). The lower temperatures 
are attained by the rapid evaporation of liquid carbon dioxide (—130°), 
or hyponitrous oxide (— 140°) with aid of air pumps ; or the strongly 
compressed gas is permitted to expand rapidly, whereby, through the 
evaporation, so much heat is absorbed that the remainder of the gas 
condenses. 
Its refraction of light and conductivity of heat is greater 
than that of other gases. Like all difficultly coercible gases, 
hydrogen is but slightly soluble in water, 100 volumes dis- 
solving 1.9 volumes H. It is the lightest of all gases, being 
14.46 times lighter than air. Its specific gravity compared 
with air as unity is == 0.06926. In many respects it is 
more convenient to compare the specific gravity of gases with 
II as unity. If the specific gravity of gases compared with 
H = 1 be represented by A, and the specific gravity com- 
pared with air = 1 by D, then i = f)X 14.46 and D 
A cubic decimeter (= 1 litre) of hydrogen weighs at 0°, and 
an atmospheric pressure of 760 millimeters (in the 49° of lati- 
tude) 0.089578 grams ; a litre of air, which is 14.46 times 
heavier, weighs, therefore, 1.2995 grams. 
That hydrogen is lighter than air is shown by a balloon of 
collodium or gum filled with hydrogen rising in it; this can 
also be seen in soap bubbles filled with hydrogen. In conse- 
quence of its levity, hydrogen may be collected in invei'ted 
vessels (opening turned down), 
by replacing the air, and can 
also be poured from one cylin- 
der into another, as is repre- 
sented in (Fig. 32). Owing to 
its levity, H 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. Owing to its levity and 
mobility, hydrogen penetrates porous bodies with ease, and 
diffuses both through animal and vegetable membranes, as well 
as through gutta percha. Metals, such as iron, platinum, palla- 
dium, allow, when in glowing condition, a free passage to hydro- 
gen, whilst for other gases they are impenetrable ; this behavior 
Fig. 32. HYDROGEN. 
37 
is probably dependent partly upon the chemical attraction of 
these metals for hydrogen. 
Chemical Properties. Hydrogen is characterized by its 
ability to burn in the air, at the same time combining with 
the oxygen of the latter and forming water; hence its name 
hydrogenium (from oowp water, and 
yswau), I produce). Its flame is faint 
blue and almost non-luminous. 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 fully escape, 
otherwise 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 
atmosphere of the gas contained in an 
Fig. 33. 
Fia. 34. 
inverted cylinder (Fig. 33), the latter will ignite at the mouth 
of the vessel, but the candle is extinguished in the hydrogen 
gas. 
Water is the product of the combustion of hydrogen in the 
air. It is a chemical compound containing hydrogen and 38 
INORGANIC CHEMISTRY. 
oxygen. To render the formation of it visible, by the combus- 
tion of hydrogen, the dame of the latter is made to burn under 
a cold glass jar (Fig. 34). The sides of the latter are soon 
covered with moisture, which collects in drops of water. To 
avoid any deception the hydrogen is first conducted through 
a tube filled with calcium chloride, to absorb all moisture. 
The immediate union of hydrogen with oxygen only 
occurs at a high temperature, either in contact with a flame 
or by the electric spark. The combination can be effected at 
ordinary temperatures with the aid of platinum sponge; the 
latter consists of finely divided metal, obtained by the ignition 
bf ammonio-platinum 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 ability 
of the metal to condense hydrogen and oxygen upon its sur- 
face, and thereby increase the ability of the gases to unite. 
Upon this behavior depends the 
action of the so-called JJwbereiner 
Lamp (Fig. 35). This consists of a 
continuous hydrogen generator. The 
outer glass cylinder, c, is filled with 
dilute sulphuric acid, into which the 
pear-shaped vessel, b, projects. This 
is open below and above, provided 
with a stop-cock, e, by means of 
which it communicates with the air ; 
in it a piece of zinc is suspended by 
a wire. On opening the stop-cock 
the sulphuric acid presses from the 
outer cylinder a into b and meets 
the zinc—when the liberation of 
hydrogen begins. The stop-cock 
directs the gas up on the support, f, 
in which is fixed some platinum 
sponge, where it is ignited. Upon 
again closing e the gas causes the 
acid to recede from the inner vessel, 
the zinc is freed of acid and the 
hydrogen evolution ceases. 
Fig. 35. 
Very characteristic for hydrogen is its absorption by the 
metal palladium. As already known, water is so decom- 
posed by the electric current that hydrogen separates at the 
electro-negative pole and oxygen at the electro-positive. If a 
piece of palladium, in sheet or wire form, be attached to the 
electro-negative pole, the disengagement of the gas ceases, HALOGENS. 
39 
because it is absorbed by the palladium, and in the ratio of over 
nine hundred times its volume of gas. The palladium expands, 
becomes lighter in weight, but retains its metallic appearance. 
Its tenacity and power of conducting heat and electricity are 
but little impaired. The compound of palladium and hydro- 
gen, Pd2H, therefore, conducts itself like an alloy of two 
metals. From the specific gravity of the compound (accord- 
ing to Graham), the specific gravity of the condensed hydro- 
gen is readily found to be 0.62 (water = 1), therefore, some- 
what heavier than the metal lithium. Also the metals 
potassium and sodium when heated from 200 to 400° absorb 
hydrogen, forming the alloys (Na2H and K2H,) in which the 
density of hydrogen also equals 0.62. From these facts we 
could conclude that hydrogen possesses a metallic character, 
and that by condensation it will form a metallic, mercury- 
like liquid—a conclusion which at present has been confirmed 
by experiment. Its metallic character is also shown, as we 
will observe later, by its entire chemical deportment. Thus, 
under great pressure, hydrogen is able to expel different 
metals from their salt solutions, the same as other metals do. 
When palladium hydride is heated it suffers decomposition, 
hydrogen escaping, just as in mercury alloys the mercury is 
expelled. 
HALOGEN GROUP. 
To this group belong chlorine, bromine, iodine and 
fluorine. The latter is not well known in a free condition. 
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. 
1. CHLORINE. 
Cl — 35.4 
Cl2 — 70.8 
It does not occur free in nature, as it is endowed with 
strong affinity for the majority of the elements. Its most 
important derivative is sodium chloride, or rock salt. The 
Swedish chemist, Scheele, discovered it in 1774. Gay-Lussac 
and Thenard, in France (1809), however, and Davy, in 
England (1810), first established its elementary character. 
Preparation. To obtain free chlorine, heat a mixture of 40 
INORGANIC CHEMISTRY. 
black oxide of manganese (Mn02) and hydrochloric acid in a 
flask (Fig. 36), provided with a so-called Welter’s safety tube, 
to equalize the gas pressure. The escaping gaseous chlorine is 
washed and freed from acid mechanically carried along by 
passing through the water in a three-necked Woulff’s bottle, 
Fig. 36. 
then collected over the same liquid. The reaction which 
occurs above is indicated in the following equation :— 
The manganous chloride formed dissolves in the water. 
The evolution of the chlorine proceeds more regularly if a mixture 
of manganese oxide (5 parts), sodium chloride (4 parts) and sulphuric 
acid (12 parts diluted with six of water) is employed :— 
MnOa -f 4HC1 = MnCl2 + Cl2 + 2H20. 
Mn02 + 2NaCl + 2H2S04 = MnS04 + Na2S04 + Cl2 + 2H20. 
Manganese 
di-oxide. 
Sodium 
chloride. 
Sulphuric 
acid. 
This reaction comprises two phases: First, the sodium chloride 
(NaCl) is decomposed by the sulphuric acid, yielding sodium sulphate 
and hydrochloric acid :— 
2NaCl + H2S04 = Na2S04 + 2HC1. 
The latter acid then acts, together with a new portion of sulphuric 
acid, upon the manganese dioxide :— 
Mu02 + H2S04 + 2H01 = MqS04 + 2H20 + Cla. HALOGEN 
41 
The second method is more advantageous for laboratory purposes ; 
the first, however, is preferred in practice, as it is cheaper. 
The resulting manganous chloride (MnCl2) is converted by the 
Process of Weldon into manganese superoxide (see this). Techni- 
cally, chlorine is also obtained by the Process of Deacon, by conducting 
HC1 mixed with air over glowing porous substances (bricks) saturated 
with metallic salts (copper sulphate). 
As chlorine gas dissolves readily in cold water it is advisable to 
collect it over warm. Over mercury it cannot be collected, as 
it readily combines with the latter. When perfectly dry chlorine is 
sought, conduct the liberated gas through WoulfTs bottles containing 
sulphuric acid, to absorb the moisture, then collect in an empty 
upright flask (compare Fig. 44, p. 54). As chlorine is so much 
heavier than air it will displace the latter. 
Physical Properties. Chlorine is a yellowish-green gas 
(from '/Xwpbq), with a penetrating, suffocating odor. Its spe- 
cific gravity compared with hydrogen (1) is 35.4; with air 
(= 1) it is A2 = 2.45. At 15° C., and a pressure of 4 atmos- 
pheres (at —40° C., at ordinary temperatures) it condenses to 
a yellow liquid. To effect the condensation of chlorine take a 
bent glass tube (Fig. 37), introduce into the leg closed at one 
end some crystals of chlorine 
hydrate (Cl2 -f 10H2O, see 
below), then seal the open 
end. The limb containing 
the compound is placed in a 
water bath; the other is cooled 
in snow. Upon heating the 
water to a little above 30° the 
chlorine hydrate is decom- 
posed into water and chlorine 
gas, which condenses to a 
liquid in the covered limb. 
On reversing the position of 
the limbs and cooling the one 
previously warmed, the chlo- 
rine distills back and is reabsorbed by the water. Charcoal 
saturated with chlorine may be substituted for the chlorine 
hydrate. This substance takes up 200 volumes of chlorine, 
which are again disengaged on heating. 
One volume of water, at 20° C., absorbs 2 volumes of chlo- 
rine ; at 8° C., 3 volumes. The aqueous solution is known as 
chlorine water (aqua chlori), and possesses almost all the 
Fig. 37. 42 
INORGANIC CHEMISTRY. 
properties of the free gas; therefore frequently employed for 
laboratory uses as a substitute for it. The yellow, scale-like 
crystals of chlorine hydrate (Cl2 + 10H2O) separate when 
water saturated with the gas is cooled below 0°. This com- 
pound is regarded as one of chlorine with water. At ordinary 
temperatures it decomposes into its constituent atoms. 
Chemical Properties. Chlorine has great affinity for almost 
all the elements. It combines, at ordinary temperatures, with 
the most of them, to form chlorides. When thin sheet copper 
(impure gold leaf), or, better, pulverized antimony or arsenic, 
are thrown into a vessel filled with dry chlorine, they burn 
with a bright light; a piece of phosphorus will also inflame 
in an atmosphere of the gas. 
Chlorine unites with hydrogen just as energetically. A 
mixture of equal volumes of the gases combines in direct sun- 
light, with violent explosion. In dispersed sunlight the action 
is only gradual; in the dark it does not occur. The great 
affinity of chlorine for hydrogen is manifested in the hydrogen 
compounds; most of these are so decomposed by the chlorine 
that it removes the hydrogen from them, and forms hydro- 
chloric acid. Thus water is decomposed by Cl into hydro- 
chloric acid and oxygen :— 
If a glass cylinder be filled with, and inverted over, chlorine 
water and exposed to the sunlight, the evolution of gas, which 
will collect in the upper portion of the vessel, will soon be ob- 
served ; this is oxygen. In diffused light the decomposition 
will not be so rapid ; it is hastened by heat. 
Chlorine alters the hydrocarbons, in that it abstracts hydro- 
gen. The reaction is sometimes so violent that carbon is sepa- 
rated in free condition. A piece of tissue paper saturated 
with newly distilled turpentine oil, introduced into a dry 
chlorine atmosphere, is immediately carbonized. An ignited 
wax candle immersed in chlorine burns with a smoky flame, 
with separation of carbon. 
The organic (containing C and H) dye stuffs are decolor- 
ized by moist chlorine gas. The same occurs with the dark 
biue solutions of indigo and litmus; colored flowers are also 
rapidly bleached by it. On this principle depends the appli- 
cation of chlorine in bleaching goods, and in destroying de- 
caying matter and miasmas (chlorine disinfection. See 
Bleaching Lime). 
H20 + Cl 2 = 2HC1 + 0. HALOGENS. 
43 
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 (see above). This property free oxygen does 
not possess, it does, however, very probably belong to that which is in 
the act of forming—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 free ; the cause for this will be ex- 
plained later. 
2. BROMINE 
Br = 79.7. 
Bra = 159.4. 
Bromine, the perfect analogue of chlorine, was discovered 
by Balard, in 1826. It occurs as sodium bromide, accompa- 
nied by sodium chloride in, however, much smaller quantity, 
in the sea (especially in the water of the Dead Sea), in many 
salt springs, as at Kreutznach and in Hall. When sea water, 
or other salt water, is evaporated, sodium chloride first sepa- 
rates ; in the mother-liquor, among other soluble salts, are 
found sodium and magnesium bromides. The principal source 
for the bromine yield are the upper layers of the rock salt 
deposits of Stassfurth, near Magdeburg, which contain bro- 
mides 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:— 
Mn02 + 2NaBr -f- 2H2S04 = MnS04 -f- Na2S04 -f- Br2 ~f* 2H20. 
The operation can be executed in the apparatus pictured in 
Fig. 38. This can also be used for many other distillations. 
The retort A, containing the mixture, is heated in a water 
bath; the tube B serves to cool the vapors which are con- 
densed by cold water or ice in the receiver C. When free 
chlorine is conducted into an aqueous solution of sodium bro- 
mide, bromine separates out. 
Bromine is a heavy, reddish-brown liquid, of an exceedingly 
penetrating, chlorine-like odor (hence the name Bromine, 
from stench). At 7.3° it crystallizes to a yellowr-green, 
scaly mass, having metallic lustre and resembling iodine. 
Liquid bromine, at 0°, has the specific gravity, 3.18 (w7ater = 
1) ; it is very volatile, yielding dark brown vapors at medium 
temperatures, and boiling at 63° C., converting itself, at the 
same time, into a yellowish-brown vapor. Its density equals 
79.7 (hydrogen = 1), or 5.53 (air = l). 44 
INORGANIC CHEMISTRY. 
Bromine is more soluble in water than chlorine. Cooled 
below 4° C., the hydrate (Br2 + 10H2O), analogous to that 
of chlorine, crystallizes out. This is decomposed at medium 
temperatures. Bromine dissolves with ease in alcohol, and 
especially in ether, chloroform and carbon disulphide. 
Fig. 38. 
In a chemical point of view, bromine is extremely like 
chlorine, combining with most metals to form bromides; but 
it possesses a weaker affinity than chlorine, and is liberated by 
the latter from its compounds:— 
KBr + Cl = KC1 + Br. 
With hydrogen it only combines on warming, not in sun- 
light. Upon hydrocarbons it acts like chlorine, withdrawing 
hydrogen from them. Bromine water gives starch an orange 
color. 
3. IODINE. 
I = 126.5 
I, = 253. 
Iodine occurs, as well as bromine, in combination with sodium* 
in sea water and some mineral springs, especially at Hall, in 
Austria, and the Adelheit Spring, in Bavaria. In most 
springs the iodine can easily be directly detected ; in sea water 
it is, however, only present in such minute quantity that its 
separation, practically, is disadvantageous. Sea-algae absorb 
it from the water, and these are then thrown by the tide on HALOGENS. 
45 
various coasts, where it is burned, yielding an ash (known as 
kelp in Scotland, as varec in Normandy) which represents the 
principal source for the manufacture of iodine. It was in this 
ash that the element was accidentally discovered, in 1811; in 
1815, it was investigated by Gay-Lussac, and its elementary 
character established. To obtain the iodine, the ash is treated 
with water, the solution concentrated and the sodium and 
magnesium iodides are further worked up. Lately, iodine has 
been obtained from the mother-liquors of the crude Chili salt- 
petre. It is set free from its compounds in the same manner 
as chlorine and bromine—by distillation with manganese 
dioxide and sulphuric acid. It is more convenient to pass 
chlorine (or, better, nitrous acid) through a solution of the 
salt, when all the iodine will separate:— 
KI + Cl = KC1 +1. 
The grayish-black powder thus liberated is collected on a 
filter, dried, and then sublimed. 
Iodine is a gray-black solid, subliming in large rhombic 
crystals, possessing strong metallic lustre. It has a peculiar 
odor, reminding one somewhat of that of chlorine; it stains 
the skin brown, and is corrosive, although not as strong as bro- 
mine. Its specific gravity is 4.95. Fuses at 113° to a dark 
brown liquid, and boils near 200°, passing at the same time 
into a dark violet vapor (hence the name iodine, from 
violet-blue). 
The vapor density of iodine equals 8.7 up to 600° C., (air = l) or 
126.5 (H = 1), corresponding to the molecular value I2 = 253. 
Above 600° the vapor density gradually diminishes and about 1500° it 
is only half the original. This is explained by the gradual decomposi- 
tion (see Dissociation of water) of the normal diatomic molecule I2 
into the free atoms I + I- In like manner, the bromine molecules 
Br2 also, at high temperatures, suffer a separation into the free atoms, 
while the vapor density of chlorine remains unchanged. 
Iodine is very slightly soluble in water, more readily in 
alcohol ( Tinctura Iodi), very easily in ether, chloroform and 
carbon disulphide, the two last assuming a deep red violet 
color in consequence. It crystallizes particularly beautiful, in 
forms of the rhombic system, 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 set free, by those elements, from its compounds. 46 
INORGANIC CHEMISTRY. 
With the metals it usually only combines when warmed ; 
with hydrogen it does not combine directly, and it does not 
remove hydrogen from its carbon compounds. 
Characteristic for iodine is the deep blue color it imparts 
to starch. On adding starch-paste 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. 
4. FLUORINE. 
Fluorine possesses such strong affinity for almost all sub- 
stances that it cannot be obtained free; it attacks glass and 
even platinum vessels. Its most important compound is 
fluorite (calcium fluoride CaFl2), which is used for the pre- 
paration of the others, which are very similar to the chlorine 
derivatives. When silver fluoride is heated in a stream of 
chlorine, or calcium fluoride in a current of oxygen, there 
escapes a colorless, very violent smelling gas, which attacks 
glass and all the metals. Upon conducting the same into 
water, oxygen is disengaged from the latter:— 
FI = 19 (Fl2 = 38). 
This gas is probably fluorine. Its affinity for the metals 
and hydrogen is much greater than that of chlorine. 
Upon the basis of theoretical observations developed later 
the specific gravity of free fluorine is 19 (hydrogen == 1). 
H20 + Fl2 = 2 H FI + 0. 
The four observed and similar elements, fluorine, chlorine, 
bromine and iodine, exhibit gradual differences in their prop- 
erties, and what is remarkable, this gradation stands in direct 
relation to the specific gravity of the elements in the state of 
gas or vapor. 
Specific Gravity 
FI 
19 
Cl 
35.4 
Br 
79.7 
I 
126.5. 
With the increase of specific gravity occurs a simultaneous 
condensation of matter, which expresses itself in the diminished 
volatility. Fluorine is a gas ; chlorine can readily be con- 
densed to a liquid; bromine is a liquid at medium tempera- HYDROGEN CHLORIDE. 
47 
tures, and iodine finally is a solid. Other physical properties, 
as seen in the following table, are also in accord with the 
preceding. 
Fluorine. 
Chlorine. 
Bromine. 
Iodine. 
Fusing point 
— 7° 
+ 107° 
Boiling point 
— 33° 
+ 63 
+ 200° 
Specific gravityin liquid or solid 
condition 
1.33 
3.18 
4.97 
Color 
Colorless 
Yellow 
Brown 
Black 
Just such a gradation, as we have seen, is observed in the 
chemical affinities of these four elements for the metals and 
hydrogen ; fluorine is the most energetic, iodine the least. 
Therefore, each higher element is separated from its soluble 
metal and hydrogen compounds by the lower. # We will dis- 
cover, later, that in the affinity of the halogens for oxygen 
and some other metalloids, the reverse is true. 
COMPOUNDS OF THE HxVLOGENS WITH HYDROGEN. 
With hydrogen the halogens form compounds of an acid 
nature, readily soluble in water. 
1. HYDROGEN CHLORIDE. 
HC1 — 36.4. Density = 18.25 
The direct union of chlorine with hydrogen takes place 
when heated and by the action of direct sunlight or other 
chemically active rays ; in diffused light the action is only 
gradual, and does not occur at all in the dark. On intro- 
ducing a flame of hydrogen ignited in the air into a cylinder 
filled with chlorine (Fig. 39), it will continue to burn in the 
latter. The opposite, the combustion of chlorine in an atmos- 
phere of hydrogen, may be easily shown in the following 
experiment (Fig. 40). An inverted cylinder is filled with 
hydrogen by displacement, the gas is ignited at the mouth, 
and a tube immediately introduced which will conduct dry 
chlorine into the cylinder. The burning hydrogen will in- 
flame the chlorine, which will continue to burn in the former. 
From these experiments, we perceive that combustibility 
and combustion are only relative phenomena ; if hydrogen is 
combustible in chlorine (or air), so, inversely, is chlorine (or 
air) combustible in hydrogen. By the term combustion, in 48 
INORGANIC CHEMISTRY. 
chemistry, is understood every combination of a body (with 
gaseous aid) which is accompanied by the phenomenon of 
light. 
A mixture of equal volumes of chlorine and hydrogen, 
made under the above conditions, of the union of the gases, 
Fig. 40. 
Fig. 39. 
explodes with very great violence. The product is hydrogen 
chloride. 
The formation of the latter compound succeeds best by al- 
lowing sulphuric acid to act upon sodium chloride, when solid 
sodium sulphate and hydrogen chloride gas will result:— 
Pour over 5 parts sod. chloride, 9 parts sulphuric acid, some- 
what diluted (2 parts), and warm the mixture gently in a 
flask, A (Fig. 41). The escaping hydrogen chloride is con- 
ducted through a Woulff’s bottle, containing sulphuric acid 
to the cylinder B (filled with pumice stone saturated with 
sulphuric acid), intended to free it from all moisture, and 
afterwards collected over mercury. 
Physical Properties. Hydrogen chloride is a colorless gas 
with a suffocating odor. In moist air it forms dense clouds. 
Under a pressure of 40 atmospheres (at 10° C., or 1 atmos- 
phere at — 80.3°), it condenses to a liquid, with a specific 
gravity of 1.27, which does not solidify at — 110° and boils 
at — 80.3°. 
2NaCl -f H2S04 = Na2S04 + 2 HC1. HYDROGEN CHLORIDE. 
49 
The specific gravity of the gas (density) is 18.25 (II = 1), 
or 1.26 (air = 1). 
Hydrogen chloride possesses an acid taste and colors blue 
litmus paper red; it is, therefore, an acid, and has received 
the name hydrochloric acid gas. It dissolves very readily in 
water, and on that account cannot be collected over it. One 
volume of water at 0° dissolves 505 volumes, and at ordi- 
nary temperatures about 450 volumes of the gas. This great 
solubility is very nicely illustrated by filling a long glass 
Fig. 41. 
cylinder with the gas and then just dipping its open end 
beneath water ; the latter rushes up into the vessel rapidly, as 
it quickly absorbs the gas. The aqueous solution of hydrogen 
chloride, in ordinary language, is known as muriatic or hydro- 
chloric acid (acidum hydrochloratum). For its preparation 
the gas is passed through a series of Woulff bottles (Fig. 42) 
containing 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 em- 
ployed in the manufacture of chlorine water, and is generally 
used in the saturation of liquids with gases. 
A solution saturated at 15° C., contains about 40 °/c hydro- 
gen chloride, has a specific gravity of 1.2, and fumes in the air. 50 
INORGANIC CHEMISTRY. 
On the application of heat the gas again escapes, and the 
temperature of the liquid rises to 110° C., when a liquid 
distills over, containing about 20 % of hydrogen chloride, 
having a specific gravity of 1.104 and almost corresponds to 
the formula HC1 -f 8H20. The composition of the dis- 
tillate varies somewhat with the pressure. On conducting 
hydrogen chloride into hydrochloric acid cooled to — 22°, 
Fio. 42. 
crystals of the formula HC1 -f- 2H20 separate, which fuse at 
— 18° and then decompose. 
Hydrochloric acid finds an extensive industrial application, 
and is obtained in large quantities, as a bye-product, in the 
soda manufacture. 
Chemical Properties.— Acids — Bases—Salts. 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 acid, reddens blue litmus paper, and saturates the 
bases (metallic oxides and hydrates) ; i. e., such bodies as 
impart a blue color to red litmus paper. If we add hydro- 
chloric acid to a solution of a base, e. g., sodium hydrate, 
until the reaction is neutral, we will obtain (besides water) a 
neutral, solid compound—sodium chloride. 
Sodium 
hydrate. 
NaOH + HC1 = NaCl + H20. 
Sodium 
chloride. HYDROGEN CHLORIDE. 
51 
HBr, HI and IIF1 deport themselves similarly to HC1. 
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 con- 
duct themselves like the former, and saturate bases, forming 
salts and acids:— 
KOH + HN03 = KN03 + H20. 
In the same manner the acids act upon the basic oxides, to 
form salts and water:— 
Potassium 
hydrate. 
Nitric 
acid. 
Potassium 
nitrate. 
Water. 
ZnO + 211Cl 2 = ZnCl2 + II20. 
Zinc 
oxide. 
Zinc 
chloride. 
ZnO + 2HN03 = Zn (N03)2 + H20. 
When acids act upon metals, generally, the hydrogen of the 
former is directly displaced ; salts and free hydrogen are pro- 
duced. Thus, by the action of hydrochloric acid upon 
sodium, its chloride and water result:— 
Zinc 
oxide. 
Zinc 
nitrate. 
From the cited examples it is manifest that acids are such 
hydrogen compounds as, by the replacement of their hydrogen 
by metals (by the action of metallic oxides, hydroxides or free 
metals), yield salts. The oxides capable of forming 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 formed by the mutual action of bases 
and acids. Salts are distinguished as haloid salts and oxygen 
salts. The first have no oxygen and arise in the direct union 
of the halogens with the metals. 
Hydrogen chloride is a very stable compound, suffering 
only a partial decomposition at 1500° C. Its composition is 
easily analytically established by the following experiments : 
Pass hydrochloric acid gas over a piece of sodium, heated 
in a glass tube, and hydrogen will escape from the latter :— 
HC1 + Na = NaCl + H 
If manganese superoxide, on the other hand, be heated in it, 
chlorine will be disengaged :— 
Na + HC1 = NaCl + H. 
Mii02 + 4HC1 — MiiC12 + 2II20 + Cl2. 52 
INORGANIC CHEMISTRY. 
If the electric current be permitted to act upon an aqueous 
solution of hydrochloric acid, the latter will be so decom- 
posed that chlorine separates at the electro-positive and 
hydrogen at the electro-negative pole. 
2. HYDROGEN BROMIDE, 
Hydrogen bromide is perfectly similar to the corresponding 
chlorine compound. As there is but slight affinity between 
Br and H their direct union will only occur at a red heat or 
in the presence of platinum sponge (see p. 38). Like hydro- 
gen chloride, hydrogen bromide can be obtained by the action 
of some acids, e. g., P205, upon bromides; sulphuric acid 
wrould not answer, as the resulting HBr is again partly decom- 
posed by it. Ordinarily it is prepared by the action of phos- 
phorus tri-bromide (see Phosphorus) upon water:— 
HBr = 80.7. Density = 40.3. 
PBr3 -f- 3H20 = H3P03 -j- 3HBr. 
Phosphorus 
tri-bromide. 
Phosphorus 
acid. 
Fig. 43. 
Place some water (1 part) in a flask (Fig. 43), gradually 
admit through the funnel, supplied with a cock, the liquid, 
PBr3 (3 parts), and warm gently. The escaping HBr gas is 
collected over mercury or conducted into water. To free it 
perfectly from accompanying PBr, vapors it is passed through 
water (the U-shaped tube, in Fig. 43, contains pieces of 
pumice stone, which are moistened with water). HYDROGEN IODIDE. 
Instead of employing prepared brom-phosphorus, we may let 
bromine vapors act upon phosphorus moistened with water. As the 
action of bromine upon ordinary yellow phosphorus is too energetic, it 
is more practical to employ red amorphous phosphorus. 
To obtain an aqueous solution of the gas, pour 15 parts 
H20 over 1 part amorphous phosphorus, and then add Br(10 
parts) drop by drop. Finally the solution is heated, filtered 
and distilled. From bromides (NaBr, KBr) hydrogen 
bromide is obtained by distillation with somewhat dilute 
sulphuric acid, with addition of phosphorus. 
Avery’s convenient method to obtain the gaseous HBr is 
the following: Heat a mixture of 6 parts sodium sulphite 
(S03Na2), 3 parts bromine and 1 part water. HBr is at 
once disengaged. 
Hydrogen bromide is a colorless gas, burning strongly in the 
air. Under great pressure it is condensed to a liquid, boiling 
at — 73.3°, and solidifying at — 120°. Its density is 40.5 
(H = 1) or 2.71 (air = 1). 
In water the gas is very readily soluble, its saturated solu- 
tion having a specific gravity of 1.78, and containing 82 % 
HBr ; at 15° it contains 49.8 %, and has the specific gravity 
of 1.515. At 125° C a solution distills over containing 
46.8 °}o HBr, and closely approximates the formula HBr -j- 
5H20; its specific gravity is 1.47 at 14° C. 
On conducting HBr into a solution of same cooled to — 20°, 
crystals of the formula HBr 2H,0 separate and melt at— 
— 11°. Chemically, HBr is the perfect analogue of HC1; it 
is, however, less stable, and suffers a partial decomposition 
at 800° C. 
3. HYDROGEN IODIDE 
The attraction of iodine for hydrogen is very slight. Their 
union (and then only partial) occurs when both elements, in 
the form of vapor, are conducted over 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:— 
HI = 127.5 Density = 63.7 
PI5 + 4H20 = PO (OH), + 3HI. 
A more convenient procedure, is to warm a mixture of 
amorphous phosphorus (1 part), iodine (20 parts), and water 
(15 parts) ; and an analogous reaction will ensue. Or add a 54 
INORGANIC CHEMISTRY. 
solution of 2 parts iodine in 1 part hydriodic acid, of specific 
gravity 1.7, drop by drop, to red phosphorus, and aid the re- 
action by heat. As HI dissolves readily in water, and is 
decomposed by mercury, we can only collect it by conducting 
it into a dry flask (Fig. 44), where in consequence of its five- 
fold greater density, it will displace the air. 
Fig. 44. 
To get an aqueous solution of HI, take more water, warm 
the solution, filter, and then distill. 
Another method of obtaining III consists in passing hydro- 
gen sulphide into water having iodine dissolved in it:— 
Filter oft'the separated sulphur and distill the liquid. 
Hydrogen iodide is a colorless gas ; it fumes strongly in the 
air; its density is 3.7 (H = 1) or 4.41 (air = l). Under a 
pressure of 4 atmospheres (at 0°) it is condensed to a liquid 
which solidifies at—55°. It is easily soluble in water; the 
solution saturated at 0° C., has a specific gravity, 1.99, and 
fumes strongly in the air. At 127° a solution of 1.56 specific 
gravity, and containing 57.7 °Jc HI distills over, corresponding 
closely to the formula HI -f- 5H20. 
Hydrogen iodide is a rather unstable compound, decompo- 
sing at 180°, into hydrogen and iodine. At high temperatures 
oxygen decomposes it into water and iodine :— 
H2S + I2 = 2HI-f S. 
On bringing a flame near a vessel containing a mixture of 
HI and oxygen, violet iodine vapors will at once fill it. The 
2HI + 0 = H20 + I2. HYDROGEN FLUORIDE. 
55 
same will be noticed when fuming nitric acid is dropped 
into a vessel containing HI; in this reaction the oxygen of 
the acid oxidizes the hydrogen and liberates iodine. All 
oxidizing bodies behave in the same way; the hydrogen 
iodide abstracts their oxygen and reduces them. The oxygen 
of the air, even at medium temperatures, gradually decomposes 
aqueous hydrogen iodide. The solution, at first colorless, 
becomes brown, owing to separation of iodine, which in the 
beginning dissolves; subsequently, however, it separates in 
beautiful crystals. 
At ordinary temperatures, mercury and silver decompose 
HI, with separation of hydrogen:— 
HI + Ag = AgI + H. 
Chlorine and bromine liberate iodine from HI. 
This compound is employed as a powerful reducing agent 
in laboratory work. 
4 HYDROGEN FLUORIDE. 
H Fl= 20. Density = 10, 
It is obtained, like hydrogen chloride, by decomposing fluo- 
rides with sulphuric acid. Finely pulverized fluorite is mixed 
with II2S04and heat applied gently:— 
CaFl2 + H2S04 = CaS04 + 2HF1. 
Calcium 
fluoride. 
Calcium 
sulphate. 
The operation is executed in a lead or platinum retort, as 
the hydrogen fluoride attacks glass and most of the metals. 
(Fig. 45). The escaping HF1 condenses in the U-shaped 
Fig. 45. 56 
INORGANIC CHEMISTRY. 
receiver containing some water. To get perfectly anhydrous 
hydrogen fluoride, heat hydrogen potassium fluoride, which 
then decomposes according to the following equation :— 
KF12H = KF1+ HF1. 
Anhydrous hydrogen fluoride is a colorless, very mobile 
liquid, fuming strongly in the air, and attracting moisture with 
avidity ; it boils at -j- 19° C., and has a specific gravity, 0.98 
at 12°. To recondense the gas it must be cooled to — 20°. 
The concentrated aqueous solution fumes in the air ; when 
heated HF1 escapes; the boiling temperature increases 
regularly until constant at 120°C., a solution distills over, 
whose sp. gr. = 1.15 and whose contents of HFl = 35.37%. 
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 anhydrides 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 glass. To perform the latter operation, coat the glass 
with a thick 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 HFl, The exposed portions of the glass appear 
etched; gaseous HFl furnishes a dim, and liquid HFl 
a smooth, transparent etching. 
Vessels of lead, platinum or caoutchouc are employed for 
the preservation of hydrofluoric acid, as they are not affected 
by it. 
The halogen derivatives of hydrogen show great resem- 
blance to each other. At ordinary temperatures they form 
strongly smelling and fuming gases, which by pressure can be 
condensed to liquids. Their fuming in moist air is due to the 
fact that the gases are condensed by the aqueous vapor. 
Readily soluble in water, they are only partially expelled 
from their solution by boiling. As acids they neutralize the 
bases and form haloid salts, which also can result by the direct 
union of the halogens with metals. 
The densities .of the hydrogen haloids exhibit a gradation 
similar to that of the densities of the halogens (page 46):— 
HF1 
10 
HG1 
18.2 
HBr 
40.5 
HI 
64. 
Density, HYDROGEN FLUORIDE. 
The difference in chemical deportment corresponds to this 
gradation. Hydrogen fluoride is the most stable, and acts 
most energetically; chlorine combines with hydrogen in 
sunlight, bromine only at a red heat, while iodine and 
hydrogen do not react at all. On the other hand, hydrogen 
iodide is decomposed at a gentle heat (180°), into its 
constituents ; the more stable hydrogen bromide at 800°, 
while hydrogen chloride remains unaltered at 1500°C. 
Corresponding to this we have the very energetic action of 
fluorine, the tolerably ready action of chlorine upon water, 
oxygen separating at the same time :— 
H20 + Cl2=2HCl + 0 
Iodine does not act upon water. The opposite reaction occurs: 
oxygen decomposes hydrogen iodide into water and iodine:— 
2HI + 0 = H20 + I2. 
Bromine stands intermediate between chlorine and iodine; in 
aqueous solution it decomposes water into HBr -(- O, while a 
concentrated solution of hydrogen bromide, on the contrary, 
is partly decomposed by oxygen into water and free bromine. 
From all the above it is evident that the affinity of fluorine 
for hydrogen is the greatest; then follows chlorine and bro- 
mine, and finally, as the least energetic element, we have 
iodine. (See p. 47.) 
A measure for the chemical affinity of the halogens to hydrogen, as 
also of the chemical elements and compounds for each other, is the 
determination of the quantities of heat liberated or absorbed in chem- 
ical unions. 
The quantity of heat is determined according to heat units or calo- 
ries, and as unit is selected that quantity of heat which is required to 
raise 1 part of water to 1° C. The quantities of the elements are ex- 
pressed in numbers corresponding to the atomic weights. Hence, in 
the union of 35.4 parts chlorine with 1 part hydrogen, 22,000 calories, 
in the union of 79.7 parts bromine with 1 part hydrogen, 8400 calories 
are developed, while in the union of 126.5 parts iodine with 1 part 
hydrogen. 6200 calories are absorbed. This may be more simply ex- 
pressed by the following: 
(H,C1) = + 22,000 (H,Br) = -f- 8400 (H,I)= — 6200. 
The first two reactions, in which heat is liberated, are exothermic, 
while the heat absorbing combination of iodine with hydrogen repre- 
sents an endothermic reaction, which, according to the principle of 
the greatest heat liberation (p. 17) can only be induced by the addition 
of external energy (heat). Consequently the thermal relations of the 58 
INORGANIC CHEMISTRY. 
halogens correspond to their varying chemical affinity, and in them the 
different deportment of halogen hydrides, in their formation and decom- 
Eosition, finds complete expression. In similar manner is explained 
ow, from the compounds of the halogens with hydrogen and the 
metals, iodine is eliminated by bromine and chlorine, and bromine by 
chlorine. In same manner, bearing in mind the thermal relations in 
the formation of water, is explained the varying decomposition of the 
hydrogen haloids by oxygen, and the reverse—that of water by the 
halogens. To illustrate these interesting relations, let us view the for- 
mation of hydrogen iodide, by action of hydrogen sulphide on iodine 
(p. 54), corresponding to the reaction: — 
H2S + I2 = 2HI + S. 
As in the production of hydrogen iodide, according to the equation, 
as given above, 12,400 (= 2 X 6,200) calories are absorbed—this re- 
arrangement cannot occur at ordinary temperatures. It does, however, 
transpire in presence of water, as the heat afforded by the solution of 
HI in water is sufficient. In the solution of hydrogen haloids, as also 
of other bodies, in water, a considerable amount of heat is set free, 
corresponding to the symbols:— 
(HCl,Aq) = 17,320. (HBr,Aq) = 19,940. (HI,Aq) =19,210. 
The heat (38,420 C.) set free by the solution of 2 HI in water exceeds 
the heat of formation of 2HI (12,400 C.) and the decomposition heat 
of H2S (9200 C.), so that the reaction can occur even with disengage- 
ment of heat. 
COMPOUNDS OF THE IIALOSENS WITH EACH 
OTHER. 
These compounds, resulting from the union of the halogens, 
like most of those of chemically similar elements, are very 
unstable. 
When chlorine is conducted over dry iodine, the latter being 
in excess, simple chlorine iodide results, and when the chlorine 
is in excess there is formed trichlor-iodine. 
Chlorine Iodide—IC1—is a red crystalline mass, fusing at 
24.7° C., and distilling a little above 100° C. Water decom- 
poses it easily, with formation of iodic acid, iodine and 
hydrogen chloride. 
Iodine Trichloride—IC13—is formed upon mixing iodic acid 
with concentrated hydrochloric acid, and by the action of 
PC15 upon P205. It crystallizes in long, yellow needles, and, 
when heated, suffers decomposition into IC1 and chlorine (at 
ordinary pressure, the dissociation commences at 25° C.) ; 
dissolves in little water without alteration; large quantities 
cause partial decomposition, with formation of iodic acid. WEIGHT PROPORTIONS. 
59 
Bromine Iodide—IBr—obtained by the direct union of the 
elements, consists of iodine-like crystals; fusing about 30°. 
Iodine Pentajhtoride—IF15—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. The Law of 
Constant Proportions. Atomic Hypothesis. 
If, in the considered halogen derivatives, as wrell as in all 
other chemical compounds, we determine the quantity of the 
elements (according to methods described in analytical che- 
mistry) we will discover that they are always combined with 
each other in the same proportions by weight. In every 
chemical compound the proportions by weight of the constituents 
contained in it are invariably the same. Thus chemical analysis 
shows the following percentage composition for the halogen 
derivatives of hydrogen :— 
H = 5.3 
Fl= 94.7 
HF1 —100.0 
H = 2.7 
Cl — 97.3 
HC1 = 100.0 
II = 1.2 
Br = 98.8 
HBr = 100.0 
H = 0.8 
I = 99.2 
HI = 100.0 
As experience has shown, H, of all the elements, enters com- 
pounds in the least quantity, therefore the latter is chosen as 
unity, and wTe calculate those weights of the elements which 
combine with one part by weight of H. In this manner we 
find the following proportions for the halogens :— 
H =• 1 
FI = 19 
HF1 = 20 
H = 1 
Cl = 35.4 
HC1 = 36.4 
II = 1 
Br = 79.7 
HBr = 80.7 
H = 1 
I = 126.5 
HI = 127.5 
Experiments have also established the remarkable fact that 
the same proportions by weight of the halogens are also 
obtained by the union of the same with other elements. 
Thus 19 parts by weight of FI combine with the following 
weights of the metals : 23 parts Na, 39 parts K, 32 parts Zn, 
31.7 parts Cu, 100 parts Hg and 35.4 parts Cl. 79.7 parts Br 
and 126.5 parts I combine with exactly the same quantities by 
weight of the metals. Let us take another example. On 
bringing copper into the solution of a mercuric salt the former 
dissolves, while Hg separates out; indeed, 31.7 parts Cu 60 
INORGANIC CHEMISTRY. 
displace 100 parts Hg. If zinc be brought into the copper 
solution thus obtained, it will dissolve, while copper separates— 
and 32 parts of Zn separate 31.7 parts Cu. Furthermore, 
zinc displaces the hydrogen in acids; from all of them 32 
parts Zn separate 1 part H. In all these reactions we observe 
the same quantities appearing, by weight, of the elements. 
These are the remarkable facts fully verified by experi- 
ment. Similar facts are formulated in a rule, and when a 
rule comprises a great number of facts—true for all and 
expressible in numbers—we designate it a law. The facts 
presented above find their expression in the empirical law of 
constant proportions, first proposed by Dalton, and reading : 
The elements combine with each other in definite proportions by 
weight, and indeed the proportions by weight of two elements 
remain the same in their combinations with other elements. 
Causes underlie facts. The cause is first expressed in the 
form of a supposition or hypothesis, and when the latter 
includes a long series of facts, if it is repeatedly substantiated 
by other phenomena and has acquired a high degree of prob- 
ability, it is termed a theory. 
If an hypothesis completely satisfies all the observations to 
which it refers, it becomes a fact, for the further explanation 
of which a new hypothesis may be necessary. Conversely, 
something which long passed as a fact or a theory can be 
shown to be erroneous, if not any longer consistent with new 
observations. Hypothesis and that which we designate a fact 
are distinguished really only by the different degree of prob- 
ability. If, for example, we make a sight observation we 
assume the hypothesis that the same has been caused by an 
external process, the reality of which (in distinction to sub- 
jective perceptions) we can only assure ourselves of being cer- 
tain by repeated observations. The hypothesis of the revolving 
of the earth, which at first was only a suitable, improbable 
supposition, proposed for simplifying calculation, has become 
a fact. The combustion theory of Lavoisier met a like result. 
Similarly may it deport itself with the supposition of atoms— 
be it that we comprehend them {is material particles or as ether 
motion. 
The law of constant proportions finds its clearest explana- 
tion in the hypothesis of the existence of atoms. Grecian phil- 
osophers even conjectured that matter consisted of indivisible 
and very small particles—atoms (from a, privative, and To/xoq, ATOMS. 
61 
divisible.) This a 'priori supposition was subsequently repeat- 
edly announced, but Dalton (1804) first gave it an actual 
confirmation, in that he applied it to the explanation of the 
law of constant proportions. According to the atomic view, 
matter consists of extremely small (although not indefinitely 
small) particles, atoms, which cannot be further divided, either 
mechanically or chemically. The atoms of different elements 
possess different weights ; all atoms, however, of one element 
have the same absolute weight and are like each other. By the 
aggregation of the elementary atoms arise the smallest particles 
of compound bodies. Upon the basis of this representation is 
very clearly explained the law of constant proportions ; we 
can comprehend that the quantities of the constituents of a 
compound should be constant, and that the relative quantities 
by weight of the elements, as they express the relative weights 
of the atoms, must in all their compounds be the same. 
As yet only the relative atomic weights of the elements have been 
determined by chemical researches ; in these the hydrogen atoms, as 
they possess the least weight, have been taken as unity. The knowl- 
edge of the absolute atomic weight has been, for chemical considerations, 
until now, unessential. Various physical phenomena allow us to fix, 
however, even at present, with very great approximate accuracy, the 
absolute size of the atom. Very different considerations lead to the 
same conclusion, that the atoms cannot be smaller than the fifty mil- 
lionth part of a millimeter (Thomson). 
If we grant that in the preceding halogen-hydrogen com- 
pound one atom of hydrogen is combined with every halogen 
atom, the conclusion follows, that the found numbers of the ratio 
express the atomic weights of the halogens. This supposition, 
however, appears questionable, in view of the more compli- 
cated proportions which occur in the union of some elements. 
Observation shows, to wit, that very frequently two elements 
unite with each other in not only one, but, indeed, several 
proportions. For example, 35.4 parts of chlorine combine 
not only with 31.7 copper and 100 parts mercury but also 
with 63.4 parts copper and 200 parts mercury. One part, by 
weight, of hydrogen, combines with 8 to form water, and 
16 parts oxygen (to form the so-called hydrogen peroxide) ; 
further, with 16 and 32 parts sulphur; oxygen forms five dif- 
ferent compounds with nitrogen, according to the following 
proportions, by weight:— 62 
INORGANIC CHEMISTRY. 
Nitrogen. 
Nitrous Oxide, 
14 parts. 
Oxygen. 
8 parts 
16 “ 
= 1X8. 
Nitric Oxide, 
14 “ 
= 2X8. 
Nitrous Anhydride, 
14 “ 
24 “ 
= 3X8. 
Nitrogen Dioxid^ 
14 “ 
32 “ 
= 4X8. 
Nitric Anhydride, 
14 “ 
40 “ 
= 5X8. 
Similar proportions are observed in the union of many 
other elements. Therefore, they combine with each accord- 
ing to several ratios by weight. As we have noticed in the 
examples given, the varying quantities of one of the elements 
(calculating upon the same quantity of the other element), 
bear a simple ratio to each other; they are mostly multiples 
of the smallest quantity. These facts are enunciated in the 
Law of Multiple Proportions, also proposed by Dalton (1807), 
which forms an essential amplification of the law of constant 
proportions. Based on the atomic hypothesis, these facts 
are explained by saying the elements can not only unite 
with each other, atom for atom, but in variable numbers. 
This considerably complicates the problem of determining the 
relative atomic weights of the elements, as these are directly 
dependent upon the conceived number of atoms in a com- 
pound. If, for example, in water, one atom of hydrogen is 
combined with one atom of oxygen, the atomic weight of 
the latter would = 8 (regarding that of hydrogen as 1). It 
is just as well possible that water consists of two atoms of H 
and O, or of one of H and two of O, etc.; in the first case the 
atomic weight of O would be = 16, in the latter = 4. 
Analytical results afford nothing positive for the solution of 
this difficulty. This was the condition in which, thirty years 
ago, the question as to the greatness of the atomic weights 
existed. To establish these correctly, various views were 
allowed to prevail, none, however, with positive foundation. 
The question can only be solved upon a new and accurate 
basis : first and foremost, for that purpose, secure the specific 
gravities of the chemical compounds in a gaseous or vapor 
form. 
Density of Bodies in State of Gas. Volume Ratio in the Union 
of Gases. Atomic Moleeular Theory. 
The halogens fluorine, chlorine, bromine and iodine, unite 
with hydrogen in only one proportion. The supposition, there- 
fore, is, that in the halogen-hydrogen compounds, 1 atom of 
Ii is combined with 1 atom of the halogen, the simplest and DENSITIES, 
most probable union. Then their weight proportions, derived 
from analysis, express directly their relative atomic weights. 
Upon now comparing these atomic numbers (referring to H = 
1, page 59), with those expressing the density in state of gas 
(also referred to H = 1), the astonishing result is that these 
two series of numbers are identical. 
Elements. 
Density. 
Air = 1 
Density. 
Hydrogen = 1 
Atomic Weights. 
Hydrogen, 
0.0693 
1 
1 
Fluorine, 
1.31 
19 
19 
Chlorine, 
2.44 
35.4 
35.4 
Bromine, 
6.39 
79.7 
79.7 
Iodine, 
8.71 
126.5 
126.5 
From this similarity of the atomic (combination) weights 
with the densities, follows the cogent conclusion that in 
equal volumes of these elementary gases there is contained the 
same number of atoms. Indeed, if in one volume of hydrogen, 
for example, are contained 1000 atoms of hydrogen, which 
equal 1000 weight units, and in a like volume of chlorine 
there are also present 1000 atoms of chlorine, which equal 
1000 X 35.4 weight units, then it is evident, that the rela- 
tion between the atomic weights, and that between the densities 
(the weights of like volumes) must be the same. 
1000 X 1 
1000 x 35.4 
These relations can be expressed by the following rule : 
The atomic weights of the halogen elements are proportional or 
equal to their densities, if referred to the same unit. Yield- 
ing to a too hasty generalization, this was incorrectly followed 
for all elements. 
1 Vol. Hydrogen. 
1 Vol. Chlorine. 
We arrive at a perfectly similar, but much more general 
conclusion, by consideration of the physical properties of gases 
or vapors. The similar deportment of the same under pres- 
sure (law of Mariotte and Boyle), their similar expansibility 
by heat (law of Charles and Dalton, ordinarily the law of 
Gay-Lussac), only appear comprehensible by the following 
suppositions. The gases consist of small portions of matter, 64 
INORGANIC CHEMISTRY. 
which are separated by equally, in proportion to the particles, 
great distances (those of the centre are alike and suffer like 
alterations). It immediately follows from this, that the same 
number of particles is contained in similar volumes of all gases 
(by like temperature and pressure.) The kinetic gas theory, 
based on the same supposition, explains the similar deportment 
of gases by the analogous existing force of the smallest 
gaseous particles. 
From the proposition, that in equal volumes an equal num- 
ber of particles is present, follows directly that their rela- 
tive weights are proportional to the volume weights or gas 
densities, and that by determination of the latter, the first are 
also given. In what ratio these smallest particles (called 
molecules) stand to the chemically smallest particles (atoms), 
remains, first of all, undetermined, and is obtained by a com- 
parison of the volume ratios according to which the bodies 
combine (p. 63). It is, however, even now perceptible, 
that, at least in case of compound bodies, the smallest gas par- 
ticles must be sums of atoms, as the same consist of complex 
atoms. 
It follows, from the similarity of the atomic weights and the 
densities, that the halogens must combine with hydrogen 
in similar volumes, since 1 part by weight of H combines 
with 35.4 parts by weight of chlorine, etc., and the weights of 
similar gas volumes stand in the same ratio. Further: 1 part 
H and 35.4 parts chlorine yield 36 4 parts HC1; one volume of 
the latter weighs, however, 18.2 (II = 1, p. 70) ; consequently, 
36.4 parts HC1 occupy 2 volumes. Therefore, equal volumes 
of H and Cl yield a double volume of HC1, or, as ordinarily 
expressed, 1 volume II and 1 volume Cl yield 2 volumes IdCl. 
In similar manner may be deduced that 1 volume H and 1 
volume Br vapor yield 2 volumes IIBr ; 1 volume H and 1 
volume I vapor, 2 volumes HI. 
These conclusions are confirmed by the following experiments:— 
1. The concentrated aqueous solution of hydrochloric acid is decom- 
posed by the action of the galvanic current, and the chlorine and 
hydrogen, separated at opposite poles, collected. The electrolysis 
may be made in an ordinary voltameter (Fig. 46). Better adapted to 
this purpose is Hofmann’s apparatus (Fig. 47). Two glass cylinders, 
provided at the top with stop cocks, are connected at the lower end 
with each other and with a funnel tube ; the latter serves to fill the 
apparatus with liquid ; and also, in subsequent additions, to let out the GAS VOLUME. 
65 
gases collected in the tubes. The platinum 
electrodes are fused into the lower part of 
both tubes. In another form of Hofmann’s 
apparatus (Fig. 48), the electrodes are intro- 
duced by means of caoutchouc corks. When 
the separating gases (in this case the chlorine) 
attack the platinum, carbon electrodes are 
substituted for the latter. 
To electrolyze hjdrogen chloride, fill the 
apparatus with concentrated hydrochloric 
acid, which is mixed with ten volumes of a 
saturated salt solution ; close the upper cocks, 
and connect the electrodes with the poles of 
the battery. Gases separate in both tubes, 
and in equal volumes-, that separated at the 
positive pole may be proven to be chlorine ; 
the other combustible gas is hydrogen. 
This experiment shows that hydrogen 
chloride decomposes into equal volumes of 
chlorine and hydrogen. The opposite—the 
production of HC1 by the union of equal 
volumes of H and Cl—appears from the next experiment:— 
Fig. 46. 
Fig. 47. 
2. Fill a cylindrical glass tube, provided with stop-cocks at both 
ends (Fig. 49), with equal volumes of chlorine and hydrogen. This is 
most conveniently done by conducting the gaseous mixture obtained 66 
INORGANIC CHEMISTRY. 
by the electrolysis of HC1 into the dry tube. (The tube should be 
filled in the dark, as the gases combine in daylight.) When the tube 
is filled with the mixture, sunlight or magnesium light is brought to 
bear upon it, when chemical union ensues. On immersing the lower 
end of the tube in water, and opening the lower cock, the water will 
rapidly fill the tube, as it dissolves; all hydrogen and all chlorine have 
disappeared. 
3. The same experiment may be 
applied in the establishment of 
another important fact which has 
reference to the proportion of the 
volume of the hydrogen chloride to 
the volumes of its constituents. If 
the tube filled with equal volumes 
of Cl and H be opened under Hg, 
after the explosion, no diminution 
Fig. 48. 
Fro. 49. 
in volume will be detected, although the mixture of Cl and H has 
been changed to hydrogen chloride. It follows from this that a 
mixture of equal volumes of Cl and H affords the same volume of HC1, 
or, as ordinarily expressed, one volume of Cl and one volume of H 
yield two volumes of hydrogen chloride. 
The following experiment confirms this conclusion : Into a bent 
tube (Fig. 50). filled with Hg, conduct dry HC1 and then introduce in 
the bend of the upper part a little piece of metallic sodium. On 
heating the latter with a lamp, the HC1 is decomposed, the Cl com- 
bining with the Na to form sodium chloride, while free hydrogen is 
separated. Upon measuring the remaining hydrogen it will be found 
that its volume is exactly the half of the volume of HC1 originally in- ATOMS AND MOLECULE 
67 
trodnced. In the same manner may be shown the fact that in two 
volumes of HBr and HI there is contained one volume of H : From 
the densities of Br and I 
vapors follows,further,that the 
quantities of these elements in 
gas form combining with one 
volume of hydrogen also oc- 
cupy one volume. Hence, 
one volume of hydrogen and 
one volume of bromine vapor 
yield two volumes of HBr, 
and one volume of hydrogen 
and one volume of iodine 
vapor two volumes of HI. 
Fig. 50. 
The volume ratios in the chemical union of gases were first investi- 
gated by Humboldt and Gay-Lussac (1805-1808). The latter derived 
the two following empirical laws from the experiments: (1) Gases 
unite according to simple volume ratios; (2) The volume of the 
resulting body stands in a simple multiple proportion to the volumes 
of the constituents. 
Comparing the fact announced by Gay-Lussac, that simple volume 
ratios occur in the chemical union of gases, with that discovered by 
Dalton (p. 60), that the quantities by weight of the combining elements 
also bear a simple ratio, and granting the atomic constitution of 
matter, it follows that the number of smallest gas particles (molecules) 
contained in similar volumes of different gases must bear a simple 
ratio to each other : the simplest supposition, however, would be that 
this number of molecules in equal volumes of all gases is the same. 
These important conclusions were deduced by Avogadro in 1811, and 
by Ampfere in 1814. 
As deduced on page 64 and confirmed by the described 
experiments, the quantities by weight of the halogen-hydrogen 
compounds, expressed by the chemical formulas, HC1, IIBr, 
HI, occupy double as large a volume as one part by weight 
of H or 35,4 parts Cl, 79,7 parts bromine, 126,5 parts iodine. 
While the gas densities of the elements are equal to their 
atomic weights (p. 63) those of the compound bodies, con- 
sequently, amount to half that expressed by their formulas. 
From this it would follow that in equal volumes of com- 
pound bodies only half as many atoms or particles are 
present as in an equal volume of an elementary form of mat- 
ter. In fact, one volume of H, containing n atoms of H, com- 
bines with one volume of chlorine, which, too, contains n 
atoms of Cl. n parts HC1 result, which fill two volumes ; 
therefore, there are only " parts of HC1 contained in one 
volume of HC1:— 68 
INORGANIC CHEMISTRY. 
This conclusion contradicts the general postulate (p. 64), 
derived from the physical properties, viz., that all gases, both 
simple and compound, contain in equal volumes the same 
number of gaseous particles. This contradiction, which for a 
long time prevented the adoption of the atomic volume theory 
in chemical science, is now easily removed by the following 
supposition of Avogadro, announced in 1811. It is necessary 
to distinguish two different kinds of particles: molecules and 
atoms. The smallest discrete particles in gases are not atoms, 
but molecules, which consist of several atoms. That the 
molecules of compounds consist of atoms, is obvious, as, indeed, 
the same represent aggregates of atoms, but the elements also 
form molecules in a free condition, which are composed of 
several, generally, of two atoms. The previously deduced rule 
(p. 63), that in equal gas volumes of the halogen elements 
there is contained the same number of atoms, must be formu- 
lated somewhat as follows : In equal volumes of all gases is 
found the same number of molecules (law of Avogadro). 
1 vol. 
nH + nCl = nHCl. 
1 vol. 
2 vols. 
The process of the combination of hydrogen with chlorine 
(and the other halogens) must, therefore, be conceived to be 
somewhat like the following: 1 molecule of H, containing 2 
atoms of H, acts upon 1 molecule Cl, also composed of two 
atoms of Cl, and there result 2 molecules of IIC1:— 
H2 + Cl 2 = 2HC1. 
Then it is understood, that hydrogen chloride in an equal 
volume contains just as many molecules, as H and Cl, as is 
manifest from the following representation :— 
nHj 
nClj 
nHCl 
nllCl 
1 volume 
1 volume 
2 volumes. 
In similar manner 2 volumes H (containing 2n molecules) give 
with 1 volume oxygen (containing n molecules), 2 volumes 
aqueous vapor ; consequently, 2n molecules of water. In 2n 
molecules of the latter (H20) are contained 2n atoms of 0; 
therefore in n molecules of oxygen, 2n atoms of oxygen—or 
one oxygen molecule consists of 2 atoms. ATOMS AND MOLECULES. 
69 
nil, 
nOa 
Yield 
nll20 
nH,0 
nH, 
1 Vol. 
2 Vola. 
2 Vola. 
In the same way it may be deduced that the nitrogen 
molecule consists of 2 atoms of (N2), the phosphorus molecule, 
of 4 atoms of (P4), etc., etc. 
This peculiar result, following from the law of Avogadro, 
that the molecules of the elements consist of several atoms, 
etc., is shown by many other circumstances founded on facts. 
For example, by the existence of the allotropic modifications 
of the elements (compare ozone), by the chemical reactions 
(compare hydrogen peroxide), by the remarkable action of 
the elements in the moment of their liberation. 
Upon p. 43 we said that the oxygen separated from water 
by chlorine acted much more energetically than free oxygen. 
Other elements, especially hydrogen, behave similarly in 
the moment of formation—in statu nascendi. As viewed by 
the atomic molecular theory, this may be very easily ex- 
plained. The free elements (their molecules) are compounds 
of similar atoms whose chemical affinity has been, of course, 
always partially satisfied. In the moment of their separation 
from compounds free atoms appear, which, before they com- 
bine to molecules, must act more energetically. 
All that has been developed in the preceding statements 
may be summarized in the following sentences: All bodies 
are composed of elementary atoms. The latter unite to pro- 
duce the molecules of the simple and compound bodies. 
Molecules are the smallest discrete particles existing in a free 
state. In equal volumes of all gaseous and vapor-forming 
bodies there is contained the same number of molecules. 
Therefore, the gas densities bear the same ratio to each other 
as the molecular weights. The density is generally compared 
with that of hydrogen = 1, therefore, the gas densities 
(the specific gravities of gases), of all bodies are one-half 70 
INORGANIC CHEMISTRY. 
their molecular weights. The atomic weights are compared 
with—H = 1, therefore, the densities of the elements whose 
molecules consist of two atoms, are equal to the atomic 
weights:— 
Atoms. 
Molecules. 
Density. 
H = 
1 
H2 
= 2 
1 
Cl = 
35.4 
Cl2 
= 70.8 
35.5 
Br = 
79.7 
Br2 
= 159.4 
80 
I = 
126.5 
I2 
= 253 
127 
HC1 
= 36.5 
18.25 
II Br 
= 81 
40.5 
HI 
= 128 
64 
0 = 
16 
o2 
= 32 
16 
h2o 
= 18 
9 
N = 
14 
n2 
= 28 
14 
nh3 
= 17 
8.5 
P = 
31 
p4 
= 124 
62 
ph3 
= 34 
17 
A simpler deduction, that the molecules of the elements consist of 
two or more atoms, is the following: Proceeding from the law of 
Avogadro, that in like volumes of all gases or vapors an equal number 
of molecules is contained, this law, or much better, hypothesis, cannot 
be, as was attempted, proven mathematically, just as little as any other 
fundamental hypothesis.* 
But it possesses, as basis of the entire later mechanical gas theory, 
a high degree of probability. It necessarily follows from this law that 
the molecular weights of all bodies are proportional to the gas densi- 
ties. Referred to hydrogen as unit, the empirical gas densities of HC1 
= 18.2, of HBr = 40.3, of HI = 63.7, etc. Analysis shows, however, 
that 35.4 parts of Cl are in union with 1 partH in HC1, 79.7 bromine in 
HBr, 126.5 iodine in HI. As the weight of one atom of H is made 
equal to 1, and 35.4 parts of chlorine are combined with it, the weight 
of a molecule of hydrogen chloride, as consisting at least of one atom 
of H and one atom of Cl, must equal at least 36.4; it is, therefore, 
twice as much as its density, 18.2. Hence the molecular weights 
also of all other bodies, as they bear the same ratio as the densities, 
must also be twice as large as the (referred to H as unit) latter. 
The hydrogen molecule is = 2, and consists of two atoms, as its 
atomic weight equals 1. The chlorine molecule weighs 70.8 units, and 
consists of two atoms (Cl2), if we suppose, that the atomic weight = 
35.4. Its atomic weight could, however, be only the half (or another 
sub-multiple) of 35.4 ; then its molecule would consist of four chlorine 
atoms (Cl4 = 70.8 when Cl is made equal to 17.6), and the formula 
* A mathematical proof is only possible upon the basis of another, more general, quanti- 
tative hypothesis (or of an axiom), which then on its side is not to he adduced. OXYGEN GROUP. 
71 
of hydrogen chloride would be HC12. From the densities of the ele- 
ments in gas form we only ascertain their molecular weights. Their 
atomic weights are revealed from the molecular weights of their com- 
pounds, as we regard the smallest quantity of the element which analysis 
discloses in the molecule of any compound as the atomic weight. Thus 
in the molecule of any compound of chlorine there is never less than 
35.4 parts by weight of Cl. That the thus derived maximum values 
have not been found too high, but corresponding 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 formed from these numbers. 
Taking one atom of hydrogen as the unit of weight and 
volume, then two parts, by weight,of H, or one molecule (Ha), 
would occupy two volumes. We say, therefore, although 
incorrectly, that the molecules fill two volumes, and designate 
the molecular formulas double volume formulas. The volume 
of molecules and atoms is, however, unknown to us ; we only 
know that in like gas volumes there is contained a like number 
of molecules. 
The totality of the convincing suppositions and conclusions 
deduced from real circumstances forms the atomic-molecular 
doctrine, which is the foundation of the chemistry of to-day. 
As this doctrine completely explains the quantitative phe- 
nomena arising in the action of the chemical elements upon 
each other, inasmuch as they have repeatedly been con- 
firmed by entirely opposite phenomena, it is only proper and 
correct that the doctrine be designated a theory (p. 60). 
OXYGEN GROUP. 
In this group are included the elements, oxygen, sulphur, 
selenium and tellurium. They are perfectly analogous in 
their chemical deportment. They unite with two atoms of 
hydrogen. 
1. OXYGEN. 
Oxygen (oxygenium) is the most widely distributed element 
in nature. Found free in the air, in combination it exists in 
water. It is an important constituent of most of the mineral 
and organic substances. 
0 = 16 *) Oa = 32. 
* The correct atomic weight of oxygen referred to H — 1 is 15.96. 72 
INORGANIC CHEMISTRY. 
It was first discovered, almost simultaneously, by Priestley, 
in England, 1774, and Scheele, in Sweden, 1775, as a re- 
markable body; but Lavoisier, in France, 1774-1781, first 
explained the important role attached to oxygen in processes 
of combustion, of respiration and of oxidation. 
Preparation. Heat red mercuric oxide, a compound of 
mercury with oxygen, in a small glass retort; in this way the 
oxide is decomposed into mercury and gaseous oxygen :— 
HgO = Hg + O. 
The following method is commonly pursued in the chemical 
laboratory. Potassium chlorate, a compound of potassium, 
chlorine and oxygen, is heated in a glass retort (Fig. 51) or 
flask, and thus decomposed into solid potassium chloride and 
oxygen:— 
KC10, = KC1 + 03.* 
Fig. 51. 
The gas evolution proceeds more regularly and requires a 
less elevated temperature if the pulverized chlorate be mixed 
with ferric oxide or manganese peroxide. The liberated 
oxygen is collected over water. 
* The chemical equations used here and previously are only intended 
to represent the manner of the reaction and to express the accom 
panying relative quantities by weight. It should not be forgotten that 
free atoms do not exist, but that they always occur combined in mole- 
cules. Molecularly written the equation would be :— 
2 KC103 = 2 KC1 + 3 02. OXYGEN. 
Very pure oxygen may also be obtained by heating potas 
sium dichromate with sulphuric acid :— 
K2Cr207 -(- 4H2S04 — Cr2(S04)3 -(- K2S04 -(- 4H20 -|- 30. 
Besides these, many other methods may be employed for 
the preparation of the gas : e. g., the ignition of manganese 
and barium super-oxides ; the decomposition of sulphuric acid 
at a high heat; the boiling of a solution of bleaching lime 
with a cobalt salt, etc., etc. These methods, applied techni- 
cally, will be considered more thoroughly later. 
Properties. Oxygen is a colorless, odorless, tasteless gas, 
wrhich is condensed to a transparent liquid, of specific gravity 
0.978, at —130°, under a pressure of 470 atmospheres. Its 
density equals 16 (H = 1)*, or 1.1056 (air = 1). One 
litre of oxygen at 0° C., and 700 mm. pressure weighs 1.4336 
grams. Only slightly soluble in water; 100 volumes of the 
latter dissolve, at 0°, 4.1 volumes; at 15°, 2.9 volumes of oxy- 
gen. More readily dissolved by absolute alcohol (28 volumes 
in 100 volumes). 
Fig. 52. 
Fig. 53. 
Oxygen combines with all the elements excepting fluorine. 
With most of them it unites directly, generally accompanied 
by an evolution of light and heat. The combustion of bodies 
which burn 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 animals are also influenced 
by the oxygen contact of the air—hence the earlier designa- 
tions of oxygen as inflammable air and vital air. In pure 
oxygen the phenomena of combustion, as well as those of res- 
piration, proceed more energetically. Ignited charcoal or a 
* Correct density is 15.96. 74 
INORGANIC CHEMISTRY. 
glowing sliver inflame immediately in oxygen, and burn with 
a bright light. This test serves for the recognition of pure 
oxygen. Sulphur and phosphorus ignited in the air burn in it 
with an intense light (Fig. 52). Even iron is able to burn 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 (Fig. 53). At once the match inflames and ignites the 
iron, which burns with an exceedingly intense light and emits 
sparks. (To protect the vessel from the fusing and glow- 
ing globules of'iron oxide, cover tbe 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. 38). Hydrogen burns in oxygen with 
a flame; vice versa oxygen must also burn in hydrogen ; this may 
be demonstrated in the same manner as indicated under hydro- 
gen chloride (p. 47). A mixture of hydrogen and oxygen 
detonates violently; most strongly if the proportions are 1 
volume oxygen and 2 volumes hydrogen; such a mixture is 
known as oxy-liydrogen gas. The explosibility may be shown 
in a harmless way by the following experiments : Fill a 
narrow-necked flask of 4-6 ounces, over water, f with hydrogen, 
and i oxygen ; close the opening with a cork, then wrap the 
flask up in a towel, remove the cork and bring a flame near 
the opening. A violent explosion ensues, generally with com- 
plete breaking of the flask. 
Fig. 54. 
The oxy-hydrogen flame is only faintly luminous; it possesses, 
however, a very high temperature, answering, therefore, for the 
melting of the most difficultly fusible substances, e. q. plati- 
num. To get a continuous oxy-hydrogen flame, efflux tubes of 
peculiar construction are employed, Fig. 54 ; through the outer 
tube, W, hydrogen is brought from a gasometer ; through the OXIDE. 
inner the oxygen is conveyed, and the mixture inflamed at a. 
Such a flame impinging on a piece of burnt lime makes the 
latter glow and emit an extremely bright light—Drummond's 
Lime Light. 
The union of oxygen with other substances, is termed 
oxidation. This term, as also the name oxygenium (from 
and yewaio), or acid producer, arises from the fact that 
sometimes in oxidation acids are formed. This the previously 
mentioned combustion experiments prove. If the vessels, for 
instance, in which carbon, sulphur and phosphorus were 
burned, be shaken up with water, the latter will give an acid 
taste, and redden blue litmus paper. It was formerly thought 
that the formation of acids is always conditioned by oxygen. 
We have, however, already noticed that the haloid acids HC1, 
HBr, HI, contain no oxygen. Some of the elements yield 
acids by their union with oxygen, or more correctly oxides, 
which with water form acids. Most of these are the metalloids. 
Thus, from the acid-forming oxides of sulphur and phosphorus 
the following corresponding acids are derived:— 
so3 + h2o = h2so4 
Sulphur 
trioxide. 
Sulphuric 
acid. 
P205 + H30 = 2HP03 
Phosphorus 
pentoxide. 
Metaphosphoric 
acid. 
The metals with oxygen mostly yield oxides, which, with 
water, give hydroxides (also hydrates) or bases :— 
K20 + H20 = 2K0H 
Pot. Potas. 
oxide. 
CaO + H20 = Ca (0H)a 
hydroxide. 
Calcium 
oxide. 
Calc. 
hydroxide. 
By the alternating action of acids and bases arise the salts 
(see p. 51). 
Thirdly, there exist the so-called indifferent oxides, which 
yield neither acids nor bases, with H20, e. g. 
N20 
NO 
Ba02 
Oxidation is not only induced by free oxygen or bodies 
rich in it, but, frequently, also by the halogens; in the latter 
Nitrous 
oxide. 
Nitrogen 
oxide. 
Barium 
superoxide. 76 
INORGANIC CHEMISTRY. 
case the halogens first decompose the water with the elimina- 
tion of oxygen, which then oxidizes further (compare p. 42). 
The opposite process to oxidation, the removal of oxygen, 
is called reduction. Hydrogen (in statu nascendi), and sub- 
stances giving it off' easily (as HI), have a reducing action. 
Most of the metallic oxides are, at red heat, reduced by 
hydrogen, e. g.:— 
Copper 
oxide. 
CuO + H2 = Cu + H20 
Copper, 
OZONE, 03. 
Ozone is a peculiar modification of oxygen, characterized 
by a remarkable odor and great ability to react, therefore, 
active oxygen. It is obtained from oxygen in various ways ; 
it is almost always produced when this gas is liberated, or 
when it takes part in a reaction, e. g., in all slow oxidations, 
in every combustion, in the action of electricity upon oxygen 
or air, in the electrolysis of water. In none of these in- 
stances is all the oxygen ever converted into ozone, only a 
small portion—in most favorable conditions 5-6%—suffers 
this change. 
The following methods serve for the preparation of ozone: 
Discovered in 1840, by SchOnbein. 
Fig. 55. 
1. Into a spacious flask bring several pieces of stick phosphorus, 
covering them about half with water, and allowing them to stand for 
some hours. Or conduct oxygen over pieces of phosphorus placed in 
a glass tube and moistened with water. Ozone is formed in large 
quantity when a potassium bichromate solution is substituted for water. OZONI 
2. Pass the electric spark (from an electrical machine or a Ruhm- 
korff’s coil) through air or oxygen. More advantageous is the so-called 
silent discharge from a powerful induction current. For this purpose 
we can employ a Siemen’s induction tube (Fig. 55) which consists of a 
glass tube covered without with tin foil, and in the interior pf which is a 
smaller tube coated within. Between the two tubes the oxygen circu- 
lates ; the two coatings are in connection with the induction spiral, or 
the poles of a Holtz electric machine. 
3. Gradually add barium peroxide in small portions (or potassium 
permanganate) to cold sulphuric acid: — 
Ba02 + H.,S04 = BaS04 + H202 + 0. 
The escaping oxygen is tolerably rich in ozone, and is collected over 
water. 
Ozone possesses a highly penetrating, chlorine-like odor 
(phosphorus odor), which by prolonged respiration produces 
bad results. When heated to 230-300° C., it reverts to ordi- 
nary oxygen. In pure water it is somewhat soluble; the larger 
portion of it is, however, converted by the water into oxygen, 
without formation of hydrogen peroxide. Opposite to 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 sepa- 
rated from potassium iodide by it:— 
2KI + H20 + 0 = 2K0H + I2. 
It also oxidizes all organic substances, like caoutchouc; 
therefore the apparatus used in its preparation must not be 
constructed of the latter. Solutions of dye stuffs, like indigo 
and litmus, are decolorized. Very characteristic for ozone is 
its ability to turn an alcoholic solution of guaiacum tincture 
blue. 
For the detection of ozone the ordinarily so-called potassium iodide 
starch paper may be used. This is. prepared by immersing white tissue 
paper in a starch solution mixed with potassium iodide. The iodine 
which the ozone liberates from the potassium iodide blues the starch 
paper. From the rapidity and the intensity of the coloration the quan- 
tity of ozone may be approximately determined. Besides the potas. 
iodide starch paper, guaiacum tincture and paper saturated with a lead 
acetate solution may be used to detect the ozone ; the first acquires a 
blue color, the second is browned. Other substances also blue potas- 
sium iodized starch and guaiacum, e. (/., chlorine, bromine, nitrogen 
dioxide, etc., etc. To distinguish ozone from these, proceed as follows : 
Take two strips of violet litmus paper, one of which is saturated with 78 
INORGANIC CHEMISTRY. 
KI, and expose it to the action of the gas ; when 03 is present potas- 
sium will be set free from the KI, and color the violet litmus blue. 
The second paper serves to show the absence of ammonia. 
Ozone is formed from pure oxygen, and is nothing more 
than the latter condensed, the molecules of which consist of 
3 atoms of 0 :— 
3 vols. oxygen. 
30 2 yield 20 3. 
2 vols. ozone. 
This is proven by the following experiments : In ozonizing 
oxygen its volume diminishes ; upon heating (whereby ozone 
is again changed to oxygen), the original volume is repro- 
duced ; when ozonized oxygen is brought in contact with oil of 
turpentine or cinnamon, all the ozone is absorbed and the 
volume of the gas diminished. Comparing this diminution, 
corresponding to the ozone volume, with the expansion which 
an equal volume of ozonized oxygen suffers after the applica- 
tion of heat, we will find (according to Soret) that the first is 
twice as large as the latter; this indicates that 1 volume of 
ozone yields II volumes of oxygen. From this it follows that 
the specific gravity of ozone must be H times greater than 
that of ordinary oxygen, and that if the molecule of O con- 
sists of 2 atoms, the molecule of ozone must contain 3 atoms. 
This conclusion is confirmed by the specific gravity of ozone 
experimentally derived from the velocity of diffusion. The 
density of ozone is found to be 24 (H = 1); the molec- 
ular weight of it, therefore, is 24 X 2 = 48; i. e., the mole- 
cule of ozone consists of 3 atoms of O. 
In the action of ozone upon oxidizable bodies like KI and 
Hg, a diminution in the volume of the gas does not occur, 
although all the ozone disappears. From this it would appear 
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 :— 
48 = 3 X 16 = 03. 
1 vol. 
03 -f- 2KI — 02 -(- K20 -f-I2. 
1 vol. 
As a consequence of this behavior, ozone is also called 
oxidized oxygen; i. e., free oxygen (0-2), which has com- 
bined with an additional oxygen atom. 
We observe, therefore, that the elementary substance oxy- 
gen occurs in free condition in two different forms—allotropic 
modifications—ordinary oxygen (02) and ozone (0:i). We OZONE. 
79 
will learn later that very frequently substances of the 
same elementary composition possess different physical and 
chemical properties ; such bodies are called isomer ides and the 
phenomenon isomerism. The isomerism of the elements is 
known as allotropy; this is accounted for (as in the case of 
oxygen and sulphur) by the different number of atoms in the 
molecule. 
The phenomena of isomerism constitute an important argu- 
ment for the atomic constitution of matter. If in the chemical 
union of two bodies the particles of matter would entirely 
permeate and blend into each other, the existence of isomeric 
bodies would be scarcely comprehensible. We can therefore 
only suppose a co-stratification of the atoms, and must con- 
sider isomerism as only a varied arrangement of the same. 
Special allotropy verifies the conclusion drawn from the gas 
density, 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, in consequence, only very slowly oxidized ; 
the ozone is contained in them in a peculiar, combined condi- 
tion. In this form it acts upon some bodies like free ozone; 
in other instances, on the contrary, the oxidizing action is 
only rendered possible by peculiar bodies which carry the 
ozone. Such substances are spongy platinum, ferrous sulphate, 
the blood corpuscles. Thus, for example, old turpentine oil 
containing absorbed ozone only acts on paper saturated with 
starch and potassium iodide, if a few drops of the ferrous 
sulphate solution has been added to it. 
As ozone is formed when electricity acts upon air, and 
indeed, probably, in all oxidation and combustion processes; 
as, further, potassium iodide starch paper is blued when ex- 
posed to the air; it was believed that ozone was a constant 
constituent of atmospheric air (1-10 milligrams in 100 litres 
of air); according to recent investigations it is, however, 
probable that the imagined ozone reactions are frequently 
produced by hydrogen peroxide. 
Antozone, which was regarded as a third peculiar modifica- 
tion of oxygen, probably does not exist. 80 
INORGANIC CHEMISTRY 
COMPOUNDS OF OXYGEN" WITH HYDROGEN 
1. WATER. 
H20 = 18. Density = 9. 
Water, the combination product of hydrogen with oxygen, 
is produced in many chemical processes, as, e. <7., in the forma- 
tion of salts from bases and acids (p. 51). 
Cavendish was the first (1781) to confirm the formation of 
Avater by the combustion of hydrogen. Lavoisier first (1783) 
determined its quantitative composition. Later (1805), Gay- 
Lussac showed that water was produced by the union of two 
volumes of hydrogen with one volume of oxygen. 
Physical Properties. Chemically pure water is obtained 
by the distillation of naturally occurring water, which always 
contains other matter dissolved in it. It appears in all three 
states of aggregation; in the liquid, gaseous (steam), and 
solid (ice, snow). When water is cooled it contracts and 
attains its greatest density at -j- 4° C., the maximum con- 
traction. The weight of a cubic centimeter of such water is 
taken as the unit of weight ( = 1 gram). By further cooling 
the water expands—the opposite of most other bodies ; its 
volume is greater, while the specific gravity decreases. Water 
Is 773 times heavier than air at 0°. 
By further cooling water solidifies to ice. The solidification 
temperature of water, or more correctly the fusing point of ice, 
is taken as the zero point of Celsius’ and Reaumur’s thermome- 
tric scales. We can, however, reduce still water considerably 
below the 0° point without its freezing, whilst the fusing 
point of ice, like all other solid bodies, is constant (at a defi- 
nite pressure). In the conversion of water into ice a consid- 
erable expansion occurs: 100 vols. H20 at 0° yield 107 vols. 
ice; the specific gravity of the latter is, therefore, 9.3. Ice 
crystallizes in hexagonal forms, as is distinctly observed in 
snow flakes. 
The various compounds require different quantities of heat 
to bring them to the same warmth. This capacity for heat is, 
for water, greater than for all other liquid and solid bodies. 
It is customary to take the quantity of heat necessary to 
raise one part by weight of TLO from 0° C., to 1° C., as the 
unit of heat, or calorie. In the passage of a liquid to the 
solid state heat is always set free, while, on the other hand, in 
the fusion of the solid heat is absorbed. The latent heat of WATER. 
81 
water equals 79 calories; that means, that for the fusion of 
one part by weight of ice, a quantity of heat is required 
which is capable of raising one part H20 from 0° to 79° C. 
AVhen heated water boils and is converted into steam, the 
boiling temperature, like that of all other liquids, 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). The temperature of the steam escaping 
from water at the ordinary pressure of 760 mm. is = 100° 
of the thermometric scale of Celsius, = 80° Ileaumur. 
One volume of water, at 100° C., yields 1696 volumes of 
vapor of the same temperature. The specific gravity of 
steam = 0.6234 (air = 1), or = 9 (H = 1). One litre of 
aqueous vapor weighs 0.8064 grams. 
The vaporization of water, and of other liquids, occurs not only at 
the boiling point, but also at lower temperatures. The tension of the 
vapors is measured by the height of the mercurial column. 
The following table gives the tension of aqueous vapor for various 
temperatures :— 
Temperature. 
Tension. 
Temperature. 
Tension. 
—20° C. 
0.93 ram. 
40° C. 
54 9 
—10° C. 
2.09 mm. 
60° C. 
148.8 
0° C. 
4.6 mm. 
80° C. 
354.6 
+ 10° C. 
9.1 mm. 
100° C. 
760.0 
20° C. 
17.4 mm. 
120° C. 
1491.0 
Moist gases, therefore, occupy a larger volume than those which are 
dry. The above table will answer to reduce the observed volume of 
a moist gas to its volume when dry, by deducting from the observed 
atmospheric pressure the tension of steam (in mm.) corresponding 
to the given temperature. 
A definite quantity of heat requisite for the conversion of a 
liquid into vapor is applied to internal and external work ; 
therefore, it disappears as heat, or becomes latent. The latent 
heat of the evaporation of water equals 536.5 heat units at 
100° C.; i. e., for the conversion of one part of water of 100° 
C., into vapor of the same temperature, a quantity of heat 
capable of raising 536.6 parts of H20 from 0° to 1° will be 
absorbed. 
In consequence of the evaporation of water, the gases sepa- 
rating from an aqueous solution are always moist. To dry 
the same, conduct them over such substances as will be able 
to take up the moisture, e. g., calcium chloride, stick potash, 
sulphuric acid, phosphorus anhydride (compare page 33). 82 
INORGANIC CHEMISTRY. 
Many solids abstract moisture from the air without chemically 
uniting with it: to dry these let them stand in an enclosed 
space over sulphuric acid (dessicators). 
The Natural Waters. As water dissolves many solid, liquid 
and gaseous compounds, all naturally occurring waters con- 
contain foreign admixtures. The purest natural water is rain 
and snow water; it contains upwards of three vols. per- 
centage of gases (oxygen, nitrogen and carbon dioxide), and 
traces of solids (the ammonium salts of nitrous and nitric 
acids). If water that has stood exposed to the air be heated, 
the dissolved gases escape as bubbles. 
River and spring waters contain, on an average, in 10,000 
parts, from 1-20 parts of solid constituents. Water having 
present in it much lime and gypsum, is ordinarily known as 
hard, in distinction to soft water, which contains very little 
of these impurities. On boiling lime waters, most of the 
impurity deposits out. Spring water generally contains in 
addition larger quantities of carbon dioxide, which imparts 
to it a refreshing, enlivening taste. Spring waters holding 
considerable quantities of solid constituents or exhibiting par- 
ticular healing properties, are called mineral waters. These 
are distinguished as saline waters (containing sodium chloride), 
bitter waters (containing magnesium salts), sulphur waters 
(hydrogen sulphide), acidulated waters (saturated with carbon 
dioxide), chalybeate waters (containing iron), and others. 
Sea water contains about 3.5 °jo of salts, of which 2.7 % are 
sodium chloride. 
To purify the natural waters they are filtered (for the 
removal of mechanical admixtures), and for chemical pur- 
poses, distilled (distilled water), in apparatus of varying form. 
Solutions. The phenomena appearing in the dissolving 
of substances indicate that solutions are not mere mechan- 
ical mixtures. In every solution alterations in the tem- 
perature of the liquid occur. The solubility of solid and 
liquid substances increases usually with the temperature, while 
that of gases diminishes. The quantity of dissolved gas is fre- 
quently proportional to the temperature; other gases, on the 
contrary, which are readily soluble in water, such as the halo- 
gen-hydrogen compounds, are exceptions to this rule. Heat 
does not perfectly expel them from their solution ; they distill WATER. 
83 
over as liquids of definite composition (compare p. 56). 
When they dissolve, a large quantity of heat is liberated, just 
as in the case of chemical compounds. Further, in the solu- 
tion of solids and liquids a contraction is always perceived ; 
the volume of the solution is less than the sum of the volumes 
of the constituents. These phenomena point to the acceptance 
of a certain affinity between the dissolving bodies. Therefore, 
solutions, like alloys, are designated undetermined compounds, 
in contrast to the determined compounds, which are com- 
bined according to constant atomic weight ratios. Among 
others, this view is confirmed by the fact that frequently 
definite compounds of bodies with water do exist in solution. 
Such compounds often separate, unaltered, upon evaporation 
of the solution ; the water present in them is known as water 
of crystallization. It is, however, impossible to draw a sharp 
line between determined and undetermined compounds, 
between chemical and physical attraction. 
The thermal phenomena appearing when solution occurs bear a close 
relation to chemical affinity. The hydrogen haloids, easily soluble in 
water, disengage in their solution large quantities of heat correspond- 
ing to the symbols— 
(HG1, aq.) = 17320; (HBr, aq.) = 19940; (HI, aq.) = 19200 
This liberation of heat is explained by the production of the hydrates 
HCl + 8H20, HBr -+- 6H20, HI -f- 5H20, which distill over unal- 
tered (p 56). The slightly soluble, so-called permanent gases do not 
form such hydrates, and when they dissolve, disengage but little heat. 
The liquid and solid bodies exhibit a like deportment. These, form- 
ing hydrates, liberate heat, while the non-hydrate-forming solid 
bodies absorb heat in their solution, which at the same time is em- 
ployed to liquefy them (latent heat of fusion). Thus, in the solution of 
the halogen compounds of potassium the following quantities of heat 
are absorbed :— 
(KC1, aq.) = — 4400; (KBr, aq.) == — 5080; (KI, aq.) = 5100. 
Upon such an absorption of heat depend the freezing mixtures, to be 
described later. 
Chemical Properties of Water. Water is a neutral sub- 
stance, i. e., it possesses neither acid nor basic properties. As 
we have already observed (p. 75), it forms bases with basic 
oxides and acids with acid-forming oxides. 
Despite the fact that the affinity of hydrogen for oxygen is 
so great, water* may, however, be decomposed by many sub- 
stances. At ordinary temperatures, metals like K, Na and 
Ca decompose it, with liberation of hydrogen :— 
2H20 fK,= 2K0H + H,. 84 
INORGANIC CHEMISTRY. 
Other metals do not decompose it, except at elevated tem- 
peratures. Steam conducted over glowing iron gives its 
oxygen to the latter, forming ferroso-ferric oxide, while hydro- 
gen is set free :— 
HHO 
3Fe + HHO = Fe304 + 4H2. 
HHO 
HHO 
Chlorine in sunlight decomposes water ; the decomposition 
is more rapid when the vapors are conducted through heated 
tubes:— 
The galvanic current separates* water into its constituents, 
oxygen and hydrogen ; the first collecting at the positive, 
and the latter at the negative pole. The oxygen thus obtained 
contains ozone. 
H20 + Cl 2 = 2HC1 + 0. 
The deportment of water at high temperature is very inter- 
esting. On pouring molten platinum into cold water, bubbles 
of oxy-hydrogen gas escape. A similar decomposition of water 
occurs when it is led through porcelain tubes raised to a white 
heat. 
Fig. 56. 
From the observations of Sainte-Claire Deville, it was found 
that the decomposition of II20 begins at 1200° C., and that 
it increases, with a rising temperature, and is complete at 2500° 
C. Such a partial decomposition increasing with the tempera- 
* Pui-e water appears not to be decomposed by the galvanic current. WATER. 
ture is known as dissociation. The following experiment 
illustrates this. Through a porous clay tube, a, puttied into a 
wider, non-permeable porcelain tube, heated to white heat in 
an oven, pass aqueous vapor, Fig. 56. The water suffers par- 
tial decomposition, the lighter hydrogen, which passes through 
into the porcelain tube more rapidly than the oxygen, escapes 
through the gas tube b. The oxygen escapes, for the most 
part, through the inner tube at a. A part of the same diffuses 
simultaneously with the hydrogen and the latter. 
To avoid this conduct a stream of carbon dioxide through the 
wider porcelain tube; this will carryout the hydrogen with it. 
The carbon dioxide will be absorbed by the alkali solution in 
the collecting vessel, and oxy-hydrogen gas be found in the 
cylinder. The quantity of the gas increases with the tempera- 
ture. A platinum tube may be advantageously substituted 
for the porous clay tube, because only hydrogen will pass 
through it (p. 36). 
Many other compounds, like ozone, ammonium chloride, 
phosphorus pentacldoride, carbon dioxide, etc., suffer a simi- 
lar partial decomposition at a glowing heat; they are altered 
to simpler molecules by heat. 
The explanation of the dissociation phenomena is found in the 
kinetic theory of gases and heat. According to it, not only have the 
gas molecules a direct oscillating movement, inasmuch as they, like 
elastic balls, rebound from each other, but even the atoms in the mole- 
cule possess heat vibrations. The velocity'of the oscillations of mole- 
cules and atoms increases with augmented temperature ; it is, there- 
fore, understood that by a determined energy of the oscillations the 
chemical affinity is overcome and the united atoms are separated from 
each other. 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; the former are warmer than 
the latter. Only the sum of the existing forces of all the molecules is 
a constant quantity at every temperature. The more highly heated 
molecules, whose number increases with the temperature, therefore, 
yield to the decomposition. From this we discover that the dissocia- 
tion is gradual and increases with the temperature. The law of dis- 
sociation is expressed by the curve of probability. 
Dissociation, i. e., the partial decomposition, increasing with 
the temperature, explains many chemical processes, previously 
appearing enigmatical, e. g., the mass action in reversed 
chemical reactions. We have already said that glowing iron 
decomposed water with the separation of hydrogen and pro- 
duction of ferrous-ferric oxide. Ou conducting H over 86 
INORGANIC CHEMISTRY. 
glowing iron oxides the opposite process occurs; the oxygen 
compound of the iron is reduced and water is formed :— 
Fe304 + 4H2 = 3Fe + 4H20. 
In the first instance the excess of water acts. Some of its 
molecules are dissociated ; oxygen combines with iron, while 
the liberated H 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. In the action of the bodies in an 
enclosed space at a given temperature there must occur a 
state of equilibrium, in which Fe304, Fe, H20 -f- II2 occur 
simultaneously. Such a state occurs in every dissociation. 
The Quantitative Composition of Water. The Atomic Weight 
of Oxygen. 
The composition by weight of water is best determined by 
a synthesis of the same, by reducing cupric oxide with 
hydrogen:— 
CuO + H2 = Cu + h2o. 
Cupric oxide. 
Copper. 
Fig. 57. 1VATER. 
87 
Heat a weighed portion of cupric oxide (containing a 
definite amount of oxygen), in a stream of pure, dry hydro- 
gen, and weigh the quantity of H20 obtained. The operation 
can be executed in the apparatus represented in Fig. 57. 
The H generated in the flask A is washed in B and then 
dried in the tubes C, D and E, 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 completely absorbed in the tube H. Hydro- 
gen is led over the cupric oxide until it is reduced to red 
metallic copper, then allow to cool, and weigh F alone and G 
and H together. The loss in weight of F expresses the 
quantity of oxygen which, combined with H, produces water. 
The increase in weight of G and II gives the quantity of water 
produced. The difference shows the amount of H in water. 
Thus we ascertain that in 100 parts by weight of water there 
are— 
11.11 Parts Hydrogen. 
88.81 “ Oxygen. 
100.00 “ Water. 
or, 1 part hydrogen and 8 parts oxygen yield 9 parts water 
The Molecnlar Formula of Water. Atomic W eight of Oxygen. 
If the molecule of water (like HC1) contains 1 atom H and 
1 atom oxygen, then its chemical formula would be HO, and 
the atomic weight of oxygen wTould be = 8. However, such a 
supposition has not been in any way proved. It would be 
just as likely that the formula H02 might be ascribed to the 
water molecule; then the atomic weight of oxygen would be 4. 
According to the formula H20 the atomic weight of O is 16, 
etc. (see p. 62). The analytical data give no decision. For 
the determination of the actual atomic weight of oxygen, and 
therefore also the number of atoms in the molecule, we must 
direct our attention to the views presented on pages 62-71. 
In equal volumes of the gases (or vapors) there is an 
equal number of molecules. The molecular weights, there- 
fore, are proportional to the gas densities, and are equal 
to double that of the densities referred to H = 1. The 
density of steam is 9 (H = 1); the weight of the water 88 
INORGANIC CHEMISTRY. 
molecule is therefore 18. Analysis, however, shows that in 
18 parts water 2 parts by weight are hydrogen (= 2 atoms) 
and 16 weight parts oxygen. According to this the molecule 
of water contains not more nor less than 2 atoms of hydrogen. 
That the 16 parts oxygen combined with the latter correspond 
to one atom (that the atomic weight does not equal 8, in which 
case the molecular formula of water would then be H202), 
follows from the fact that the analysis of none of the innumer- 
able oxygen derivatives has shown less than 16 parts oxygen 
in the molecule (see p. 70). The molecular formula of water, 
therefore, is H20 — 18. The gas density of oxygen is 16, the 
molecular weight 32, therefore the oxygen molecule consists of 
2 atoms 02 = 32. 
After having thus derived the molecular formula of water 
and the atomic weight of oxygen, we deduce the following 
conclusions: (1) 16 parts by weight of 0 occupy the same 
volume as 1 part by weight of H, as 16 parts of the former 
unite with 2 parts of the latter; from this, in the formation of 
water 1 volume of 0 must combine with 2 volumes of II. 
(2) In equal volumes the same number of molecules is con- 
tained; n molecules 0 unite therefore with 2 n molecules of 
hydrogen (H2) ; the same yield 2 n molecules water; conse- 
quently 2 volumes of aqueous vapor:— 
nH2 + nH2 + n02 = nH20 + nH20. 
According to the above, 2 volumes of hydrogen and 1 volume 
oxygen condense in their union to 2 volumes of aqueous vapor. 
The same result follows from the gas density of water. 
As 1 volume of steam weighs 9, 2 weight parts of H, however, 
form with 16 parts by weight of O, 18 parts by weight of 
steam; then this quantity of water in form of vapor must 
occupy two volumes. Conversely from these volume ratios 
it is shown that the molecule of oxygen consists of two atoms. 
These conclusions are confirmed by the following experi- 
ments :— 
1 vol. 
1 vol. 
1 vol. 
1 vol. 
1 vol. 
1. When water is decomposed by the electric current in a voltame- 
ter, or more suitably, in Hofmann’s apparatus (Fig. 48, p. 66), it will 
be found that the volume of the separated hydrogen is double that of 
the oxygen. This can also be proven synthetically. Into an eudiome- 
ter tube filled with mercury (see Air), introduce 1 volume oxygen and 
2 volumes hydrogen, and let the electric spark pass through the mix- 
ture. This will unite the two gases, a small quantity of water forming 
at the same time ; all gas has disappeared, and the tube fills perfectly WATER. 
89 
with mercury. In place of the 
eudiometer the following apparatus 
may be advantageously employed 
in this experiment (and also in 
many others Fig. 58). It consists 
°f a V-shaped glass tube, one limb 
ot which, open above, is provided 
below with an exit tube. The 
other limb really represents an eu- 
diometer; it is divided into cubic 
centimeters, having two platinum 
wires fused into the upper end, and 
provided with a stop-cock to let out 
the gases, and thus test them. Fill 
the tube to the stop-cock with mer- 
cury, and run into the eudiometer 
limb 1 volume O and 2 volumes II. 
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 formed water existing as 
aqueous vapor, it is only necessary 
after the explosion to convert it, by 
heat, into steam. The subjoined 
Fiq. 58. 
Fig. 59. 90 
INORGANIC CHEMISTRY. 
apparatus will answer for this purpose (Fig. 59). This is essentially 
the same as that pictured in Fig. 32, 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. (aniline). These, 
then, pass through the envelope B, and are again condensed in the 
spiral tube C. The quantities of IT and 0 used are heated to the 
same temperature, their volume noted, the explosion produced, and 
the volume of the resulting aqueous vapor determined. From this it is 
found that the volume of hydrogen is f of the volume of the gas 
mixture ; that thus from 3 volumes of oxy-hydrogen gas, 2 volumes 
aqueous vapor result. 
From the specific gravity of hydrogen and of oxygen, and from the 
volume ratios according to which they combine, the composition by 
weight, of water, is easily deduced :— 
1 volume of oxygen weighs 16 weight parts 
2 volumes ofhydrogen weigh 2 weight parts 
the resulting H20 weighs 18 weight parts. 
18 parts water, therefore, contain 16 parts oxygen and 2 parts hydrogen, 
or in 100 parts there are 88.88 parts oxygen aud 11.12 parts hydrogen. 
2. HYDROGEN PEROXIDE. 
II2 02 = 34 
In addition to water oxygen forms another compound with 
hydrogen, known as hydrogen peroxide. It is produced by 
the action of dilute acids upon certain peroxides, such as 
those of potassium, calcium and barium. It is most con- 
veniently obtained by the action of hydrochloric acid upon 
barium peroxide:— 
Ba02 + 2HC1 = BaCl 2 + H202. 
Barium peroxide made to a paste with a little water (better the 
hydrate—sec Barium) is introduced gradually, in small quantities, into 
cold hydrochloric acid, diluted with three volumes of water. Hydro- 
gen peroxide and barium chloride, both soluble in water, result. To 
remove the second from the solution, add to the latter a solution of 
silver sulphate, as long as a precipitate is formed. By this reagent two 
compounds, barium sulphate and silver chloride, perfectly insoluble in 
water, are produced:— 
Barium 
peroxide. 
Barium 
chloride. 
Remove the precipitate by filtration and concentrate the aqueous solu- 
tion, now containing only hydrogen peroxide, under the air pump. 
In making the peroxide carbon dioxide may be allowed to 
act on barium peroxide suspended in water:— 
BaCl2 + Ag2S04 — BaS04 + 2AgCl. 
BaO 2 -f C02 + H20 = BaCOs + H202. 
The insoluble barium carbonate is filtered off and the filtrate 
concentrated. HYDROGEN PEROXIDE. 
91 
Hydrogen peroxide is most practically obtained by adding 
moist barium hydrated peroxide (see Barium) to cold dilute 
sulphuric acid. The reaction occurs according to the following 
equation :— 
Ba02, H20 + H2S04 =BaS04 -f H202 + H20. 
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 hydrate solution, then 
concentrate the filtrate. Dry commercial hydrate of the per- 
oxide of barium is not applicable for the above. 
Further, worthy of remark, is the formation of hydrogen 
peroxide from water and ozone aided by ether. Ozone is 
conducted into the latter, and the ozonized liquid shaken with 
pure water, which dissolves the peroxide that has been pro- 
duced. 
Besides these methods, all dependent on the decomposition of per- 
oxides, others exist for preparing hydrogen peroxide (although in 
much smaller quantity). Thus, also, it arises in almost all slow oxida- 
tions in which ozone is produced at the same time. This simultaneous 
production of hydrogen peroxide and ozone is explained to some extent 
by the composition of these bodies. The oxygen molecule splits, in its 
chemical action, into two atoms of oxygen ; one combines with the 
water, the other with free oxygen, to yield ozone:— 
02 + H20 02 — H202 -f- 03. 
Thus, in the oxidation of moist phosphorus in the air, oxygen is 
ozonized, while hydrogen peroxide appears in the water. When a 
glowing platinum spiral is hung in a flask containing some water and 
ether, it will continue to glow, as the ether vapors burn slowly ; ozone 
is then perceptible in the air of the flask—hydrogen peroxide, how- 
ever, in the water. On shaking zinc filings or zinc amalgam with 
water, hydrogen peroxide may be detected in the latter. Oxygen 
obtained by electrolysis contains ozone, and the water some hydrogen 
peroxide. Bain water almost invariably contains hydrogen peroxide, 
which appears to form by the action of sunlight. (Scheme.) 
The latter, concentrated as much as possible under the air 
pump, is a colorless, syrupy liquid, with a specific gravity of 
1.45, and does not solidify at —30° C.; from mere dilute 
solutions pure water freezes out. It possesses a bitter, astrin- 
gent taste, miscible in all proportions with water, and vapor- 
izes in vacuo. 
In concentrated solutions hydrogen peroxide is very un- 
stable, and easily decomposed with liberation of oxygen ; in 
more dilute acidulated solutions it may be preserved longer. 
Decomposition occurs, even at ordinary temperatures ; by heat- 92 
INORGANIC CHEMISTRY. 
ing the point of explosion can be reached. In consequence of 
this ready decomposition, hydrogen peroxide oxidizes power- 
fully, since oxygen appears (p. 69) in statu nascendi. It 
converts selenium, chromium and arsenic into their corres- 
ponding acids ; sulphides are changed to sulphates (PbS to 
PbSO<) ; from lead acetate solutions the superoxide is precipi- 
tated ; organic dyestuffs are decolorized and organic tissues 
decomposed. From hydrogen sulphide sulphur, from hydro- 
gen chloride and iodide, chlorine and iodine are set free: 
H202 + 2HI = 2H20 + Ia. 
Thus h)rdrogen peroxide acts in a manner analogous to 
ozone ; in both there exists a slightly bound 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, ferrous sulphate, or blood corpuscles (see p. 79) be 
added to the solution, iodine immediately separates out and 
colors added starch-paste deep blue. 
In all these cases the action of hydrogen peroxide is oxidi- 
zing. Some substances, on the other hand, are reduced by 
H202, oxygen separating at the same time ; this is true of cer- 
tain unstable oxides, peroxides, and the highest oxidations of 
some metals. Thus argentic oxide, mei’curic oxide and gold 
oxide are reduced to a metallic state with an energetic evolu- 
tion of oxygen :— 
Lead peroxide is changed to lead oxide:— 
Ag20 + H202 — 2Ag + H20 + 02 
Pb02 “l- H202 = PbO -f- H20 -(- 02 
In presence of acids, the solution of potassium permanganate 
is decolorized and changed to a manganous salt. In the same 
way chromic acid and its salts are altered to chromic oxide :— 
2Cr03 -f- H202 = Cr203 + H20 02. 
Ozone and hydrogen peroxide decompose themselves into 
water and oxygen :— 
03 -(- H202 — 02 H20 -j- 02. 
All these reactions can only be explained by supposing that the 
oxygen atoms (also those of other elements), possess a certain affinity 
for each other. Those present in other compounds, therefore, sepa- 
rate and unite with each other, and form free oxygen molecules 
—00. The conclusion derived from the gas density, viz., that the 
molecules of the free elements consist of two or more atoms, HYDROGEN PEROXIDE. 
93 
is corroborated by these reactions. The unstable union of the one 
atom of oxygen in ozone and in hydrogen peroxide also finds expres- 
sion in their thermo-chemical behavior. In the production of ozone 
from oxygen, as also of hydrogen peroxide from water and oxygen, 
heat is absorbed :— 
(02,0) = — 29600 (H20,0) = — 11200 
Both compounds are endothermic (see p. 57), therefore, little stable, 
and the one oxygen atom in them is readily capable of reaction. 
The production of both substances, according to the above sj'mbols, 
can also only be effected by the addition of external energy. In case 
of ozone this may be accomplished by electricity, with hydrogen perox- 
ide (in its production from barium superoxide and dilute acids) by 
heat, which in the formation of the barium salts becomes free. 
Third and finally, hydrogen peroxide suffers decomposition 
into water and oxygen by many bodies, especially when the 
latter exist in a divided condition, and they themselves are 
not in the least altered. Gold, platinum, silver, manganese 
peroxide, carbon and others, act in the above way. Such 
reactions, in which the reacting substances undergo no per- 
ceptible changes, are designated catalytic. In many cases these 
may be explained by the previous formation of intermediate 
products, which subsequently react upon each other. Thus, 
we can suppose that in the action of silver and gold upon 
H202 oxides at first result, which afterwards are reduced in 
the above mentioned manner, by the hydrogen peroxide. 
H202 decomposes potassium iodide very slowly, but in the presence 
of iron sulphate, however, iodine at once separates, and is recognized 
by the blue color it yields with starch paste. In the same way, in 
the presence of ferrous sulphate guaiacum tincture is at once colored 
blue and an indigo solution decolorized. The most characteristic test 
for the peroxide is the following: Introduce some H202 into a chromic 
acid solution, add a little ether and shake thoroughly ; the 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 delicate reagent which yields an orange yellow color 
with hydrogen peroxide. 
Quantitatively hydrogen peroxide is determined by oxidation with 
potassium permanganate (see Manganese b The latter is added to the 
solution, and acidified with sulphuric acid until a permanent coloration 
occurs. The reaction takes place according to the equation :— 
2Mn04K + 3S04H2 + 5H202 = 2S04Mn + S04K2 -f 8H20 + 502 
Or the liquid to be examined (rain water) for hydrogen peroxide is 
shaken in a stoppered glass with a five per cent, solution of potassium 
iodide and some starch paste, allowed to stand several hours, and the 
separated iodine determined, colorimetrically (Schone). 
Reactions for the Detection of Hydrogen Peroxide. 94 
INORGANIC CHEMISTRY. 
Hydrogen peroxide occurs in slight quantity in the air and is 
detected in almost every rain water frequently, also in snow—but not 
in natural dew and frost. Its quantity varies from 0.05 to 1 milli- 
gram in a litre of rain. Its formation in the air is probably induced 
by the action of ozone upon ammonia, whence ammonium nitrite, 
hydrogen peroxide and oxygen result (Carius). 
Analysis shows that II202 consists of one part hydrogen and 
sixteen parts of oxygen ; its simplest formula will, therefore, 
be HO. The difficult volatility of the compound, as well as the 
above mentioned reactions, let us conceive that the molecule 
of hydrogen peroxide is more complicated, and is expressed 
by H202. It is supposed that the peroxide is composed of 
two groups of OIT, called hydroxyl, combined with each other. 
S = 32 S = 64 (above 1000° C.). S6 = 192 (at 500° C.). 
2. SULPHUR. 
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 gypsum, 
calcite and marl. The compounds of sulphur with the 
metals are known as blendes or glances. In combination 
with oxygen and calcium, sulphur forms calcium sulphate, the 
widely distributed gypsum. It is also present in many organic 
substances. 
To obtain it, the natural sulphur in Sicily, arranged in 
heaps and covered with earth, is melted or distilled from 
earthen retorts. This commercial, crude sulphur, for further 
purification, is (in the manufactory) again distilled from cast- 
iron retorts, and when molten, run into cylindrical forms— 
stick sulphur. If the sulphur vapors during distillation are 
rapidly cooled (which occurs by conducting them into a stone 
chamber through which cold air circulates), they condense to 
a fine yellow powder, known as flowers of sulphur (Flores sul- 
phuris). 
By heating the well known pyrites (FeS2) sulphur may be 
obtained. 
Free sulphur forms several allotropic modifications (see 
page 79). 
1. Ordinary octahedral or rhombic sulphur exists in nature 
in beautiful, well crystallized rhombic octahedra (Figs. 21 and 
22, p. 27). It is pale yellow, hard and very brittle; on 
rubbing, it becomes negatively electrified. Specific gravity 
2.05. Difficultly soluble in alcohol and ether ; more readilv SULPHUR. 
95 
soluble in hydrocarbons and ethereal oils. The best solvents 
are sulphur monochloride (S2C12) and carbon disulphide (CS2); 
100 parts of the latter at 22° C., dissolve 46 parts sulphur. 
By slow evaporation of the solutions sulphur crystallizes in 
transparent, lustrous, rhombic octahedra, like those occurring 
naturally. Sulphur fuses at 111.5° C., (113° C.) to a yellow, 
mobile liquid, which upon further heating becomes dark and 
thick, and at 250° C., is so viscid that it cannot be poured from 
the vessel holding it. Above 300° C., it again becomes a thin 
liquid, boils at 440° C., and is converted into an orange yellow 
vapor. 
2. The prismatic or monoclinic sulphur results from the 
rhombic when the latter is heated to its point of fusion ; gener- 
ally on cooling it assumes the monoclinic form (sulphur 
heated too high crystallizes at 90°C., in rhombic forms). The 
monoclinic crystals are best obtained as follows: Fuse sul- 
phur in a clay crucible, allow it to cool slowly until a crust 
appears on the surface; break this open near the side and 
pour out the portion yet in a liquid state. The walls of the cru- 
cible 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 
disulphide, in a sealed tube, is heated to 100° C., and then 
gradually allowed to cool; monoclinic crystals at first sepa- 
rate, and later, at low temperatures, rhombic octahedra. The 
monoclinic crystals separated from the solution are almost 
colorless and perfectly transparent. 
From the supersaturated solution of sulphur in benzene prismatic 
or octahedral crystals may be obtained as desired, by bringing some 
of the same into the liquid. 
This form of sulphur has a specific gravity of 1.96 and fuses 
at 120°. Soluble in the same solvent as the rhombic. It is very 
unstable; the transparent prisms and needles at ordinary 
temperatures become opaque and pale yellow, specifically 
heavier (heat is evolved), and pass over into an aggregate of 
rhombic octahedra retaining the external prismatic form. 
Stick sulphur deports itself similarly; the freshly moulded 
sticks are composed of monoclinic prisms, but in time their 
specific gravity changes and they are converted into the rhom- 
bic modification. 
3. Soft, plastic sulphur appears to consist of two modifica- 
tions. It is obtained when sulphur melted above 230° is 96 
INORGANIC CHEMISTRY. 
poured in a thin stream into water; it then forms a soft, fusible 
mass, of a yellowish-brown color. In a few days it hardens, 
and is converted into rhombic sulphur. At 95° the conver- 
sion is instantaneous and accompanied by the evolution of 
considerable heat. It is only partly soluble in carbon disul- 
phide, leaving behind an amorphous powder—amorphous in- 
soluble sulphur—insoluble in carbon disulphide. Also pro- 
duced when light acts upon dissolved or fused sulphur, and 
in the decomposition of the halogen-sulphur compounds by 
1LO. Flowers of sulphur are for the most part insoluble in 
carbon disulphide. 100° C., will convert the amorphous insol- 
uble sulphur again into the ordinary variety. 
On adding hydrochloric acid to polysulphide solutions of 
potassium or calcium, sulphur separates as a fine, white 
powder, known as milk of sulphur (Lac sulfuris) :— 
K2S5 + 2HC1 = 2KC1 + H2S + 4S. 
This is amorphous, soluble in carbon disulphide, and gradually 
passes into the rhombic form. 
The existence of these various modifications of sulphur, like 
ozone, may be attributed to the presence of a varying number 
of atoms in the molecules. This supposition is confirmed by 
the deportment of sulphur vapor. The vapor density of 
sulphur at 500° C., has been found to be equal to 96 (H = 
1). With increase of temperature, from 700° C., up, the 
vapor density steadily diminishes, becoming constant at 900° 
C., and equalling 32; the molecular weight, therefore, is 64. 
As now the atomic weight of S (as we will see) = 32, it 
follows that at 1000° C., the molecules of S consist of two 
atoms (S2 = 64 = 32 X 2). At 500°, however, when the 
vapor density = 96, molecular weight equalling 192, the 
molecule consists of six atoms (S6 = 6 X 32 = 192). 
According to this the hexatomie sulphur molecules dissociate 
(see p. 85) on further heating, and fall into normal 
diatomic molecules; the dissociation beginning at 700° and 
being complete at 1000° C. As, then, the sulphur molecules 
in vapor formed at very high temperatures consist of two atoms 
and at lower of six atoms, we can take for granted that the 
molecules in liquid and solid condition are more complicated, 
and that the various allotropic modifications of sulphur are 
influenced by the number of atoms contained in the molecules. 
Just like sulphur, other solid metalloids, e. g., selenium, phos- HYDROGEN SULPHIDE. 
97 
phorus, arsenic, carbon, silicon, occur in different modifica- 
tions. Up to this time we have no means of ascertaining the 
size of the molecular elements in liquid and solid conditions; 
there is much, however, which favors the idea of these 
elements in free condition consisting of complex atomic 
groups. 
Chemical Properties. In its chemical behavior sulphur is 
very similar to oxygen, and its compounds have the same 
constitution as the corresponding oxides. It unites directly 
with most of the elements. Heated to 260° in the air, it 
inflames and burns with a pale bluish flame, giving sulphur 
dioxide (S02). Nearly all the metals combine with sulphur 
to form sulphides. By tubbing mercury, flowers of sulphur 
and water together, black mercury sulphide is obtained. 
A moist mixture of iron filings and sulphur glows after a time. 
Cu and Fe burn in sulphur vapor. The sulphides are 
analogous to the oxides, exhibit similar reactions, and possess, 
for the most part, a similar composition, as is seen from the 
following formulas:— 
H20, Water. 
KOH, Pot. hydrate. 
BaO, Barium oxide. 
CO 2, Carbon dioxide. 
C03K2, Pot. carbonate. 
H2S, Hydrogen sulphide. 
KSH, Pot. sulphydrate. 
BaS, Barium sulphide. 
CS2, Carbon disulphide. 
CS3K2, Pot. sulpho-carbonate. 
COMPOUNDS OF SULPHUR WITH HYDROGEN. 
1. HYDROGEN SULPHIDE. 
H2S = 34. Density = 17. 
In nature hydrogen sulphide occurs principally in volcanic 
gases and in the so-called sulphur waters. It is always pro- 
duced in the decomposition of organic substances containing 
sulphur, and also in the reduction of alkaline sulphates by 
decomposing carbon compounds. It may be formed directly 
from its constituents, although in small quantity, if hydro- 
gen gas be conducted through boiling sulphur, or if sulphur 
vapors, together with, hydrogen, be conducted over porous 
substances (pumice stone, bricks) heated to 500° C. Many 
sulphides are reduced upon ignition in a stream of hydrogen, 
with separation of hydrogen sulphide :— 
Ag2S -f H2 = 2Ag + H2S. 98 
INORGANIC CIIEMISTRV 
For its production acids are allowed to act upon sulphides. 
Ordinarily iron sulphide with diluted sulphuric acid is 
employed; the action occurs at ordinary temperatures:— 
FeS + H2S04 = FeS04 + H2S. 
The operation is performed either in Kipp’s apparatus (p. 
33) or in the one pictured in Fig. 30. 
Hydrogen sulphide thus obtained contains admixed hy- 
drogen, in consequence of metallic iron existing in the sul- 
phide. Perfectly pure hydrogen sulphide is obtained by 
heating antimony sulphide with hydrochloric acid :— 
Sb2S3 + 6HC1 = 2SbCls + 3H2S. 
Properties. Hydrogen sulphide is a colorless gas, having 
an odor similar to that of rotten eggs; inhaled in large 
quantities it has a stupefying effect, and is very poisonous. At 
medium temperatures it condenses under a pressure of 14 
atmospheres (under ordinary pressure at— 74°) to a colorless 
liquid of specific gravity 0.9, which at — 85° C., solidifies to a 
white crystalline mass. Its density equals 17 (H = 1) 
or 1.177 (air = 1). Water dissolves 3-4 times its 
volume of gas, the solution possessing all the properties of 
gaseous hydrogen sulphide, and hence called hydrogen sul- 
phide water. 
Ignited in the air the gas burns with a blue flame, water 
and sulphur dioxide resulting :— 
With insufficient air access, or when the flame is cooled by 
the introduction of a cold body, only hydrogen burns, while 
sulphur separates out in a free condition. In aqueous solu- 
tion hydrogen sulphide is decomposed by the oxygen of the 
air at ordinary temperatures, sulphur separating as a fine 
powder:— 
H2S + 30 = h2o + so2. 
For this reason hydrogen sulphide becomes turbid upon 
exposure to the air. 
The halogens behave like oxygen; the hydrides of the 
halogens are formed with separation of sulphur :— 
H2S + 0 = H20 + S. 
This reaction serves for the production of hydrogen iodide 
(p. 54). 
As hydrogen sulphide has a great affinity for oxygen, it 
withdraws the latter from many of its compounds, hence 
H2s + I2 = 2H[ + S HYDROGEN SULPHIDE. 
99 
acting as a reducing agent (p. 76). Thus chromic, man- 
ganic and nitric acids are reduced to lower stages of oxida- 
tion. On pouring fuming nitric acid into a dry vessel con- 
taining hydrogen sulphide, the mixture will ignite with a 
slight explosion. 
Hydrogen sulphide possesses weak acid properties, red- 
dens blue litmus paper, forms salt-like compounds with bases, 
and is, therefore, termed hydrosulphuric acid. Nearly all 
the metals liberate hydrogen from it, yielding metallic sul- 
phides :— 
Pb + H2S = PbS + H2. 
AV'ith the oxides and hydroxides of the metals H2S yields 
sulphides and sulphydrates:— 
KOH + H2S = KSH + H20 
CaO + H2S = CaS + H20. 
Sulphides, therefore, like the compounds of the halogens 
with the metals, may be viewed as the salts of hydrosulphuric 
acid. The sulphides of almost all the heavy metals are in- 
soluble in water and dilute acids ; therefore they are precipi- 
tated by fLS from solutions of metallic salts:— 
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. Paper saturated with a lead 
solution is at once blackened by II2S, lead sulphide being 
formed—sensitive test for H,S. 
CuS04 + H2S = CuS + II2S04. 
Molecular Formula of Hydrogen Sulphide. Atomic Weight of 
Sulphur. 
The analysis of hydrogen sulphide shows that it consists of one part 
hydrogen and sixteen parts sulphur. If the molecular formula of 
hydrogen sulphide were HS, the atomic weight of sulphur would be six- 
teen (compare p. 87). The great analogy of the sulphur compounds, 
with those of oxygen (p. 97), permits us to accept a formula for the 
former similar to those of thedatter. The molecular formula of hydro- 
gen sulphide would, therefore, be H2S, and the atomic weight of sul- 
phur would equal 32. Hence the gas density of hydrogen sulphide must 
be 3/ =17 (H = 1), or 1 177 (air= 1), which is also confirmed by direct 
experiment. Conversely, it follows from the gas density that the molecu- 
lar weight of hydrogen sulphide = 34. As, then, according to analysis, 
there are in thirty-four parts of the same, two parts hydrogen, the mole- 
cule will contain two atoms of hydrogen. That the thirty-two parts, 
sulphur combined with them correspond to one atom of sulphur 100 
INORGANIC CHEMISTRY. 
follows, because in the molecule of no sulphur compound have there 
been found less than thirty-two parts of that element (see p. 87). 
From the molecular formula H2S, further, we conclude that the 
contained hydrogen in one volume hydrogen sulphide, in free condi- 
tion, occupies the same volume :— 
nH2S contains nH2. 
This conclusion is experimentally verified as follows: In a bent 
glass tube filled with mercury (Fig. 60), introduce dry hydrogen sul- 
phide gas; then in the bent portion place a piece of tin, which is heated 
by a lamp. The sulphur of the H2S, now 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 sulphur, 32 parts, 
in vapor form, at 1000 C., 
when the density is 32 (p. 
96) combined with hydro- 
gen (2 parts) will equal ex- 
actly half the volume of the 
hydrogen ; at 500° C., how- 
ever, when the vapor density is three times as great, it will equal 1 
volume of the hydrogen. 1 volume H2S, therefore, consists of £ 
volume H and £ volume sulphur vapor (at 500°), or as ordinarily 
expressed, 2 volumes H2S consist of 2 volumes H and J volume 
sulphur vapor. Molecularly written, we have:— 
1 vol. 
1 vol. 
Fig. 60. 
At 500°C. : S6 + 6H2 = 6H,S. 
1 vol. 
6 vols. 
6 vols. 
At 1000° C., however : S2 -f- 2H2 = 2H2S. 
1 vol. 
2 vols. 
2 vols. 
2. HYDROGEN PERSULPHIDE. 
Just as hydrogen peroxide is formed by the action of acids 
upon some superoxide, so hydrogen persulphide may be 
obtained from metallic persulphides. Calcium persulphide is 
most suitable, and when its aqueous solution is poured into 
aqueous hydrochloric acid, 
CaS2 + 2HC1 = CaCl2 + H2S2, 
a yellow, oily, disagreeable liquid, insoluble in water, sepa- 
rates. It decomposes gradually at medium temperatures, more 
rapidly on warming, into hydrogen sulphide and sulphur:— 
It is generally thought that H2S2 is analogous to H202, and 
is hydrogen disulphide containing excess of sulphur dissolved 
h2s2 = h2s + s. SULPHUR 
101 
in it. As, however, the calcium supersulphide used repre- 
sents a mixture of CaS2, CaS3 and CaS5; it is probable that 
the oily liquid is a mixture of H2S2, H2S3 and H2S5. We 
must at least conclude that H_,S3 is present in it, because it 
unites with strychnine to form a crystalline compound. 
COMPOUNDS OF SULPHUR WITH THE HALOGENS. 
Sulphur and chlorine unite to form three compounds: 
SC12, SC14 and S2C12. 
Sulphur Dichloride—SC12—produced when S2C12 is satu- 
rated with chlorine in the cold :— 
S2C12 + Cl2 = 2SC12. 
The excess of chlorine is removed by conducting a stream of 
CO 2 through it. 
A dark, red colored liquid, with a specific gravity 1.62; 
boils at 64° C., with partial decomposition into S2C12 and Cl2; 
the dissociation commences at ordinary temperatures. 
Fio. 61. 
Sulphur Tetrachloride—SC14—only exists at temperatures 
below 0° C. It is formed by saturating SCI. with Cl at 30° 
C., and readily decomposes into SC12 and Cl2; the dissociation 
commences at — 20° C., and is complete at -f- 6°. It yields 
crystalline compounds with some chlorides, e. g., SnCl4, AsC13, 
SbCl3. 102 
INORGANIC CHEMISTRY, 
The most stable of the sulphur chlorides is 
Sulphur Mono-chloride—S2C12—formed when chlorine is 
conducted over molten sulphur. (Fig. 61.) It distills over 
and condenses in the receiver E; the product is redistilled, to 
purify it. 
Sulphur mono-chloride is a reddish-yellow liquid with a 
sharp odor, provoking tears, having a specific gravity of 1.68, 
and boiling at 189° C. Its vapor density equals 67 (H = 1) 
corresponding to the molecular formula S2C12 = 135. Fumes 
strongly in the air, and by water is decomposed into sulphur 
dioxide, sulphur and hydrochloric acid :— 
Sulphur mono-chloride readily dissolves sulphur and serves in 
the vulcanization of caoutchouc. Bromine forms analogous 
compounds with S. S2Br2 is a red liquid, boiling at 190-209° 
C. Iodine, upon gentle warming, unites with S to form S2I2. 
2S2C12 + 2H20 = S02 + 4HC1 + 3S 
Se = 79.5. Se2 = 159 (at 1400° C.). 
8. SELENIUM. 
This element occurs rarely in nature, and then only in very 
small quantities, principally in certain iron pyrites (in Sweden 
and Bohemia). In roasting this ore of iron for the prepara- 
tion of sulphuric acid selenium settles out in the chimney dust 
or in the deposit of the lead chambers (compare Sulphuric 
Acid), and was found in this by Berzelius, in the year 1817. 
Like sulphur, selenium forms different allotropic modifi- 
cations. Amorphous selenium, obtained by the reduction of 
selenium dioxide (Se02) by means of sulphur dioxide (S02), 
is a reddish-brown powder, soluble in carbon disulphide, with 
a specific gravity 4.26. From carbon disulphide selenium 
crystallizes in brownish-red crystals. From the brownish-red 
solution of potassium selenide on standing exposed to the air, 
leafy black selenium crystals separate, which are isomorphous 
with sulphur. Upon suddenly cooling fused selenium it so- 
lidifies to an amorphous, glassy, black mass, with a specific 
gravity of 4.28, and is also soluble in carbon disulphide. When 
selenium (amorphous) is heated to 97° C., its temperature sud- 
denly rises above 200° C.; it is converted into a crystalline, 
dark gray mass with a specific gravity 4.8. It possesses 
metallic lustre, conducts electricity, and is insoluble in carbon TELLURIUM. 
103 
disulphide. The crystalline, insoluble modification is obtained 
by slowly cooling the molten selenium. 
Selenium melts at 217°, and boils toward 700°, converting 
itself into a dark yellow vapor. The vapor density diminishes 
regularly with increasing temperature (similar to sulphur), 
becoming constant at 1400°C. Then it equals 79.5; the 
molecular weight is, therefore, 159, i. e. the molecule of 
selenium at 1400° C., consists of two atoms (2 X 79.5 = 159.) 
Selenium is a perfect analogue of sulphur. In the air it 
burns with a reddish blue flame, forming Se02 and emitting 
a peculiar odor resembling rotten horse-radish. In concen- 
trated sulphuric acid selenium dissolves with a green color, 
forming selenious acid. 
Hydrogen Selenide — H2Se — produced like hydrogen 
sulphide, is a colorless, disagreeable smelling gas with 
poisonous action. In the air the aqueous solution becomes 
turbid and free selenium separates. 
With chlorine selenium forms SeCl4 and Se2Cl2 perfectly 
analogous to the sulphur compounds; SeCl4 is a solid and 
sublimes without decomposition. 
4. TELLUKIUM. 
Of rare occurrence, either native or in combination with 
metals; with gold and silver in sylvanite, with silver and lead 
in altaite. Occurs principally in Transylvania, Hungary, 
California, Virginia, Bolivia and Brazil. 
The tellurium precipitated by sulphurous acid from a 
solution of tellurous acid (see this) is a black powder of 
specific gravity 5.928. 
According to its physical properties tellurium is a metal. 
It is silver white, of a perfect metallic lustre, and conducts 
electricity and heat. It crystallizes in rhombohedra, having 
a specific gravity 6.25. Fuses at 500° and vaporizes in a 
stream of hydrogen. Heated in the air it burns, with a 
bluish-gray flame, to tellurium dioxide (Te02). 
Te = 128* 
* The atomic weight of tellurium, formerly taken as = 128, has 
been determined to be 126.8; from the law of periodicity, to be 
presented later, it, however, seems, with much certainty, that the 
same will yet be found to be somewhat lower (indeed, some less than 
that of iodine — 126.5). 104 
INORGANIC CHEMISTRY. 
The vapor density of tellurium at 1380° C., has been 
discovered to be about 126, corresponding to the molecular 
formula Te2. 
Hydrogen telluride H2Te, is a colorless, very poisonous 
gas, with disagreeable odor. With chlorine tellurium yields 
two compounds TeCl2 and TeCl4; with bromine it gives 
TeBr2 and TeBr4. 
Summary of the Elements of the Oxygen Group 
The elements oxygen, sulphur, selenium and tellurium form 
a natural group of chemically similar bodies. Especially 
marked is the similarity of the last three, while oxygen, 
possessing 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 latter, the elements of the oxygen group 
present a gradation in their properties corresponding to their 
atomic weights :— 
Atomic weights, 
0. 
16 
S. 
82 
Se. 
79.5 
Te. 
126.8. 
With the increase in the atomic weight there occurs a simul- 
taneous condensation of substance, the volatility diminishes, 
while the specific gravity, the point of fusion and boiling 
increase, as seen in the following table :— 
Oxygen. 
Sulphur. 
Selenium. 
Tellurium. 
Specific gravity, 
1.95-2.07 
4.2-4.8 
6.2. 
Melting point, 
111.5° 
217° 
500°. 
Boiling point, 
420° 
700° 
White heat. 
Gas density, 
16 
32 
79.5 
Oxygen is a difficultly coercible gas, while the others, at 
ordinary temperatures, are solids. We must, however, bear in 
mind that sulphur, selenium and tellurium in free state are 
probably composed of larger complex atomic groups (see p. 96 ). 
Further, with rising atomic weight the metalloidal passes 
into a more metallic character. Tellurium exhibits the phys- 
ical properties of a metal; even selenium in its crystalline 
modification possesses metallic properties. In chemical 
deportment, however, the metalloidal character shows scarcely 
any alteration. All four elements unite directly, at elevated 
temperatures, with two atoms of hydrogen, to form volatile 
gaseous compounds having an acid nature; only the oxy- OXYGEN GROUP. 
105 
gen derivative—water—is liquid at ordinary temperatures and 
shows a neutral reaction. At a glowing heat the hydrogen 
compounds are decomposed into their elements. The affinity 
of hydrogen for oxygen is greatest ; therefore, H2S, H*Se and 
H2Te are decomposed in aqueous solution by the air. 
A measure for the chemical affinity of the elements of the oxygen 
group is afforded, as in the case of the halogens (p. 57), by their 
thermo-chemical deportment. By the union of 2 parts hydrogen with 
16 parts oxygen to steam of 100° (H2,0—vapor), 57200 calories are 
disengaged. By the condensation of steam to water of 100° 9600 addi- 
tional (= 18 X 536.5) calories become free (latent heat of evaporation, 
p. 81), and by the cooling of water for each 1 degree C., further also 
18 calories ; so, on the whole, by the production of 1 molecular weight of 
water from its elements, 69000 calories* are disengaged. 
The heat disengagement is less in the formation of hydrogen sulphide, 
while in case of hydrogen selenide it is even absorbed, corresponding 
to the symbols :— 
(H2,0—vapor) = 57200 (H2,S) = 4600 (H2,Se =) = — 5400. 
The previously mentioned fact, viz., the decomposition of H2S, 
H2Se, and also of H2Te, by oxygen, is then explained according to the 
idea of the greatest liberation of heat. We see, consequently, that 
here, as also with the halogens, the chemical affinity of the homologous 
elements successively diminishes with increasing atomic weight (de- 
crease of negative character). It must, however, be borne in mind, 
that the heat liberation represents no direct measure for the affinity of 
the free atoms. If, for example, water be produced, corresponding to 
the molecular equation 2H2 + 02 = 2H20, the hydrogen and oxygen 
molecules must first be broken up into individual atoms, for which 
a definite quantity of heat is necessary (heat of decomposition). The 
directly observed heat disengagement of 57200 calories only indicates 
that the energy of hydrogen to oxygen (211,0) is so much greater than 
the energy of the elementary molecules. 
2H,0 > H,H + 0,0. 
As the solid sulphur molecules consist of a greater number of atoms, 
it is probable their heat of decomposition is greater than that of the 
diatomic oxygen molecules, so that the energy of 2H to S certainly is 
more than 4600 calories. From this decomposition of the molecules 
standing in connection with the absorption of heat, is explained why 
in every reaction, for its introduction and continuation, definite thermal 
conditions are necessary ; that, for example, sulphur and hydrogen 
unite first at more elevated temperature, and then only partially. 
That, indeed, in the breaking of the molecules, heat is absorbed, 
follows, among others, from the fact that, in the process of combustion, 
in nitrogen oxide more heat will be set free than in oxygen (02). 
All the energy masses derived from thermal data are, therefore, only 
* 683C0 calories, according to Tliomsen. 106 
INORGANIC CHEMISTRY. 
relative ; it is only in the most recent times that it has become possible 
to calculate the heat of decomposition of the carbon molecules, and, 
consequently, the real chemical energy of the carbon atoms (Thomsen). 
NITROGEN GROUP. 
Here belong nitrogen, phosphorus, arsenic, antimony and 
bismuth. The latter possesses a decidedly metallic character. 
With three atoms of hydrogen these elements, bismuth ex- 
cepted, yield gaseous derivatives. 
1. NITROGEN. 
Free, it exists in the air, of which it constitutes f and oxy- 
gen the remaining £. Combined, it is principally contained 
in the ammonium and nitric acid compounds, as well as in 
many organic substances of the animal kingdom. 
To isolate nitrogen from the air the latter must be deprived 
of its second constituent. This is effected by such bodies as 
are capable of ab- 
sorbing oxygen with- 
out acting upon the 
nitrogen. It is most 
simply brought about 
by the combustion of 
phosphorus. Several 
pieces of the latter 
are placed in a dish 
swimming on water, 
then ignited, and a 
glass bell jar placed 
over them (Fig. 62). 
In a short time, when 
all the oxygen is 
absorbed from the air, the phosphorus will cease burning ; the 
phosphorus j)entoxide produced dissolves in water, and the 
residual gas consists of almost pure N; its volume will equal 
four-fifths of the air taken. Another procedure consists in 
conducting air through a glowing tube filled with copper 
turnings; the copper unites with the oxygen and pure nitro- 
gen escapes. At ordinary temperatures the removal of O from 
the air may be accomplished by the action of phosphorus, 
a solution of pyrogallic acid and other substances. 
N = 14. Na = 28. 
Fig. 62. NITROGEN. 
107 
A very convenient course for the direct preparation of 
nitrogen is the following. Heat ammonium nitrite in a small 
glass retort; this decomposes the salt directly into water and 
nitrogen:— 
In place ofammonium nitrite a mixture ofpotassium nitrite and NH4C1 
may be used ; by double decomposition upon warming, these salts yield 
Kotassium chloride and ammonium nitrite (KN02 + NH4C1 = NH4 
r02 -j- KCl), which decomposes further. As potassium nitrite usually 
contains free alkali, to combine the same, some potassium bichromate is 
added. Practically the solution consists of 1 part potassium nitrite, 1 
part ammonium nitrate, and 1 part potassium bichromate, in 3 parts 
water, and then boiled ; to free the liberated nitrogen from every trace 
of oxygen, the gas is conducted over glowing copper. 
The action of chlorine upon aqueous ammonia produces 
nitrogen. The chlorine combines with the H of the ammonia, 
forming HC1, which, with the excess of NH3 forms ammonium 
chloride. The nitrogen that was in combination with the 
hydrogen is set free. The following equations express the 
reactions:— 
NH4N03 =N2 + 2H20. 
and 
2NH3 + SC12 = N2 + 6HC1 
6HC1 +■ 6NH3 = 6NH4C1 
Ammonium 
chloride. 
The apparatus pictured in figure 36, page 40 will serve 
to carry out the experiment. The disengaged chlorine is 
conducted through a Woulff wash bottle containing ammonia 
water, the free nitrogen being collected over water. 
In this experiment the greatest care should be exercised that 
an excess of chlorine is not conducted into the solution, because 
its action upon the ammonium chloride will cause the form- 
ation of an exceedingly explosive body (nitrogen chloride, 
NC13) separating in oily drops. 
Properties. Nitrogen is a colorless, odorless, tasteless gas, 
which condenses at —130° and a pressure of 280 atmospheres. 
Its density =14 (H = 1) or 0.9695 (air = 1). Water dis- 
solves about 2% by volume. In its chemical deportment it is 
extremely inert, combining directly with only a few elements, 
and entering chemical reaction but slowly. It does not 
support combustion or respiration; a burning candle is 
extinguished in nitrogen; animals are suffocated in it. This 
is not due to the activity of the N, but to absence of O—a 
substance which cannot be dispensed with in combustion and 108 
INORGANIC CHEMISTRY. 
respiration. The presence of N in the air moderates the 
strong oxidizing property of the pure oxygen. 
THE ATMOSPHERE. 
The air, or the envelop encircling the earth, consists 
principally of a mixture of nitrogen and oxygen ; in addition 
it always contains, but in slight and variable quantities, 
aqueous vapor, carbon dioxide and traces of other substances, 
as accidental constituents. The pressure exerted by the air is 
measured by a column of mercury which holds it in a state of 
equilibrium ; the height of the barometric column at the sea 
level and 0° C., equals, upon an average, 700 millimeters. As 
1 c.c. of mercury weighs 13.596 grams, 76 c.c. will equal 
1033.7 grams, and the last number would indicate the pressure 
which the column of air exercises upon one square centimeter 
of the earth’s surface. 
1 c.c. air weighs (at 0° C. and 760 mm. pressure), 0.001295 
grams; 1000 c.c., therefore, or one litre, would weigh 1.295 
grams. As one litre of H20 weighs 1000 grams, air is 
consequently 773 times lighter than water. Air is 14.46 
times heavier than hydrogen. The specific gravities of the 
gases and vapors were formerly referred to air (= 1); 
compared with H = 1, they are, therefore, 14.46 times greater 
than before. 
Remarks. From these data, with the aid of the specific gravity 
derived from the molecular weights, the absolute weight of definite 
volumes of all gases may be readily determined, a problem frequently 
presented for solution in practice. One litre of air weighs 1.295 grams, 
one litre of hydrogen 0.0896 grams. To ascertain the weight of a 
litre of any other gas or vapor, its specific gravity referred to air = 
1 must be multiplied by 1.295, or if compared with H = 1 by the factor 
0.0896. 
History. In ancient times air, like fire and water, was regarded as 
an element. In the beginning of the seventeenth century it became 
known that by combustion and respiration in an enclosed space a por- 
tion of the air disappeared, and that the part remaining was no longer 
suitable for the support of the above processes; hence to it was 
ascribed the name destroyed air; to the first fire air. In the second 
half of the eighteenth century Scheele, in Sweden, and Priestley, in 
England, found that when, to the so-called destroyed air (nitrogen), a 
certain amount of the gas, set free by heating mercuric oxide (oxygen), 
was added, a mixture resulted possessing all the properties of atmos- 
pheric air. Although thus both constituents of air were separately 
won and by their mixture air again regenerated, yet at that time THE ATMOSPHERE. 
109 
views regarding the nature of both ingredients and the nature of com- 
bustion and oxidation processes prevailed which, throughout, were 
perfectly false. It was believed that combustion and oxidation were 
destructive processes; that the combustible and oxidizable bodies 
enclosed within themselves a peculiar substance, called phlogiston, 
which, in combustion and oxidation, escaped as fire and heat (phlo- 
giston theory of Stahl, 1723). These erroneous opinions were ex- 
plained and corrected by Lavoisier (in 1774) by the following cele- 
brated experiment bearing upon the composition of the air: A glass 
sphere, provided with 
a long, twice bent 
neck (Fig. 63), was 
filled with a weighed 
quantity of mercury. 
The open end of the 
neck dipped into a 
mercury trough, PS, 
and was closed com- 
pletely by a glass bell 
jar. Then the balloon 
A was heated for some 
days at a temperature 
near the boiling point 
of mercury. By this 
means the mercury 
absorbed the oxygen 
of the air contained 
in A and the bell jar 
P, forming mercuric 
oxide. In the course 
of several days, when, upon further heating, no additional decrease 
in the volume of air was observable, the experiment was interrupted 
and the residual gas volume in A and P measured. Upon com- 
paring this with the volume before the experiment, it was discovered 
that vol. of the air had disappeared and combined with the 
mercury to red mercuric oxide. Lavoisier now strongly ignited 
the thus obtained mercuric oxide, and obtained a volume of oxy- 
gen equal to that withdrawn from the air during the experiment. 
By mixing this with the residual volume of N the original vol- 
ume of air was again recovered. Thus it was demonstrated that 
air consists of £ volumes N and volume 0 gas. The elemen- 
tary character of nitrogen was first established by Lavoisier, in 1787. 
It was called azote (from life and a privative), by him. The 
symbol derived from azote, is used in France and England for 
nitrogen. The name nitrogenium (from which the symbol N) was 
given to nitrogen because it was a constituent of saltpetre (nitrum). 
The above experiment served Lavoisier for another important 
deduction As he determined the weight, both of the employed mer- 
cury and the resulting mercuric oxide, he discovered that the weight 
increase was exactly equal to that of the oxygen withdrawn from the 
Fig. 63. 110 
INORGANIC CHEMISTRY. 
air, and by glowing the mercuric oxide the same weight of oxygen 
was again separated. Thus was it demonstrated that the process of 
oxidation was the union of two bodies (not a decomposition), *and 
that the weight of a compound body equals the sum of the weights of 
its constituents ; the principle of the indestructibility of matter. 
Quantitative Composition of Air. Its composition is ex- 
pressed by the quantity of oxygen and nitrogen contained in 
it, as its remaining admixtures are more or less accidental and 
variable. Boussingault and Dumas determined the accurate 
weight composition of the air by the following experiment: 
A large balloon, with a capacity of about 20 litres (Fig. 64), 
is connected with a porcelain tube a b filled with metallic 
copper. Balloon and tubes, closed by stop-cocks, are pre- 
viously emptied and weighed apart. The bent tubes A, B 
and C contain KOH and sulphuric acid, and serve to free the 
air undergoing analysis from aqueous vapor, carbon dioxide 
Fig. G4. 
and other impurities. The porcelain tube, filled with copper, 
is heated to a red heat, and by carefully opening the stop- 
cocks ft, 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 Cu, forming 
cupric oxide, so that only pure nitrogen enters V. Now close 
the cocks and weigh the balloon and porcelain tube. The 
increase in weight of the latter represents the quantity of THE ATMOSPHERE. 
111 
oxygen in the air; the increase in V the quantity of nitrogen. 
In this manner we find that in 100 parts by weight, of air, 
there are contained— 
Nitrogen 76 87 parts by weight. 
Oxygen 23.13 “ “ “ 
100.00 “ “ “ 
As we know the specific gravity of nitrogen (14) and of 
oxygen, we can readily calculate the volume composition of 
air from that, in parts by weight, and can also determine its 
specific gravity (14.46). 
Directly, the volume composition 
of air may be found by means of 
the absorptiometer. The latter is 
a tube carefully graduated, and 
sealed at one end. This is filled 
with mercury, and air, whose vol- 
ume is determined by reading off 
the divisions on the tube, permitted 
to enter. Now introduce up into 
the tube, through the mercury, a 
platinum wire having a ball of 
phosphorus attached to the end 
(Fig. 65), (or a ball of coke satu- 
rated with an alkaline solution of 
pyrogallic acid). The phosphorus 
absorbs the oxygen of the air, and 
only nitrogen remains, whose vol- 
ume is read off by the graduation. 
The eudiometric method affords greater accuracy. It is 
dependent upon the combustion of the oxygen with hydrogen 
in an eudiometer. The latter is an absorptiometer, having 
two platinum wires fused in its upper end (Fig. 66). Into 
the eudiometer are introduced 100 volumes of air and 100 
volumes of hydrogen, and the electric spark then passed 
through the wires ( Fig. 67). All the oxygen in the air com- 
bines with a portion of the hydrogen to form water. On 
cooling, the aqueous vapor condenses and a contraction in 
volume occurs. Upon now measuring the volume of the 
residual gas, with regard to all corrections (see below), we 
will discover that the same equals 137.21 volume units ; con- 
sequently, of the 200 employed volumes of gas, 62.79 volumes 
Fia. 65. 112 
INORGANIC CHEMISTRY. 
Fig. 66. 
disappeared in the formation of water. As water 
results from the union of 1 volume of oxygen and 2 
volumes of hydrogen, therefore, the 100 volumes of air 
employed in the analysis contained~ = 20.93 volumes 
of oxygen. From this the air consists of 
79.07 volumes Nitrogen, 
20.93 “ Oxygen. 
100.00 “ Air. 
Fio. 67. 
Measuring Gases. The volume of gases is influenced by 
Pressure, temperature, and the moisture contained in them. 
'he volume of dry gases, at 760 mm. barometric pressure 
and 0° C., is accepted as the normal volume. If a gas has 
been measured under any other conditions, it must be reduced 
to the normal volume. According to the law of Boyle and 
Mariotte, the volumes of the gases are inversely proportional 
to the pressure ; therefore, if the volume of the gas at pres- 
sure h has been found equal to V, its volume by 760 mm. equals IT. 
According to Gay-Lussac’s law, all gases expand in proportion to 
the temperature. Their coefficient of expansion is — 0.003665 ; 
i e., one volume of gas at 0° occupies at 1° the volume 1 000365. If 
Yt represents the observed gas volume at t°, Yo, however, its volume at 
0°, then 113 
THE ATMOSPHERE. 
Vt 
Vo = *, 
1 + 0.003665.t 
and, considering the pressure, 
Vh 
760(1 + 0.003665. t) 
Further, the gas volume is enlarged by moisture, as the tension of 
the aqueous vapor opposes the atmospheric pressure. The moisture 
may be removed by introducing into the gas a ball of coke saturated 
with sulphuric acid, which dries it. More convenient, however, is to 
make the correction 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; now measure the gas and 
calculate its normal volume from the above formula, after deducting 
from the observed pressure h the number of millimeters corresponding 
to the tension of the aqueous vapor for the given temperature (p. 81). 
While formerly it was supposed that everywhere upon the 
earth’s surface, as also in the highest attainable media, the air 
contained the same relative quantity of oxygen and nitrogen, 
it is at present proven, by accurate investigations, that varia- 
tions to 0.5 °Jo occur. These are dependent upon the wind 
currents; and it also appears that in the tropics a stronger 
oxygen absorption occurs than in the polar regions (v. Jolly). 
From this invariability in composition one might conclude that 
air is a chemical compound of oxygen and nitrogen. 
This supposition is, however, opposed by the following 
circumstances. All chemical compounds contain their constitu- 
ents in atomic quantities, which in the case of air is not so. In 
the mixing of nitrogen and oxygen, to form air, neither the sepa- 
ration nor absorption of heat occurs, which is always observed 
in chemical compounds. Further, the air absorbed by water 
or other solvents possesses a composition different from the 
atmospheric; this depends upon the dissimilar solubility of 
nitrogen and oxygen in water. The air expelled from water 
by heating consists of 34.9 volumes oxygen and 65.1 volumes 
nitrogen (Bunsen). From this, air is not a chemical com- 
pound, but a mechanical mixture of its two constituents. 
The great constancy in composition of the air depends on the mutual 
diffusion of the gases. As the gas molecules possess a direct pro- 
gressive movement, they distribute themselves, without limitation, into 
space, and intermingle regularly among each other. The velocity of 
* Vo = Vt — Vo. 0.00366t, consequently Vo + Vo, 0.00366 -f- = Vt and Vo (1 + 
0.00366t) = Vt. 114 
INORGANIC CHEMISTRY. 
the diffusion of gases is approximately inversely proportional to the 
square root of their densities—the law of the diffusion of gases. The 
density of hydrogen = 1; the density of oxygen = 16 ; therefore, 
hydrogen diffuses 4 times more rapidly than oxygen. The unequal 
diffusion of gases may be per- 
ceived if they are allowed to 
pass through very narrow 
apertures, or through porous 
partitions. The following ex- 
periment very clearly illus- 
trates this: In the open end 
of an unglazed clay cylinder 
(as used in galvanic elements) 
there is puttied a glass tube 
about one meter long, its open 
end terminating in a dish 
containing water (Fig. 68) ; 
the cylinder and tube are filled 
with air. Over the porous 
cylinder is placed a wider 
vessel filled with hydrogen. 
The latter presses almost four 
times faster into the cylinder 
than the air escapes from it; 
the air in the tube and cylin- 
der are displaced and rise in 
the water in bubbles. When 
the separation of gas ceases, 
tube and cylinder are almost 
filled with pure hydrogen. 
On removing the larger hy- 
drogen 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 weaker than the 
external, and water ascends 
in the glass tube. 
Fro. 68. 
In addition to N and O, air constantly contains aqueous 
vapor and carbon dioxide in very small quantities. The 
presence of the former can readily be recognized by the 
fact that cold bodies in moist air are covered with dew. Its 
quantity depends on the temperature and corresponds to the 
vapor tension of water (see p. 81). 1 c.c. of air perfectly 
saturated with aqueous vapor contains 22.5 grams water at 
25° C. ; on cooling to 0° 17.1 grams of these separate as rain. 
Generally the air contains only 50-70 % of the quantity of vapor AMMONIA. 
115 
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 deter- 
mined. 
To detect the carbon dioxide in the air, conduct a portion 
of the latter through solutions of barium or calcium hydrates, 
and a turbidity will ensue. To determine its quantity, pass a 
definite and previously dried amount through a weighed cal- 
cium hydrate tube, and ascertain its increase in weight. 
10,000 parts of atmospheric air contain, ordinarily, from 2-6 
parts carbon dioxide. 
Besides the four ingredients just mentioned, air mostly con- 
tains, although in small quantities, ozone, hydrogen peroxide, 
and ammonium salts (ammonium nitrite). Finally, air, 
especially in its lower layers, contains microscopic germs of 
lower organisms, whose presence influences the processes of the 
decay and fermentation of organic substances. 
COMPOUNDS OF NITROGEN WITH HYDROGEN. 
NIIS = 17. Density = 8.5. 
AMMONIA 
Ammonia occurs in the air in combination with some 
acids, in natural waters and in the earth, only, however, in 
small quantities. The formation of ammonia by the direct 
union of nitrogen and hydrogen occurs under the influence of 
the silent electric discharge. Its compounds are frequently 
produced under the most varying conditions. Thus, by the 
action of the electric spark upon moist air ammonium nitrate 
is formed:— 
N2 + 0 + 2H2 = NH4 N03 
Ammonium 
nitrate. 
By the evaporation of water in the air small quantities of 
ammonium nitrite result:— 
N2 + 2H20 = NH4N02 
Am. nitrite. 
The same salt is formed in every combustion in the air; by 
the rusting of iron and in the electrolysis of water. The 
white vapors which moist phosphorus forms in the air, consist 116 
of ammonium nitrite. Further, ammonium salts are produced 
in the solution of many metals in nitric acid, in consequence 
of a reduction of the acid by the liberated hydrogen :— 
INORGANIC CHEMISTRY. 
HN03 + 4H2 = 3H20 + NHS. 
A fruitful source for ammonia production is the decompo- 
sition of nitrogenous organic substances and their dry distilla- 
tion. Formerly, uj) to the preceding century, the bulk of 
the technically most important ammonium salt, ammonium 
chloride, was obtained by the distillation of camel’s dung 
(in Egypt in the oasis of Jupiter Ammon—hence the name 
Sal ammoniacum). At present ammonia is almost exclu- 
sively prepared by the dry distillation of stone coal, as a 
bye-product in illuminating gas, by combining ammonia sepa- 
rating from the latter with sulphuric or hydrochloric acid. 
Fig. 69. 
To prepare ammonia heat a mixture of ammonium chloride 
and slaked lime in a glass or iron flask :— 
Ammonium 
chloride. 
2NH4C1 + Ca(OH)2 = CaCl2 + 2H20 + 2NHS. 
Calcium 
hydrate. AMMONIA. 
117 
The disengaged ammonia gas is collected over mercury, as 
it is readily soluble in water (Fig. 69). For perfect drying 
conduct it into a vessel filled with burnt lime (CaO). Cal- 
cium chloride is not applicable here for drying, as it enters 
into combination with the gas. In consequence of its levity, 
ammonia, like hydrogen, may be collected by displacing the 
air in inverted vessels. 
Physical Properties. Ammonia is a colorless gas with a 
suffocating, characteristic odor. Its density is 8.5 (H = 1), 
or 0.591 (air =1). Under a pressure of 6.5 atmospheres (at 
10°C.), or by cooling to — 40° C., it condenses to a colorless, 
mobile liquid, with a specific gravity of 0.613 at 0°, which 
solidifies at 80°. 
Ammonia gas may be con- 
densed, just like chlorine. Take 
ammonium silver chloride (Ag Cl 
2NH3), obtained by conducting 
ammonia over silver chloride, 
and fuse it in a tube with a knee- 
shaped bend (Fig. 70). The limb 
containing the compound is now 
heated in a water bath, while the 
other limb is cooled. The com- 
pound is decomposed into silver 
chloride and ammonia, which 
condenses in the cooled limb. 
Fig. 70 
Ammonia gas is exceed- 
ingly readily soluble in water, 
with the liberation of heat. 
One part of water at 0° and 
760 mm. pressure absorbs 
1050 volumes (== 0.817 parts by weight); at 15° 730 
volumes 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 into a vacant 
space; a piece of ice melts rapidly in the gas. The 
aqueous solution possesses all the properties of the free gas, 
and is called Liquor ammonii caustici. The specific gravity of 
the solution is so much less, the larger the quantity of 
ammonia present. The solution saturated at 14° contains 
about 30 °Jo NH;), and has a specific gravity of 0.897. Upon 
warming, all the gas escapes. 
The condeused liquid ammonia, upon evaporation, absorbs a great 
amount of heat and answers, therefore, for the artificial production of 118 
cold and ice in Carrfe’s apparatus. The simplest form of the latter is 
represented in Fig. 71. The iron cylinder A is filled about half with a 
concentrated aqueous ammonia solution, and is connected, by means of 
the tubes a and 6, with the conical vessel F, in the middle of which is 
the empty cylindrical space E. The entire internal space of A and F 
is hermetically shut off. A is 
heated upon a charcoal fire until 
the thermometer a, in it, in- 
dicates 130° C., while Fis 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 con- 
denses in B, of the receiver F, 
to a liquid. The cylinder A is 
removed from the fire, cooled 
with water, and the vessel B, 
constructed of thin sheet-metal 
and filled with water, placed in 
the cavity E, which is sur- 
rounded with a poor conductor, 
e. g., felt. The ammonia con- 
densed in B evaporates, and is 
reabsorbed by the water in A. 
By this evaporation a large quantity of heat, withdrawn from F 
and its surroundings, becomes latent; the water in D freezes. 
The method of Carrb for the artificial production of ice has 
acquired great application in the arts ; in recent times, however, it is 
being more and more replaced by the method of Windhausen. The 
latter depends upon the expansion of compressed air. 
INORGANIC CHEMISTRY. 
Fig. 71. 
Chemical Properties. By a glowing heat and continued 
action of the electric spark, ammonia is decomposed into nitro- 
gen and hydrogen. On conducting ammonia gas over heated 
sodium or potassium, the nitrogen combines with these metals 
and hydrogen escapes :— 
NH3 + 3K = NK3 + 3H. 
In the air ammonia will not burn; in oxygen, however, it 
burns with a yellow flame :— 
2NHS + 30 = N2 -f 3H20. 
Ammonium nitrite and nitrogen dioxide are formed simulta- 
neously. A mixture of ammonia and oxygen burns, when 
ignited, with explosion. 
To show the combustion of NH3 in O, proceed as follows : 
A glass tube, through which ammonia is conducted, is brought 
into a vessel with oxygen, bringing the opening of the latter AMMONIA 
119 
near a flame at the moment of the introduction of the glass 
tube. In contact with oxygen, the ammonia gas ignites and 
continues to burn in it. 
The following experiment (of Kraut) very conveniently 
shows the combustion of ammonia. Place a somewhat con- 
centrated ammonia solution in a beaker glass, heat over a 
lamp until an abundant disengagement of ammonia vapors 
occurs, and then, by means 
of a tube dipping into the 
liquid, run in oxygen gas. 
Upon approaching the mix- 
ture with a flame, it ignites 
writh a slight explosion. The 
ignition may be induced with- 
out a flame, by sinking a 
glowing platinum spiral into 
the mixture (Fig. 72) ; here 
explosion also occurs. Sim- 
ultaneously, the glass at 
first is filled with white 
vapors of ammonium nitrite 
(KH4N02) ; later, when oxy- 
gen predominates, red vapors 
of nitrogen dioxide (K02) 
and nitrous acid (HN02) ap- 
pear. 
If chlorine gas be conducted 
into the vessel with ammonia, 
it immediately ignites and continues to burn in the latter, at 
the same time forming fumes of ammonium chloride (NII4C1). 
The chlorine combines with the hydrogen of the ammonia to 
yield hydrochloric acid, which, with the excess of ammonia, 
produces 
Fio. 72. 
NH3 + 3 Cl = 3HC1 + N 
and 3NH3 + 3HC1 = 3NH4C1. 
In gaseous form, as well as in solution, ammonia possesses 
strongly basic properties ; it blues red litmus paper, neutralizes 
acids, forming salt-like compounds with them, which are very 
similar to the salts of the alkalies—sodium and potassium. 
The following illustrates the similarity :— 
NH3 + HC1 = NH4C1 KC1 
Ammonium 
chloride. 
Potassium 
chloride. 120 
INORGANIC CHEMISTRY 
2NH3 + H2S04 = (NH4)2S04 k2so4 
Am. sulphate. 
Potassium 
sulphate. 
nh3 + h2s = nh4sh ksh 
Am. 
sulphydrate. 
Pot. 
sulphydrate. 
In these ammonia derivatives NH4 plays the role of the metal 
potassium. Hence the group (NH4) has been designated 
Ammonium and its compounds, ammonium salts. By the action 
of strong bases ammonia is set free from the ammonium com- 
pounds :— 
The metallic character of the ammonium group is con- 
firmed by the existence of the ammonium amalgam and its 
entire deportment in compounds. Therefore, the ammonium 
derivatives will be considered with the metals. 
2NH4C1 + CaO = 2NH3 + CaCl2 + H20. 
Quantitative Composition of Ammonia. Atomic Weight of 
Nitrogen. 
The quantitative analysis-of ammonia shows that it consists 
of 1 part hydrogen and 4.66 parts nitrogen ; hence we con- 
clude that the atomic weight of N is a multiple of the last 
number (see p. 62). 
H = 1 
N = 4.66 
NH = 5.66 
2H = 2 
N = 9.33 
NH2 = 11.33 
3H = 3 
N = 14 
NHS = 17 
As the density of ammonia equals 8.5 (H = 1), then its 
molecular weight = 17. In 17 parts of ammonia are con- 
tained 3 parts, and, therefore, 3 atoms, of hydrogen. That 
the 14 parts nitrogen united with them correspond to one 
atom of N is a consequence, as never less than 14 parts of N 
are present in the molecular weight of any nitrogen deriva- 
tive. The density of nitrogen equals 14 and its molecular 
weight 28 ; therefore, the molecule of N consists of two atoms 
(N2). This is found from the volume ratios occurring in the 
formation of ammonia (see below). 
From the molecular formulas NH3 and N2 follows, further, 
that 1 vol. N and 3 vols. H form 2 vols. ammonia gas, or that 
2 vols. ISTH- decompose into 3 vols. H2 and 1 vol. N2, corres- 
ponding to the molecular equation :— 
N2* + 3H2 = 2NH3 
1 vol. 
3 vol 
2 vol. AMMONIA. 
121 
The following experiments prove these conclusions :— 
1. Decompose an aqueous ammonia solution, mixed with 
sulphuric acid to increase its power of conductivity, in a Hof- 
mann’s apparatus (Fig. 47), by the galvanic current. Hydro- 
gen will separate at the negative and nitrogen at the positive 
pole; the former will have three times the volume of the 
latter. 
2. The electric (induction) sparks are permitted to strike 
through dry ammonia gas enclosed in an eudiometer, or the 
apparatus represented in Fig. 58 (p. 89). In this way the 
ammonia is decomposed into nitrogen and hydrogen, whose 
volume is twice as large as that of the ammonia employed. 
That 3 vols. H are present in the mixture for each 1 vol. N 
is easily shown by the volumetric method, by burning the H 
with oxygen (p. 111). 
The volume ratios in the formation of ammonia confirm the 
conclusion drawn from the density of nitrogen (see above), 
that the molecule of the latter consists of two atoms. In two 
volumes of ammonia there are 2n molecules of NH3, therefore, 
2u atoms of N. The nitrogen contained in these 2 volumes of 
NH3 in a free condition occupies 1 volume, which contains n 
molecules and therefore 2n atoms of N. 
Hydroxylamine. NH30 = NH20H. This compound, very 
analogous to ammonia, was discovered (Lossen) in the reduc- 
tion of ethyl nitrate by zinc and hydrochloric acid. 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 eliminated by the tin which, in statu nascendi, 
reduces the nitric acid:— 
The hydroxylamine remains in solution in combination with 
nitric acid. Like ammonia, hydroxylamine unites directly 
with acids, to form salts :— 
HN03 + 3H2 = H3NO + 2H20. 
H3NO + HC1 = H3NO, HC1. 
It cannot be obtained in free condition, as it is very un- 
stable. On adding to the aqueous solution of the sulphate 
of hydroxylamine sufficient barium hydrate to remove all 
the sulphuric acid, an aqueous solution of the base is 
obtained, which, like the ammonia solution, possesses strongly 
9 122 
INORGANIC CHEMISTRY. 
basic properties, and blues red litmus paper. The solution is, 
however, very unstable, and readily decomposes into water, 
ammonia and nitrogen:— 
Owing to its great similarity to ammonia, it is supposed that 
hydroxvlamine represents ammonia in which, 1 H is replaced 
by the hydroxyl group OH ; therefore the name hydroxyl- 
amine :— 
3NHS0 = NH3 + 3H20 + N2. 
nh3o = nh2oh. 
COMPOUNDS OF NITROGEN WITH THE HALOGENS. 
Nitrogen Chloride. NC13. As we have seen, nitrogen is 
liberated when chlorine acts upon excess of ammonia; when, 
however, the chlorine is in excess, it acts upon the previously 
formed ammonium chloride, to produce nitrogen chloride :— 
For the preparation of a small quantity of nitrogen chlo- 
ride, 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. 
NH4C1 + 3C12 = NC1S + 4HC1. 
Nitrogen chloride is an oily, yellow liquid, with a disagree- 
able odor; its specific gravity equals 1.65. Of all chemical 
compounds this is the most dangerous, as it, by the slightest 
contact with many substances, frequently decomposes without 
any perceptible external cause, accompanied by an extremely 
violent explosion. 
The formation and explosibility of nitrogen chloride may be harm- 
lessly illustrated in the following manner: Decompose a saturated 
ammonium chloride solution with the electric current. Nitrogen 
chloride rising in small drops from the liquid will separate at the posi- 
tive pole. Upon covering the surface of the solution with a thin layer 
of turpentine oil, each drop, as it comes in contact with the latter, 
will explode. 
Nitrogen iodide. Upon saturating finely divided iodine 
powder with ammonium hydrate, or upon pouring an alcoholic 
solution of iodine into ammonium hydrate, a brownish black 
substance is obtained, which is extremely explosive. Its ex- 
plosibility may be shown without danger in the following 
manner: The precipitate is collected on a filter, washed 
with water, the filter opened 'out and torn into small pieces, PHOSPHORUS. 
123 
which are then allowed to dry; upon the slightest disturb- 
ance these pieces explode with a violent noise. According 
to the method of preparation the compound possesses the 
composition NHI2 or NIS; it is regarded as ammonia in 
which the hydrogen is partly or entirely replaced by iodine. 
P = 31. P4 = 124. Density = 62. 
2. PHOSPHORUS. 
This element does not occur free in nature, owing to 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 remain in their ash. In the 
animal kingdom calcium phosphate occurs in the bones. 
Brand and Kuukel, in Hamburg (1669), first obtained 
phosphorus by the ignition of evaporated urine. In 1769, 
Scheele, in Sweden, demonstrated its isolation from bones. 
Its name is derived from its power of giving light in the dark 
—<pajq<p6pos, i. e., light-bearer. 
To obtain phosphorus from bones the latter are burned, thereby 
destroying all organic admixtures and leaving bone ashes, which con- 
sist principally of tertiary calcium phosphate (P04)2Ca3 (see Phos- 
phoric acid). The ashes are now heated with f of their weight of 
sulphuric acid, the tri phosphate thus becoming primary calcium 
phosphate, and gypsum (cal. sulphate) resulting as a bye-product:— 
Ca3(P04)2 + 2H2S04 = CaH4(P04)2 + 2CaS04 
The gypsum, which is very difficultly soluble 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. This converts the primary phosphate, with loss 
of water, into calcium metaphosphate :— 
Tertiary 
calcium phosphate. 
Primary 
calc, phosphate. 
Calcium 
sulphate. 
CaH4 (P04)2 = Ca (P03)2 + 2H20 
Calcium 
metaphosphate. 
The ignited residue is then raised to a white heat, in retorts of in- 
fusible clay. The carbon partly reduces the metaphosphate to phos- 
phorus, by forming carbon oxide gas with oxygen, and half of the 
phosphorus contained in the metaphosphate remains as calcium pyro- 
phosphate :— 
2Ca (P03)2 + 5C = 2P + 5C0 + Ca2P207 
Calcinra 
pyrophosphate. 124 
INORGANIC CHEMISTRY. 
The liberated phosphorus escapes in vapor, and is collected under 
water in receivers of peculiar construction, and there condensed. To 
remove mechanically admixed impurities the phosphorus is again dis- 
tilled from retorts and fused under water; it is then moulded into 
sticks. 
The crystalline or yellow phosphorus obtained by distillation 
is a waxy, transparent, slightly yellow-colored substance, with 
specific gravity of 1.83 at 10° C. At ordinary temperatures 
it is soft and tough ; at 0° it becomes brittle. Fuses under 
water at 44° and boils at 290° C. By the action of sunlight 
it becomes yellow and is coated with a non-transparent, 
reddish-white layer. In water, phosphorus is insoluble; 
slightly insoluble in alcohol and ether; very readily soluble 
in carbon disulphide. From the latter solution it crystallizes 
in forms of the isometric (rhombic dodecahedra) system. 
Exposed to the air, it readily oxidizes to phosphorous acid 
(H:,POa) ; the white vapors which arise contain ammonium 
nitrite (NH+N02), ozone and hydrogen peroxide. Its odor 
resembles that of ozone. In the air it phosphoresces at night. 
Also shines in other gases, but only in such as contain oxy- 
gen. It appears the phosphorescence is influenced by the 
formation and combustion of the self-inflammable phosphine, 
as all substances which destroy the latter also prevent and put 
an end to the former. It is noteworthy that in pure oxy- 
gen the oxidation of phosphorus occurs first at 27°. If the 
oxygen be diluted by removal over an air pump or the addition 
of neutral gases, so that its quantity is not more than 40 %, 
the absorption will be very energetic at 20°, but cease entirely 
at 70°. 
Another modification—the red or amorphous phosphorus— 
possesses properties entirely different from the ordinary vari- 
ety. It is a reddish-brown amorphous powder, of specific 
gravity 2.14; insoluble in carbon disulphide, non-phosphores- 
cent, does not alter in the air, and is, indeed, very stable. 
While ordinary phosphorus is very poisonous, this variety is 
perfectly harmless. Infusible at a red heat, not even under 
strong pressure, and vaporizes very slowly (above 260°), and 
only partially, the vapors passing over into ordinary phos- 
phorus. 
To prepare the red, the first variety is heated for some 
minutes to 300°, in closed, air-tight iron vessels ; a slo\vT con- 
version occurs already at 250°. The resulting mass is then PHOSPHORUS. 
125 
treated with carbon disulphide or sodium hydrate, to withdraw, 
unaltered, ordinary phosphorus. If some iodine be added to 
the ordinary phosphorus, the change will occur below 200°. 
A third modification—metallic phosphorus—is formed, if the amor- 
phous variety be heated in a glass tube, free of air, to 530°. Micros- 
copic needles then sublime into the upper, less heated portion of the 
tube. More easily obtained if phosphorus be heated with lead, in a 
closed tube, to a red heat. The molten metal dissolves the phosphorus, 
which, on cooling, separates in black, metallic, shining crystals. Me- 
tallic phosphorus possesses the specific gravity 2.34, vaporizes with 
difficulty, and is less active than the amorphous variety. 
Two green lines characterize the spectrum of phosphorus. 
On conducting hydrogen gas over a small piece of it, heated 
in a glass tube, the escaping gas burns with a bright green 
flame, in consequence of the phosphorus it contains. When 
ordinary phosphorus is distilled with water, some passes over 
with the steam and, in the dark, phosphoresces. This proce- 
dure serves for the detection of phosphorus in poisoning with 
this substance. 
The density of P equals 62 (H = 1), or 4.42 (air = 
1); the molecular weight is, therefore, 124. As the atomic 
weight of P is 31, it follows that the molecule in form of vapor 
consists of 4 atoms : P4 = 124 (31 X 4). We saw that the 
sulphur molecule at 500° consists of 6 atoms (S6), at 900° of 
2 atoms (S2). Such a dissociation does not, however, occur 
with phosphorus ) even at 1040° its vapor density remains 
unaltered. 
When burned in oxygen or in air, phosphorus forms 
the pentoxide (P205). The ordinary variety inflames at 
40°, and also by gentle friction; the amorphous not below 
260°. Even under water the first burns with a bright 
flame. To this end heat some pieces of P in a flask with 
water, until they fuse, and conduct a current of oxygen 
through the water. Phosphorus combines with Cl, Br and 
I very energetically at ordinary temperatures ; by throwing a 
small piece of it 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 throws out some metals from solutions of 
their salts. From a silver nitrate solution, it precipitates 
silver and phosphorus-argentide (PAg3); this solution, 
therefore, answers as a counter-irritant in phosphorus burns. 126 
INORGANIC CHEMISTRY. 
COMPOUNDS OF PHOSPHORUS WITH HYDROGEN 
ph3, p2h4, p4h2. 
On warming yellow phosphorus with potassium or sodium 
hydrate, spontaneously inflammable hydrogen phosphide is 
produced. 
The liberated gas mixed with air in a closed vessel explodes 
violently, hence, to make it, proceed as follows: Fill a glass flask 
almost full of aqueous KOH, add a few pieces of P, and heat over a 
lamp (Fig. 73). 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 
Fig. 73. 
the delivery tube, the other end of which dips under warm water, to 
prevent any obstruction arising in it from phosphorus that has been 
carried over and solidified by cooling. Each bubble rising from the 
liquid inflames in the air, and forms white cloud-rings which ascend. 
The gas thus produced consists of gaseous phosphine (PH3) 
and hydrogen, with which is mixed a small quantity of a 
liquid substance (P2H4), whose presence imparts the sponta- 
neous inflammability to the gas. On conducting the latter 
through a cooled tube, the liquid phosphine (P2H4) is con- 127 
densed to a liquid, and the escaping gas no longer inflames 
spontaneously. Similarly the liquid compound may be isolated 
if the gas, is conducted through alcohol or ether, by which 
the compound P2II4 is absorbed. 
Liquid Phosphine, P2H*, separated from the gas by cooling, 
presents a colorless, strongly refracting liquid, insoluble in 
water, and boiling at 30°. Inflames spontaneously in the air, 
and burns with great brilliancy to phosphorus pentoxide and 
water. Its presence in combustible gases, such as hydrogen, 
marsh gas, and PH3, gives to them their spontaneous inflam- 
mability. In contact with some compounds, like carbon and 
sulphur, and by the action of sunlight, it decomposes into 
gaseous and solid phosphine :— 
PHOSPHORUS 
Solid Phosphine, P4H2, is a yellow powder inflamed at 
160° or by a blow. Formed in the decomposition of calcium 
phosphide by hydrochloric acid. 
5P2FI4 =6PH3 + P4H2. 
Gaseous Phosphine, PH:!, may be formed, in addition to 
the manner previously described, by the action of water or 
hydrochloric acid upon calcium phosphide :— 
Ca3P2 + 6HC1 = 3CaCI2 + 2PH3 
Further, by the glowing of phosphorous and hypophos- 
phorous acid:— 
4H3P03 = PH3 -f- 3H3P04 
Phosphorous 
acid. 
Phosphoric 
acid. 
A colorless gas with a disagreeable, garlic-like odor, and 
somewhat soluble in alcohol. Its density is 17 (H = 1), or 
1.185 (air = 1). Pure—freed of P2H4—it ignites at 100° C. 
Oxidizing agents convert it again into the spontaneously in- 
flammable variety, owing to the production of P2H4. It is 
extremely poisonous. By heat and the action of the electric 
current PH3 decomposes into phosphorus and hydrogen. 
Mixed with chlorine it explodes violently, with production of 
phosphorus tri-chloride and hydrogen chloride :— 
PH3 + 3 Cl 2 = PC13 +3HC1. 
Like ammonia, gaseous hydrogen phosphine possesses faint 
alkaline properties, and combines with hydrogen iodide and 
bromide to yield compounds similar to ammonium chloride:— 
PH3 + HI = PH J. 128 
INORGANIC CHEMISTRY. 
With ITC1 it combines at 30°-35°, or at the ordinary tem- 
perature under a pressure of 20 atmospheres. The group 
PIi4, figuring in the role of a metal in these compounds, is 
analogous to ammonium (p. 120), and termed Phosplionium. 
Phosphonium Iodide, PHJ. It is best prepared by the 
decomposition of phosphorus di-iodide (PI2), by a slight quan- 
tity of water. It sublimes in colorless, shining, cube-like rhom- 
bohedra; fumes in the air, and, with water, decomposes into 
PH3 and HI. Decomposed by potassium hydrate it yields 
pure hydrogen phosphide, which is not spontaneously inflam- 
mable :— 
PHJ + KOH = KI + PH3 + H20. 
Molecular Formula of Phosphine. Atomic Weight of Phos- 
phorus. 
Analysis of gaseous phosphine shows that it consists of 1 part 
hydrogen and 10.33 parts phosphorus. Were its molecular formula 
PH the atomic weight of P would be 10.33. The great analogy of 
phosphine, with NH3, as also all the P compounds with those of N, 
argues, however, for the formula PH3. The atomic weight of P, 
therefore, is 31 (= 3 X 10.33), and the molecular weight of the gase- 
ous phosphine is 34 :— 
II3 = 3 
P = 31 
PH 3 = 34 
This view is confirmed by the density. Corresponding to 
the formula, it must be V = 17, which follows from direct 
experiment. Further, from the formula PHS, it follows that 
in 2 volumes of the gas 3 volumes of hydrogen are present:— 
2PH3 contain 3H2 
Or in 1 volume are contained II volumes hydrogen. On 
decomposing the gas in an eudiometer by the passage of 
electric sparks, it will be found that the volume increases 1 £ 
times ; the gas consists, then, of pure hydrogen, while phos- 
phorus separates in a solid condition. As the phosphorus 
molecule in gaseous condition is composed of 4 atoms, the 
phosphorus (31 parts) separated from 2 volumes PHS, in 
vapor form, will fill I volume; hence in 2 volumes of PH3 3 
volumes of H and 2 volume phosphorus vapor are present:— 
Or, written molecularly:— 
2 vols. 
3 yols. 
P4 -}- 6H2 == 4PHS 
1 vol. 
6 vols. 
4 vols. PHOSPHORUS CHLORIDE 
129 
COMPOUNDS OF PHOSPHORUS WITH THE 
HALOGENS. 
Phosphorus combines directly with the halogens, to yield 
compounds of the forms PX3 and PX5, in which X indicates 
an halogen atom. 
Phosphorus Trichloride — Phosphorous Chloride — PC13. 
Conduct dry chlorine gas over phosphorus gently heated in 
the retort D (Fig. 74). The phosphorus ignites in the stream 
of gas and distills over as trichloride, which is collected 
in the receiver E and condensed. The product is purified by 
a second distillation. It is a colorless liquid, boiling at 74° 
Fig. 74. 
C., and has a sharp, peculiar odor. Its specific gravity equals 
1.616 at 0°. It fumes strongly in the air, decomposing with 
moisture to phosphorous and hydrochloric acids :— 
PC13 + 3H20 = H3P03 + 3HC1. 
The vapor density of the trichloride equals 68.7 (H = 1), 
corresponding to the molecular formula PC13 = 137.5. 
Phosphorus Pentachloride—Phosphoric Chloride—PC15. 
Produced by the action of an excess of chlorine upon the 
liquid trichloride. It is a solid, crystalline, yellowish-white 
compound. Fumes strongly in the air and sublimes by heat- 
ing, without fusion, but w’ith partial decomposition, into tri- 
chloride and chlorine. 130 
At lower temperatures (in an atmosphere of chlorine) the vapor 
density of the pentachloride has been found to be 104.2, corresponding 
to the molecular formula PC15 (-af-~ = 104 2). At increased tempera- 
tures the vapor density steadily diminishes, a gradual decomposition 
occurs—dissociation (p. 85) of the molecules PC15 into the molecules 
PCI3 and Cl2. At 236° the dissociation is complete, and then equals 
the vapor density 52.2 ; i. e.,the vapor then fills doub'e as large a 
volume as at a lower temperature. The breaking up of PC15 into PCI3 
and Cl2 explains this :— 
INORGANIC CHEMISTRY. 
That, in fact, such a decomposition of the penta- into trichloride and 
chlorine does occur, is, among other things, proven by the originally 
colorless vapor, gradually assuming the yellow color of chlorine with 
increasing temperature. The decomposition products—PC13 and Cl2 
—may be separated from each other by diffusion (p 114). 
1 vol. PC15 = 1 vol. PC13 -f- 1 vol. Cl2. 
With water PC15 acts very energetically. With little water 
it forms the oxychloride and hydrochloric acid:— 
Phosphorus Oxychloride—POCl3, is a colorless liquid, fum- 
ing strongly in the air, wTith a specific gravity, at 12°, of 1.7. 
Boils without decomposition at 110°. Its vapor density 
equals 72.7, corresponding to the molecular formula POCl3 = 
145.5. Water decomposes it into phosphoric and hydro- 
chloric acids:— 
PC15 + H20 = PC130 + 2HC1. 
Phosphorus oxychloride is most practically obtained by dis- 
tillation of PCI5, with excess of phosphorus pentoxide:— 
P0C1* + 3H20 = H3P04 + 3HC1. 
3PC15 + P205 = 5P0C13. 
Or chlorine gas is conducted into a mixture of 3PC13 and 
P205 
3PC1S + P205 + 3C12 = 5P0C13. 
Particularly interesting is the production of it by conduct- 
ing ozonized air through phosphorus chloride (Rernsen). In 
the same manner PC13 unites with S when heated to 130°, to 
form phosphorus sulpho-chloride—PSC13— an oily liquid, sink- 
ing in water, which with the latter decomposes into SH2, C1H 
and P03H. 
The bromine and iodine phosphorus compounds are per- 
fectly analogous to the chlorine derivatives. They are obtained 
by uniting the constituents in the proportions by weight ex- 
pressed by their formulas. As the union is exceedingly ener- 131 
getic, it is best to proceed as follows: Dissolve the phosphorus 
in carbon disulphide, add gradually the calculated amount of 
Br or I, and then distill off the volatile carbon disulphide. 
ARSENIC. 
Phosphorus Tribromide PBr3, is a colorless liquid, boiling at 
175°, and having a specific gravity of 2.7. The pentabromide, 
formed by gradual addition of 2Br to PBr3 is a yellow, crystal- 
line substance, which by heat fuses, and breaks up into PBr3 
and Br2. Water decomposes both compounds, as it does the cor- 
responding chlorides. Phosphorus oxy-bromide (POBrs), is a 
colorless crystalline mass, fusing at 45°, and boiling at 195°. 
Phosphorus Chlor-bromide, PC13 Br2 is produced by the 
union of PCI 3 with Br3 in the cold. A yellowish red mass, 
which decomposes at 35° C., into PC13 and Br2. 
Phosphorus Tri-Iodide—PI3, forms red crystals, fusing 
at 55° and distilling, with partial decomposition, at a higher 
temperature. The so-called phosphorus iodide, PI2, or P2I4 
(corresponding to P2H4), crystallizes in beautiful orange red 
crystals or prisms, and fuses at 110°. A little water decom- 
poses it into phosphorous acid, PH3 -f- HI, and the last two 
bodies form phosphonium iodide, PH4I (p. 128). 
Of interest is the recently discovered Phosphorus Pentaflu- 
oride—PF15—which results upon heating PC13 or PC15 with 
arsenic trifluoride, AsF13. 
3PC15 + 5AsF13 = 3PF15 + 6AsCls. 
It is a colorless gas, fuming in moist air, which, with water, 
breaks up into phosphoric acid and hydrogen fluoride. Its 
density is 63, corresponding to the molecular formula PF15 = 
126. 
Remarkable is the fact that, while phosphorus penta- 
iodide could not be obtained, the stability of the compounds, 
PBr5, PC15, PF15, gradually increases with diminution of the 
atomic weight of the combined halogens. 
3. ARSENIC. 
Arsenic is a perfect analogue of phosphorus, but possesses a 
somewhat metallic character. Free, it is similar to metals. 
Arsenic is found free in nature. More frequently in com- 
bination with sulphur (realgar, orpiment), with oxygen 
(arsenolite, As203), and with metals (mispickel, FeSAs, 
As = 75. As4 = 300. Vapor density, 150. 132 
INORGANIC CHEMISTRY. 
cobaltite, CoAsS). To prepare it, heat mispickel with some 
iron ; when free, arsenic will sublime. Or, in the customary 
way of isolating metals from their oxides, heat the trioxide 
(arsenolite) with charcoal:— 
As203 + 3C = 2As + 3C0 
Arsenic appears in two modifications. The amorphous 
arsenic is an almost lustreless, black mass, with a specific 
gravity 4.71. It is brittle and can be easily pulverized. 
Crystallized arsenic, obtained by continuous heating to 210- 
220°, forms steel gray hexagonal rhombohedra, with strongly 
metallic lustre ; its specific gravity equals 5.7. 
Arsenic volatilizes at 180° without previously fusing; 
under strong pressure and in a sealed tube it is fusible. Its 
vapor possesses a lemon-yellow color. The vapor density is 
150 (H = 1), the molecular weight, therefore, 300. As its 
atomic weight equals 75, it follows that its gas molecule, like 
that of phosphorus, consists of four atoms (As4 = 300 = 4 
X 75). _ 
Arsenic does not change in dry air. Heated, it burns with 
a blue colored flame and disseminates the garlic-like odor of 
arsenic tri-oxide (As203). It combines directly with most 
elements. Powdered arsenic projected into chlorine gas in- 
flames. With the metals it yields arsenides. 
It is remarkable that arsenic, which belongs to the nitrogen group 
and generally forms entirely differently constituted compounds from 
sulphur, is analogous to the latter in its metallic combinations. Thus 
the sulphides and arsenides have similar formulas, and in them sulphur 
and arsenic can mutually replace each other in atomic ratios, e. g. 
FeS2, FeAs2 and Fe(SAs). 
COMPOUNDS OP ARSENIC WITH HYDROGEN. 
Arsine, AsH3 = 78. Like nitrogen and phosphorus, arse- 
nic furnishes a gaseous compound containing 3 atoms of 
hydrogen. It is obtained pure by the action of dilute sul- 
phuric acid or hydrochloric acid upon an alloy of zinc and 
arsenic:— 
A.S2Z113 ~f" 6HC1 = 3ZnCl2 2AsH3. 
It also results in tlie action of nascent H (zinc and sulphuric 
acid), upon many arsenic compounds, like the tri-oxide :— 
As203 + 6H2 = 2AsH3 + 3H20. 133 
Arsine is a colorless gas, of strong, garlicky odor, and ex- 
tremely poisonous action, and at—40° is condensed to a liquid. 
Its density equals 39 (H = 1), or 2.69 (air = 1). Ignited it 
burns with a blueish white flame, with the evolution of 
white fumes of arsenic tri-oxide :— 
ARSENIC WITH HYDROGEN. 
2AsH3 “f* 302 ==: As203 ~f~ 3H20. 
Bv gentle glowing or by the electric spark it is decom- 
posed into arsenic and hydrogen. On conducting the gas 
through a heated tube the arsenic deposits itself behind the 
heated part as a metallic coating (arsenic mirror). On hold- 
ing a cold object, e. g., a piece of porcelain, in the flame of 
the gas the arsenic forms a black deposit (arsenic spots). In 
its chemical behavior arsine is very similar to PH,, pos- 
sesses, however, scarcely any more basic properties, and does 
not furnish any derivatives with the halogens. 
According to analysis, arsine consists of 1 part by weight of hydro- 
gen and 25 parts arsenic. If, owing to its analogy to PH3, we ascribe 
to it the formula AsH3, then the atomic weight of arsenic would be 75 
(3 X 25), and the molecular weight of AsH3 = 78. Hence the 
density must be -7/- = 39, which is confirmed by experiment. From 
the formula it follows that 3 volumes of hydrogen are contained in 2 
volumes AsH3:— 
2AsH3 contain 3H2 
We may become convinced of this by decomposing the gas by elec- 
tricity in an eudiometer (see p. 128). 
2 vols. 
3 vols. 
Marsh’s Method for the Detection of Arsenic. Resting upon 
the above described method of forming AsH3 and its charac- 
teristic properties, arsenic may be readily detected in the 
following manner—a task which, owing to the poisonous nature 
of this element, is important: Hydrogen is generated in a 
flask a (Fig. 75), by action of dilate sulphuric acid upon zinc, 
and through the funnel tube is then introduced some of the 
solution to be examined for arsenic. The liberated gas, a 
mixture of hydrogen and arsine, is dried in the calcium chlo- 
ride tube c and escapes through the difficultly fusible glass tube 
d, which is at various places drawn out into narrow parts. 
Upon igniting the escaping hydrogen (after all the air has been 
previously expelled from the vessel, as otherwise oxy-hydrogen 
gas will be present) it will burn, providing AsHs is present, with 
a bluish-white flame and dissemination of white vapor. The 
dark arsenic spots are obtained by holding a cold porcelain 134 
INORGANIC CHEMISTRY. 
dish in the flame. If the tube d be heated (as shown in 
Fig 75), an arsenic mirror is formed upon the adjoining con- 
traction of the tube. The slightest traces of arsenic may be 
detected by this method. 
Fig. 75. 
Besides the ordinary arsine, AsH3, the existence of As2H4 and 
As4H2, corresponding to the liquid and solid phosphines (P2H4 and 
P4H2) might be expected. The first is not known, but derivatives of it 
are, which, in place of hydrogen, contain hydrocarbon groups, such as 
cacodyl As2(Cff3)4 = (CH3)2As-As(CH3)2 Nitrogen affords similar 
compounds—(CH3)2N-NH2 and (CH3)NH-NH2, which are derived 
from diamine (N2H4 = H2N-NH2) not known in a free condition. 
The solid arsine As4H2 is obtained by the action of nascent hydro- 
gen upon arsenic compounds in presence of nitric acid. It forms a 
red-brown powder, which decomposes when heated. 
COMPOUNDS OF AKSENIC WITH THE HALOGENS. 
These are perfectly analogous to the corresponding phos- 
phorus compounds, and are the result of the direct union of 
their constituents; however, the compounds with the formula 
AsX5 (see p. 129) are not known. The metallic character of 
arsenic appears in the fact that arsenic chloride, like other ANTIMONY. 
135 
metallic chlorides, may be obtained by the action of hydro- 
chloric acid upon the oxide:— 
Upon boiling a solution of As203 with strong HC1, arsenic 
chloride is set free. 
As203 + 6HC1 = 2As013 + 3H20. 
Arsenic Trichloride—AsC13. A colorless, oily liquid, 
fuming in the air and having a specific gravity of 2.2. It 
solidifies at — 30° and boils at 134°. The vapor density 
equals 90.7 (H = 1), corresponding to the molecular 
formula AsC13 = 181.5. In a small quantity of H20 the 
chloride dissolves without change; much water converts it 
into oxide and hydrochloric acid :— 
2AsC13 + CH20 = As203 + 6HC1. 
Arsenic Tribromide, AsBr3, is a white crystalline mass, 
fusing at 20°, and boiling at 220° C. The Tri-iodide, Asia, 
forms red crystals; the Trifluoride, AsFl3, is a liquid, turn- 
ing strongly in the air. It results in the distillation of AsC13 
or As203 with calcium fluoride and sulphuric acid. 
4. AMTMOXY. 
Sb = 122. 
The metallic character which appeared with arsenic, ex- 
hibits itself more distinctly with antimony, which at the same 
time retains its complete analogy with the metalloidal ele- 
ments, arsenic and phosphorus. As far as concerns its physical 
properties, it is already entirely metallic. 
Antimony (Stibium) occurs in nature, principally in union 
with sulphur, as stibnite, Sb2S3, and with sulphur and metals, 
almost always accompanied by arsenic, in many ores. To pre- 
pare it, stibnite is roasted in a furnace, i. e., heated with air 
access, whereby sulphur burns and antimony trioxide 
remains:— 
The residual oxide is ignited together with carbon, which 
causes a reduction to metal (general procedure for the sepa- 
ration of metals). Antimony may also be obtained by heating 
its sulphide with iron, which combines with the sulphur :— 
Sb2S3 + 90 = Sb203 + 3S02. 
The resulting commercial crude antimony is further puri- 
fied in the laboratory by fusing it with nitre, whereby admixed 
Sb2S3 + 3Fe = 2Sb + 3FeS. 136 
INORGANIC CHEMISTRY. 
arsenic, sulphur and lead are removed. Chemically pure 
antimony is obtained by reducing the pure oxide. 
It is a silver-white, strongly shining metal, of leafy crystal- 
line structure ; specific gravity 6.715. Like arsenic it crystal- 
lizes in rhombohedra, is very brittle, and may be easily 
broken. Fuses at 430°, and distills at a white heat. It is not 
altered in the air at ordinary temperatures ; upon heating it 
burns with a blue flame, with production of white vapors of 
antimonic oxide, Sb203. Like phosphorus, it combines directly 
with the halogens ; powdered antimony inflames in chlorine 
gas. It is insoluble in hydrochloric acid ; by nitric acid it 
is oxidized to antimonic oxide. 
Hydrogen Antimonide—Stibine—(SbH3), produced like 
arsine, and very similar to the latter ; thus far only obtained 
mixed with hydrogen. A colorless gas of peculiar odor, 
which ignited, burns with a greenish-white flame, and separa- 
tion of white vapors of antimonic oxide. A red heat decom- 
poses it into antimony and hydrogen. In Marsh’s apparatus 
(Fig. 75, p. 134) it gives, like AsH3, an antimony mirror and 
spots. The mirror is distinguished from that of arsenic by its 
black color, lack of lustre, its insolubility in a solution of 
sodium hypochlorite (NaCIO); further, by its slight volatility 
in a current of hydrogen. 
When a solution of SbCl3 is decomposed by the galvanic current, 
metallic antimony, strongly impregnated with hydrogen, separates. 
This is more readily obtained if antimony be made the negative elec- 
trode in the galvanic decomposition of water. When heated (or on 
breaking the rod) such antimony throws off the hydrogen with explo- 
sion (compare palladium hydride). 
COMPOUNDS OF ANTIMONY WITH THE HALOGENS. 
Antimonous Chloride—Trichloride—SbCl3, results from the 
action of chlorine upon the metal or its sulphide; better by 
the solution of the oxide or sulphide in strong hydrochloric 
acid:— 
This solution is evaporated to dryness and the residue dis- 
tilled. 
Sb2S3 + 6HC1 = 2SbCl3 + 3H2S. 
A colorless, crystalline, soft mass (Butyrum Antimonii), 
fusing at 73° and boiling at, 223°. Its vapor density equals 
114.2° (H = 1), corresponding to the molecular formula, ANTIMONY. 
137 
SbCl3 = 228.5. In the air it attracts water and deliquesces. 
Dissolved without alteration by water acidified with hydro- 
chloric acid. Much water decomposes it; the solution be- 
comes turbid and a white powder—algar red—separates:— 
The composition of this powder varies with the conditions 
under which it is formed, and generally corresponds to the 
formula 2 (SbOCl), Sb203. Pure Antimony Oxychloride, 
SbOCl, obtained by heating SbCl3 with alcohol, presents 
colorless crystals and is further decomposed by water. 
While the metallic chlorides are not decomposed by water 
at ordinary temperatures, yet, in the ready decomposition of 
the halogen derivatives of antimony, is exhibited the partial 
metalloidal character of this element. 
SbCl3 + H20 = SbOCl + 2HC1. 
Antimonic Chloride—Pentachloride—SbCl5, results from 
the action of excess of chlorine upon antimony or the tri- 
chloride. Forms a yellowish liquid which fumes in the air, 
becomes crystalline when cold and fuses at 6°, and crystal- 
lizes in three different forms. Heat partially decomposes it, 
like PC15, into SbCl3 and Cl2:— 
Water changes it to pyroantimonic acid (H4Sb207), and 
hydrochloric acid. 
Antimony Tribromide—SbBr3—is a white, crystalline 
substance, fusing at 94° and distilling at 270°. The Tri- 
iodide, Sbl3, is a red compound, crystallizing in three distinct 
forms. 
SbCl5 = SbCl3 + Cla 
1 vol. 
1 vol. 
1 vol. 
To the group, nitrogen, phosphorus, arsenic and antimony, 
belongs yet bismuth, which gives rise to similarly constituted 
compounds, e. g., BiCl3, Bil3, BiOCl. The metallic character, 
however, of bismuth considerably exceeds its metalloidal. 
Thus it does not unite with hydrogen, and the oxide (Bi203), 
similarly constituted to the acid-forming As203, possesses only 
basic characters. Therefore, we will consider bismuth and 
its derivatives with the metals. 
Vanadium. 
Vd = 51 
Niobium, 
Nb = 95 
Tantalum, 
Ta = 182 
In close relation to the phosphorus group stand three rare 
elements, vanadium, niobium and tantalum. They yield de- 138 
INORGANIC CHEMISTRY. 
rivatives very much like those of P; possess, however, a 
more metallic character, and do not combine with hydrogen. 
They exhibit many characteristics similar to those of chro- 
mium, iron and tungsten, with which, in their naturally 
occurring compounds, they are frequently associated (compare 
the Periodic System of the elements). 
Vanadium occurs in nature principally in the form of salts 
of vanadic acid (vanadium lead ore) and in some iron ores. 
It is very difficult to prepare it in a pure state, and is a white- 
gray, metallic, lustrous powder, of specific gravity 5.5. It is 
difficultly fusible, and does not change in the air. Heated, it 
burns to Vd205. 
Vanadium Trichloride—VdCl3—forms red plates, which 
readily deliquesce in the air ; it is not volatile. 
Vanadium Oxychloride—VdOCl3—results from heating a 
mixture of Vd203 and C in chlorine gas, forming a lemon- 
yellow liquid, of specific gravity 1.84, and boiling at 120°. It 
fumes strongly in the air and with water (analogous to phos- 
phorus oxychloride), it decomposes into vanadic acid and 
hydrochloric acid. Its vapor density equals 85 (VdOCl3 = 
173). 
Vanadium Oxide—Vd203—is a black powder obtained by 
heating VdaOs in hydrogen. It combines with O, to form 
Vd205. 
Vanadium Pentoxide—Vd205—or vanadic anhydride, is a 
mass obtained by fusing the naturally occurring vanadates 
with nitre, etc. It is soluble in the alkalies, yielding with the 
metals salts of Vanadic H3Vd04, and metavanadic acids 
HVd03 (compare Oxygen compounds of P). All these com- 
pounds are analogously constituted to those of P. In addition 
to these, vanadium forms other compounds, constituted analo- 
gously to those of sulphur and chromium. To these belong 
VdCl2 (dichloride), the tetrachloride VdCl4, vanadious oxide 
VdO, vanadium dioxide Vd02 and VdOCl2. The tetrachlo- 
ride VdCl4 is a red-brown liquid, boiling at 154° ; its vapor 
density equals 96 (\rdCl4 = 192). Niobium and tantalum 
are not known in a free state. The chlorides NbCl5 and 
TaCl5 are volatile and decomposed by water. The oxides 
Nb205 and Ta205 form, with bases, salts of niobic (H3Nb04) 
and tantalic (H3Ta04) acids. ANTIMONY. 
139 
Tabulation of the Elements of the Nitrogen Group 
The elements belonging here—nitrogen, phosphorus, arsenic, 
antimony and bismuth—present similar graded differences in 
their physical and chemical properties, like the elements of 
the chlorine and oxygen group, and this gradation is intimately 
related to the atomic weights. With the increase of the latter 
the substance condenses, the fusibility and volatility decrease, 
and the metallic character becomes more prominent:— 
N 
P 
Ab 
Sb 
Atomic weight, 
14 
31 
75 
122. 
Specific gravity, 
1.8-2.1 
4 . V . V 
67. 
Fusion point, 
44° 
heat 
Vapor density, 
0.972 
4.32 
10.3 
... 
With exception of metallic bismuth, the first four elements 
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 with HBr and HI to form salt-like compounds. 
AsH3 and SbH3 no longer show basic properties. There are, 
however, compounds of arsenic and antimony, as well as of 
the two preceding elements, with hydrocarbons (e. g., CH3 and 
C2H5), which are analogously constituted to the hydrides, and 
very similar to them in character. These compounds (As 
(CH3)3 and Sb (CH3)3) which will be described in organic 
chemistry, possess basic properties and yield salts correspond- 
ing to those of ammonium. 
The oxygen derivatives of these elements exhibit a similar 
gradation. With increase of atomic weight corresponding to 
the addition of metallic character, the oxides, which in the 
lower series form strong acids, acquire a more basic nature. 
The thermal relations of the elements of the nitrogen group are as 
yet but little investigated. In the production of one molecular weight of 
ammonia gas from hydrogen and nitrogen (H3N—gas) 11890 calories are 
set free ; the union, however, occurs almost only in an indirect way. It is 
very probable that the heat disengagement of PH3, AsH3 and SbHs 
is gradually a less one, for which the increasing decomposition of the 
same would argue. On the opposite hand, in the formation of the 
chlorides with increase of atomic weights and metallic character the 
heat disengagement successively increases, corresponding to the 
symbols:— 
(N, Cl3) = — 38100 (P, Cl3) =75800 (As, Cl3) = 74600 
Sb, Cl3 = 86300. 140 
INORGANIC CHEMISTRY. 
The great instability of the halogen derivatives of nitrogen finds 
explanation in the remarkable heat absorption taking place in their 
formation. 
CARBON GROUP 
Here belong the two non-metals, carbon and silicon, and 
the metal Tin. These unite, each, with four atoms of hydro- 
gen and four of the halogens. 
1. CARBON. 
Free carbon occurs in nature as the diamond and graphite. 
It constitutes the most important ingredient of all the so- 
called organic substances originating from the animal and 
plant kingdoms, and is especially contained in the fossilized 
products arising from the slow decomposition of plant matter 
—in turf, in brown coal, stone coal and in anthracite In 
combination with hydrogen it forms the so-called mineral oils 
—petroleum and asphaltum. It occurs, further, as carbon 
dioxide (CCh) in the air and in the form of carbonates (mar- 
ble, calcite, dolomite), comprises many minerals and entire 
rock formations. 
In free condition it is found in different allotropic modifica- 
tions, which may be referred to the three principal varieties: 
diamond, graphite and amorphous carbon. In all these forms 
it is a solid, even at the highest temperatures, non-fusible 
and non-volatile substance, which is only explainable by the 
supposition that its free molecules are composed of a large 
number of atoms of elementary carbon combined with each 
other. (See p. 97.) As far as concerns their chemical deport- 
ment, all the modifications of carbon are very stable and 
little capable of chemical reaction. When burned, all yield 
carbon dioxide. 
C — 12. 
1. The diamond occurs in alluvial soils in certain districts (in India, 
Brazil and South Africa); less frequently in micaceous schist. It has 
great lustre, strong power of refraction, and the greatest hardness of 
all substances. It crystallizes in forms of the regular system, mostly 
rhombohedra, seldom in octahedra. Ordinarily it is perfectly color- 
less and transparent, sometimes, too, colored by impurities. Its 
specific gravity equals 8.6. Only in the strongest heat, between the 
poles of a powerful galvanic battery, does it soften any, and is con- 
verted into a graphitic mass. Heated in oxygen gas, it burns to carbon CARBON 
141 
dioxide. By the action of a mixture of nitric acid and potassium 
chlorate it is scarcely at all attacked. 
2. The fact that graphite heated with nitric acid and potassium 
chlorate yields graphitic acid—and oxidized by potassium perman- 
ganate to mellitic acid—is characteristic. Native graphite is found in 
the oldest rock formations, and of especially good quality at Altai, in 
Siberia. It occurs, too, at many places in the United States, in consider- 
able quantities. Occasionally it appears crystallized in six-sided forms, 
mostly, however, as an amorphous, grayish-black, glistening, soft 
mass, used in the making of lead pencils. The specific gravity is 2.25. 
It conducts heat and electricity well. When away from air contact it 
does not alter, even at the highest temperatures. Heated in an oxygen 
atmosphere it usually burns, with more difficulty than the diamond, to 
carbon dioxide, leaving a residual ash of 2-5 per cent. For the 
purification of the poorer, more impure kinds of graphite, the latter is 
pulverized, and heated with a mixture of IvC103 and H2S04 ; the pro- 
duct is washed with water and the residue glowed (Brody’s Graphite). 
Artificially, graphite may be obtained by fusing amorphous carbon 
with iron; on cooling the latter a portion of the dissolved carbon 
separates in hexagonal, shining leaflets. 
3. Amorphous carbon is produced by the carbonization of organic 
(containing carbon) substances, and is found in a fossilized state. 
Nitric acid and potassium chlorate in the cold convert it into brown 
substances soluble in water. The purest amorphous carbon is soot, 
which is obtained by the imperfect combustion of resins and oils (like 
turpentine) rich in carbon. Gas carbon, also called metallic carbon, 
which deposits in the manufacture of gas in the glowing tubes, is very 
hard, possessing metallic lustre, and conducting electricity well, hence 
used in galvanic batteries. Coke, resulting from the ignition of stone 
coal, forms a sintered mass, conducting heat and electricity well. 
Charcoal is very porous and can absorb many gases and vapors. 1 
vol. of it absorbs 90 vols. NH3, 55 vols. H2S, and 9 vols. 0 ; at 100° 
and under the air-pump the absorbed gases are again liberated. Char- 
coal will also take up many odorous substances and decaying matter, 
hence employed as a disinfectant. Animal charcoal, obtained by the 
carbonization of animal matter (bones, blood, etc.), possesses in a high 
degree the power of removing many coloring substances from their 
solutions ; hence it serves, in the laboratory and practice, for the decol- 
orization of dark solutions. 
All these varieties of carbon contain, in smaller or larger quantities, 
nitrogen, hydrogen and mineral substances, which, in combustion, re- 
main behind, as ash. Hydrochloric acid will withdraw almost all the 
mineral constituents. 
The fossil coal varieties, stone coal, lignite, turf, are the products 
of a peculiar, slow decay of wood fibre, which gradually separates 
oxygen and hydrogen, and enriches itself in carbon. Fossil coal con- 
tains 90 per cent., brown coal, 70 per cent., of carbon. The fossil 
coal richest in carbon, the last product of the alteration, is anthracite, 
which has lost all its organic structure, and contains 96-98 per cent, 
of carbon. 142 
INORGANIC CHEMISTRY. 
COMPOUNDS OF CABBON WITH HYDBOGEN. 
With hydrogen, carbon forms an unlimited number of com- 
pounds, into which, also, all other elements, especially oxygen 
and nitrogen, can enter. These carbon compounds have been 
termed organic compounds, because formerly they were exclu- 
sively obtained from vegetable and animal organisms, and the 
idea was entertained that they were produced by the influence 
of forces other than those producing the mineral substances. 
At present, the most carbon derivatives are artificially pre- 
pared from the elements by simple synthetic methods ; we are 
aware that they do not essentially differ from mineral sub- 
stances. 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 carbon compounds are so 
numberless, and possess so many peculiarities, that, from di- 
dactic reasons, it appears necessary to treat them apart from 
the other compounds, in a separate portion of chemistry, 
which we, pursuing the old custom, term organic chemistry. 
In opposition to this, the chemistry of all other bodies is 
designated Inorganic Chemistry. Hence, 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, whereby acetylene 
(C2H2) results, may be effected. All other hydrocarbons are 
obtained indirectly, in various ways. 
Methane—Marsh Gas—CH4. This simplest hydrocarbon, 
containing but one atom of carbon, is formed in the decay of 
organic matter under water (in swamps and coal mines), and 
escapes in large quantities in many regions of the earth (thus 
at Baku, on the Caspian Sea). Synthetically, it may be 
formed, among other methods, by conducting vapors of carbon 
disulphide and hydrogen sulphide over glowing copper filings: 
For its preparation heat a mixture of sodium acetate with 
sodium hydrate :— 
CS2 + 2H2S + 8Cu =4Cu2S + CH4. 
Methyl is a colorless, odorless gas, insoluble in water ; it 
has not yet been condensed. Ignited, it burns with a faintly 
C2H3Na02 + NaOH — CII4 + Na2C03. METHANE. 
143 
luminous flame. With two volumes of oxygen (or 10 vols. 
of air) it gives a violently explosive mixture (fire damp of the 
miners):— 
CH4 + 20 2 = C02 + 2H20. 
1 vol. 
2 vols. 
1 vol. 
2 vols. 
Molecular Formula of Methyl. Atomic Weight of Carbon. 
The quantitative analysis of methyl shows that for every 1 
part of hydrogen in it there are 3 parts 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 (H,0) then carbon would equal 6, etc., (see 
(p. 87):— 
H = 1 
C = 3 
CH = 4 
2H = 2 
C = 6 
CH, = 8 
311= 3 
C= 9 
CH3 = 12 
4H •■= 4 
C = 12 
CH4 = 16 
In this case the analysis yields (as in former instances) 
no conclusive answer. We cannot either derive the atomic 
weight of C from its vapor density, as it is not volatile, but it 
may be arrived at from the density of methane. The latter 
equals 8 (H = 1) hence the molecular weight is 16. In 16 
parts by weight of methyl there are 4 parts by weight, hence 
4 atoms of hydrogen and 12 parts carbon. The atomic 
weight of C is, then, presuming that only one atom of it is 
present in methane, 12 :— 
4 atoms hydrogen H4 = 4 
1 atom carbon C = 12 
Methyl molecule CH4 = 16 
That indeed the atomic weight of carbon is 12, is proven by 
the fact that of all its innumerable derivatives not one con- 
tains less than 12 parts, by weight, of this element. It 
follows, with certainty, from the periodic system of elements 
(p. 70;. 
From the formula CH4 it follows that in 1 volume meth- 
ane, 2 volumes of hydrogen are present, (CHt contains 2H2). 
1 TOl. 2 vols. 
This is indirectly proven by the combustion of methyl with 
oxygen in a eudiometer (see p. 111). Four atoms of hydro- 
gen yield two molecules of TLO; 1 atom of C gives 1 mole- 144 
INORGANIC CHEMISTRY. 
cule C02. Hence the volume relation in the combustion of 
CH4 in oxygen is expressed by the equation :— 
CH4 + 20. = C02 + 2H20 
In two volumes of aqueous vapor there are 2 volumes of 
hydrogen ; hence in one volume CH4, 2 volumes of H. The 
result of the eudiometric analysis confirms these conclusions. 
1 vol. 
2 vols. 
1 vol. 
2 vols. 
Ethyl —C2 H5. Formed when hydrogen in statu nascendi 
acts upon ethyl chloride :— 
Or by the action of potassium or sodium upon methyl iodide 
C2H5C1 + 2H = C2II6 + HC1 
2CH3I -f- Na2 — C2H6 -f- 2Na I2. 
This is a colorless gas, insoluble in water, which ignited burns 
with a feebly luminous flame. Its density equals 15 (II =1) 
or 1.036 (air = 1) corresponding to the molecular formula C2 
IIfi = 30. 
Besides methyl (CH4) and ethyl (C2H6) there exists a long 
series of hydrocarbons of the general formula CnH2n -f- 2, 
(e. g., C3H8, C4H]0, C5H12 etc.,) in which each member differs 
from the preceding and next following by 1 C and 2 H (CH2). 
Bodies belonging to such a series greatly alike in their chemi- 
cal behavior are termed homologues. In addition to this series 
of saturated hydrocarbons exist others, with less quantity of 
hydrogen, and which, by the addition of the latter, pass into 
the saturated, and may hence be termed unsaturated. The 
first unsaturated series is composed according to the formula 
CDH2n, the second according to CnH2n_2, etc. The lowest 
member of the series CnHm is ethylene (see Chemical Structure 
p. 162). 
Ethylene—C2H4. Formed in the destructive distillation of 
wood, stone coal and many carbon compounds, hence con- 
tained in illuminating gas. Most easily obtained by the 
action of sulphuric acid upon alcohol, whereby the acid with- 
draws H20 from the alcohol:— 
A lcohol 
c2h6o-h2o = c2h4 
Ethylene. 
It is a colorless gas, of weak, ethereal odor, condensing at 
— 110° to a liquid. Density equals 14 (H = 1) or 0.978 (air 
= 1), corresponding to the molecular formula C2H4 = 28. It NATURE OF FLAME. 
145 
burns with a bright, luminous flame, at first being decomposed 
into marsh gas and free carbon :— 
The CH4 burns and heats the particles of coal in the flame 
to incandescence ; these then are consumed to carbon dioxide 
(C02). 
c2h4 = ch4 + c. 
The unsaturated compound, ethylene, unites directly with 
two atoms of chlorine and bromine :— 
C2H4 + Cl2 - C2H4C12. 
The resulting compounds C2H4C12 and C2H4Br2 are oily 
liquids; hence the name olefiant gas, for ethylene. 
The lowest member of the second unsatprated series is 
C2H2. 
Acetylene—C2H2—is produced in the dry distillation of 
many carbon compounds, and is present in coal gas, whose 
peculiar penetrating odor is due to this. Density = 13 (H 
= 1) corresponding to the formula C2H2 = 26. It combines 
directly with 2 and 4 atoms of chlorine and bromine. 
The three hydrocarbons above considered, methyl (CH4), 
ethylene (C2H4), and in slight amount (C2H2), constitute, 
together with H and carbonous oxide (CO), ordinary 
illuminating gas, produced in the dry distillation 
of stone, brown coal or wood. The illuminating power 
is influenced by its quantity of ethylene and acetylene 
(and their homologues). 
THE NATURE OF FLAME, 
We are aware that every chemical union which occurs in a 
gaseous form and accompanied by the evolution of light is 
designated combustion. In this we observe that some bodies, 
like sulphur and phosphorus, when burned in the air or in 
some other gas, yield a flame. These are those substances 
which, at the temperature of combustion, are converted into 
gases or vapors. Carbon burns without a flame, with only 
a glow, because it is non-volatile. The carbon compounds, 
wood, stone coal and lard, are, indeed, not even volatile, but 
burn with a flame, because under the influence of heat they 
develop combustible gases. Flame is, therefore, nothing more 
than a combustible gas heated to glowing, and which is 
continually formed anew. We observed, too, that the com- 146 
INORGANIC CHEMISTRY. 
bustibility was only a relative phenomenon ; if hydrogen burns 
in oxygen and chlorine, conversely oxygen and chlorine burn 
in hydrogen (p. 47). 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 be very plainly 
illustrated by means of the following contrivance. An ordin- 
ary lamp globe (Fig. 76) may be closed at its lower end with 
a cork, through which enter two 
tubes; the narrow tube, a, somewhat 
contracted at its end, is connected 
with a gas stop-cock ; the other 
tube, b, (best a cork borer) is about 
5 mm. wide, and communicates with 
the air. The gas issuing from the 
tube a is ignited, and the chimney 
is then dropped over the not too 
large flame; it continues to burn 
along quietly, as sufficient air enters 
through the wide tube b. Upon in- 
creasing the supply of gas, the flame 
becomes larger, the globe constantly 
fills more with illuminating gas, 
while air is crowded out; now the 
gas flame is extinguished, and upon 
the wider tube, b, an air-flame ap- 
pears, as the entering air continues 
to burn farther in the atmosphere 
of illuminating gas. The excess of 
the latter pouring out from the upper portion of the 
globe may be ignited, and we then have above a gas flame, 
while within the globe we have an air flame. On lessen- 
ing again the gas flow, the air flame will distribute itself, 
extend to the exit of the tube a, and then upon the latter will 
appear again the gas flame, while the flame above the globe 
is extinguished. In this, at will, we may repeat the return 
process of flames. That the air in the air flame actually 
burns, may be plainly proven if we introduce a small gas 
flame, c, through the wide metallic tube, b ; the little flame 
will continue to burn in the air flame, but will be extin- 
guished, if it be introduced higher up into the atmosphere of 
illuminating gas. 
Fig. 76 NATURE OF FLAME. 
147 
Ordinarily we designate only such bodies as combustible 
which, because they are capable of combining with oxygen, 
burn in an atmosphere of oxygen or of atmospheric air. If 
we imagine, however, an atmosphere of hydrogen, or illumi- 
nating gas, then, in such an one, bodies rich in oxygen must 
be combustible. In fact, nitrates, chlorates, etc., burn in an 
atmosphere of gas with production of an oxygen flame. This 
may readily be demonstrated, as follows: An argand lamp 
chimney (Fig. 77) is closed at its lower end by a cork, bearing 
a gas-conducting tube. The gas, issuing above, through the 
Fig. 77. 
opening of the sheet covering, a, is ignited. Now the substance 
(potassium or barium chlorate, etc.) is introduced into a flame 
by means of an iron spoon provided with a handle, heated to 
the temperature of decomposition (disengagement of oxygen), 
and the spoon then plunged through the opening into the gas 
atmosphere. The substance now burns with brilliant disen- 
gagement of light, as the resulting oxygen flame is brightly 
colored by the vaporizing and reduced metallic salts. 
The brilliancy or luminosity of a flame is influenced by the 148 
INORGANIC CHEMISTRY. 
nature of the substances contained m it, also by its tempera- 
ture and density. Glowing gases, especially in dilute condi- 
tion, shine very faintly per se. Thus hydrogen, ammonia, and 
methyl burn with a pale flame. Even sulphur burns in the 
air with a slightly luminous flame. If, on the contrary, sul- 
phur or phosphine be permitted to burn in 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 nitrogen of the air, therefore is more condensed, devel- 
ops a higher temperature, and the combustion products (S02, 
P205, PC13) or the evaporating substances are not immediately 
gasified. That the density of the flame of gases is of important 
influence upon the luminosity is proven by the fact that also 
hydrogen, with oxygen, compressed into a contracted space, 
burns with intense light display. 
A slightly luminous flame may be rendered intense by 
introducing solid particles into it. For example, if hydrogen 
be passed through liquid chromium oxychloride (Cr202Cl2) it 
burns v7ith a bright luminous flame, because the volatile 
Cr202Cl2 in it is changed by the oxygen of the flame into 
solid, non-volatile chromium oxide, Cr203, whose particles are 
heated to glowing by the hydrogen flame. Similarly is 
explained the illuminating power of the various hydrocarbons 
and carbon compounds. Marsh gas, CIT4, and ethyl, C2H5, 
afford a pale flame, because 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 de- 
composes first into CH4 and carbon, whose particles 
glow7 in the flame. (See p. 150.) 
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 fibres and con- 
verted into gaseous hydro-carbons, which ignite, 
and by chemical union with the oxygen of the air, 
produce the flame. In the latter, three distinct 
zones may be perceived. In the inner non-volatile 
zone, a a' (Fig. 78), are contained unaltered gases, 
which, owing to lack of air access, cannot burn. 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. In the middle, brightly luminous part, 
Fig. 78. NATURE OF FLAME. 
149 
f e g, a partial combustion of these occurs; ethylene C2H4 
breaks up into CH4 and C: the first burns completely, while 
the C is heated to a white glow7, as not sufficient oxygen is 
present for its combustion. The presence of carbon particles 
in the luminous part may be easily proven by placing in it a 
cold body, like a wire net; it will at once be coated with soot. 
Finally, in the outer, only very feebly luminous and almost 
invisible mantle, bed, of the flame, which is completely sur- 
rounded by air, the perfect combustion of all the carbon to 
carbon dioxide occurs. 
An entirely identical structure is possessed by the ordinary 
illuminating gas flame. By bringing as much air or oxygen 
into it as is necessary for the perfect 
combustion of all the carbon, none 
of the latter separates (see below), 
and there is produced a faintly 
luminous but very hot flame. Upon 
this principle is based the construc- 
tion of the Bunsen burner, the flame 
of which is employed in laboratories 
for heating and ignition. Fig. 79 
represents a form of the same. The 
upper end, e, is screwed into the 
lower portion, and in the figure is 
only separated for the sake of ex- 
planation. The gas enters through 
the narrow opening, a, from the side 
gas tube, and mingles with air in 
the tube e, which makes its entrance 
through the openings of the ring b. 
In this way only a slightly non-luminous but intensely heating 
flame is obtained. On closing the openings in b, the air is cut 
off, and the gas burns at the upper end of the tube e, with a 
bright, strongly smoking flame. The non-luminous flame con- 
tains excess of oxygen, and hence oxidizes—oxidizing flame. 
It is employed to effect oxidation reactions. The luminous 
flame, on the other hand, is reducing in its action, and is 
designated the reduction flame, because the glowing carbon in 
it abstracts oxygen from many substances. 
Upon the same occurrences depends the construction and 
application of the ordinary and the gas blowpipe, which, 
however, at present, in most cases, are replaced by gas 
burners. 
Fig. 79. 150 
INORGANIC CHEMISTRY. 
The non-luminosity of the Bunsen burner flame, due to addition of 
air, depends on a more complete combustion of the separated carbon 
or ofthe yet undecomposed hydrocarbons. The flame, in consequence, 
is smaller, more intense, and the combustion extends itself even to the 
inner cone of the flame. The non-luminosity becomes more difficult in 
case of pure oxygen, as the flame is then not diluted by nitrogen, 
therefore it is far smaller, the temperature much higher, and the flame 
gases are more condensed. 
Another variety of non-luminosity of hydrocarbon flames is induced 
by the admixture 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. In consequence of 
the dilution there are present fewer combustible particles in an equal 
space, and therefore can be more completely consumed by the more 
readily intruding oxygen of the air; further, the temperature is lowered 
and attains probably not the decomposition temperature of eth}dene 
(C2H4) in the adjoining, continually renewing inner cone. Also by sim- 
ple extension of an illuminating flame upon a plate it may be rendered 
non-luminous, since then the air comes in contact with a larger flame 
plane. On heating a gas made non-luminous by admixture of nitrogen, 
and then letting it burn, its flame illuminates as the increased temper- 
ature may induce the decomposition of ethylene. 
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:— 
C02 + C = 2C0. 
Indeed, but a few per cent, of C02 in a gas flame suffices to decrease 
its luminosity, 
c2h4 + co2 = ch4 + CO 
1 vol. 
1 vol. 
1 vol. 
2 vols., 
while the presence of nitrogen is far less detrimental. 
Every substance requires for its ignition a definite tempera- 
ture—temperature of ignition. When a substance is once 
ignited it generally burns further, because by the heat of com- 
bustion more particles are raised to the temperature of igni- 
tion. By rapid cooling (e. g., by introduction of a piece of 
metal into a small flame) every flame may be extinguished. 
By holding a metallic net over the opening of a gas lamp, 
from which gas issues, and igniting the same above the wire 
(Fig. 80), 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 phenomenon depends the construction 
of Davy’s safety lamp, which is used in coal mines, to avoid 
ignition of the fire-damp (Fig. 81). 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 explo- 151 
CARBON WITH THE HALOGENS. 
Fig. 81 
Fig. 80. 
sive mixture, or into a combustible gas (e. g., in a large jar, 
in which ether is present), the gas penetrating into the lamp 
will burn, but the combustion will not extend to the external 
gases. 
COMPOUNDS OF CARBON WITH THE HALOGENS. 
Carbon does not combine directly with the halogens; the 
compounds result, however, by the action of the halogens upon 
the hydrocarbons. We have seen that chlorine and bromine 
act upon water, ammonia, H2S, PH3, etc., in such a manner 
as to unite with the hydrogen to hydrogen chloride, etc., while 
the other element is either set free or is also united with the 
chlorine. Chlorine and bromine act similarly upon the 
hydrocarbons; here hydrogen, atom after atom is displaced 
by chlorine, forming HC1 and chlorine derivatives:— 
CH4 -f Cl2 = CH3C1 +HC1 
CH4 + 2 Cl = CH2C12 + 2 HC1, etc. 
Such a process is termed substitution, and the products substi- 
tution products. In this way we obtain from methyl, CH4, 
CH3C1, CH2C12, CHC13 (chloroform) and finally CC14, carbon 
tetrachloride. The last compound is a colorless, ethereal 
liquid, boiling at 77°. Vapor density equals 77 (H == 1), cor- 
responding to the molecular formula CC14 == 154. The chlo- 
rides of carbon are not decomposed by water. 152 
INORGANIC CHEMISTRY. 
The compound C2C16, ethyl hexachloride, obtained by the 
action of chlorine upon ethyl, C2H6, is a crystalline mass, fusing 
and boiling at 182°. On conducting the vapors through a 
red-hot tube they decompose into C2C14 and Cl2. Ethylene 
tetrachloride, C2C14, is a liquid, boiling at 122° C. Bromine 
and iodine yield similar compounds; they will be treated more 
extensively in organic chemistry. 
2. SILICON. 
Si = 28. 
Next to oxygen this is the most widely distributed element 
in nature. Owing to its affinity for the former it does not 
occur in free condition. In combination with oxygen as silicon 
dioxide (Si02), and in form of salts of silicic acid (silicates) 
it comprises many minerals and almost all crystalline rocks. 
Obtained in a free condition by heating the silicon fluoride 
(SiFl4), or sodium-silicon fluoride (Na2SiFlo) with metallic 
sodium:— 
Na2SiFl6 + 4Na = 6NaFl + Si, 
The ignited mass is treated with water, which dissolves the 
sodium fluoride and leaves the silicon as a brown, noil- 
lustrous, amorphous powder. Heated in the air it burns to 
silicon dioxide (Si02). 
Another modification—the crystalline silicon—is obtained by 
fusing a mixture of Na2SiFl6, sodium and zinc. The separated 
silicon dissolves in the molten zinc, and on cooling, deposits 
out in crystals, which remain on dissolving the metal in hydro- 
chloric acid. In this form the metal exhibits black, strongly 
lustrous octahedrons, of specific gravity, 2.49, and of very great 
hardness. Upon ignition in the air or oxygen it is not oxid- 
ized ; it is not attacked by acids. On boiling it with a sodium 
or potassium hydrate solution it dissolves, forming a silicate 
and liberating hydrogen:— 
Si + 4K0H = K4Si04 + 2H2. 
Heated in chlorine gas, silicon burns to the chloride. 
Hydrogen Silicide—SiH4—the analogue of CH4, forms 
like arsine by dissolving an alloy of silicon and magnesium 
in dilute hydrochloric acid :— 
SiMg, + 4HC1 = SiH4 + 2MgCl2. 
The escaping hydride contains admixed hydrogen, has a SILICON CHLORIDE 
153 
disagreeable odor, ignites spontaneously in the air, and burns 
to the dioxide and water :— 
Entirely pure silicide, free of H, is obtained by heating a 
compound, SiH (O. CJiib belonging in the province of 
organic chemistry. It only ignites in the air at ordinary 
pressure upon warming; if, however, the gas, by diminution 
of pressure or by the addition of H, is diluted, it becomes 
spontaneously combustible at ordinary temperatures. Glow- 
ing heat decomposes the hydride into amorphous silica and 
hydrogen. Mixed with chlorine it inflames and probably 
forms substitution products similar to those of methyl 
Pure hydrogen silicide condenses, at —5° and 70 atmospheres, 
to a liquid. 
Silicon Chloride—SiCl4—results from the action of chlorine 
upon silicon, or by conducting chlorine over a mixture of the 
dioxide and carbon (Fig. 82):— 
SiH.t + 20 2 = Si02 + 2H20. 
Si02 + 2C + 2C12 = SiCl4 + 2C0. 
Fig. 82. 
The mixture is placed in a porcelain tube, which is heated to 
a red heat in a charcoal furnace. The chlorine generated in 
the flask is washed in a three-necked bottle and dried in a 
glass tube filled with calcium chloride. While carbon or 
chlorine do not act upon the Si02, the reaction or their simul- 
taneous action is induced by the mutual supporting affinities 
of carbon for oxygen and of chlorine for silicon. 154 
INORGANIC CHEMISTRY. 
The silicon chloride which distills over is a colorless 
liquid, having a specific gravity of 1.52, and boiling at 57°. 
In the air it fumes, and is decomposed by water into silicic 
and hydrochloric acids:— 
SiCl4 + 4H20 = II4Si 04 -f 4HC1. 
This compound may be employed in determining the atomic 
weight of silicon. Analysis shows that in it, for 35.4 chlorine, 
there are seven parts of silicon. Supposing, from the great 
analogy of the silicon compounds to those of carbon, that its 
formula is SiCl4 the atomic weight of the silicon would be 28. 
Si = 28 
Cl* = 141.6 (4 X 35.5) 
Si Cl 4 = 169.6 
This supposition is confirmed by the vapor density of the 
compound. This equals 85 (H = 1), hence the molecular 
weight is 2 X 85 = 170. Since, according to the analysis, there 
are 142 parts chlorine in 170 parts of silicon chloride, the 
atomic weight of the metal must be 28. 
Silicon Bromide—SiBr4 and Silicon Iodide—Sil4, are 
formed in the same manner as the chloride. The first is a 
colorless liquid, of specific gravity 2.8, becoming solid at—12° 
and boiling at + 153°. The iodide forms colorless octahe- 
dra, fusing at 120°, and boiling at 290°. Like the chloride, 
both are decomposed by water. 
Besides these compounds, which may be viewed as hydro- 
gen silicide, in which all the hydrogen is replaced by halo- 
gens (see page 151), others exist, in which only a part of this 
element is replaced. Thus, silicon chloroform, SiHCl3 cor- 
responds to the chloroform (CHC13) derived from methyl. 
It is produced by the action of phosphorus pentachloride, or 
antimony chloride, on hydrogen silicide:— 
SiH4 + 3SbCl5 = SiHCl3 + 3SbCl3 + 3HC1. 
Also upon heating silicon in dry hydrogen chloride gas: 
in this instance a mixture of SiCl4 and SiHCl3 results. These 
compounds may be separated by fractional distillation. Silicon 
chloroform is a colorless liquid, of specific gravity 1.6, and 
boils at 85-37°. The vapor density equals 67.7 (H = l), 
corresponding to the molecular formula, SiHCl3 = 135.5. It 
fumes in the air, and with water decomposes into silicic and 
hydrochloric acids. HYDROGEN SILICON-FLUORIDE. 
155 
Very similar to silicon chloroform are the bromoform, 
SiHBr3, and the iodoform, SiHI3, which correspond to the 
analogous compounds obtained from carbon. 
The compound Si2I6, analogous to C2C16, is known. From 
all these data we observe the great analogy between silicon 
and carbon. 
Silicon Fluoride, SiFl4, is formed when HF1 acts upon 
Si02:— 
Si02 + 4HF1 = SiFl4 + 2H20 
To prepare it, a mixture of fluorite and powdered glass, or 
sand (Si02), is warmed with sulphuric acid ; by the action of 
the H2S04 upon the fluorite hydrogen fluoride (p. 55) is dis- 
engaged, which then reacts upon the silicon dioxide, in the 
manner indicated in the above equation. The liberated gas 
is collected over mer- 
cury. It is colorless, 
has a disagreeable odor, 
and fumes strongly in 
the air. Its vapor den- 
sity is 3.57 (air = 1), 
or 52, (H = 1), cor- 
responding to the mole- 
cular formula SiFl4 = 
104. Very character- 
istic is its deportment 
with water, by which 
it is decomposed into 
silicic acid (H4Si04) 
and hydrogen silico- 
fluoride : 3SiFl4 
+ 4H20 = II4Si04 + 
2H2SiFl6. For the ex- 
ecution of this method, 
conduct the SiFl4 formed, through a glass tube, into a vessel 
containing water (Fig. 83). Gelatinous silicic acid separates 
out, and the gaseous SiH2FlB remains dissolved in the water. 
As the separating silicic acid may readily obstruct the open- 
ing of the glass tube, the latter is allowed to project a slight 
distance into mercury. By filtration the solid silicic acid is 
separated from the aqueous solution. 
Hydrogen Silico-Fluoride—H2 Si Fl6 (or 2HF1, Si Fl4) is 
only known in aqueous solution. Upon evaporating at a low 
Fxo. 83. 156 
INORGANIC CHEMISTRY. 
heat it decomposes into Si Fl4 and 2HF1. In its chemical 
deportment it is an acid similar to the hydrogen-halogen 
acids. Its aqueous solution reddens blue litmus paper, dis- 
solves many metals, and saturates bases, forming salts with 
them, in which two hydrogen atoms are replaced by metal. 
The potassium and barium salts are insoluble in water. 
To the same group as that of carbon and silicon belongs, 
as regards its chemical nature, TiN=Sn= 118, which forms 
perfectly analogous compounds, like SnCl4, SnBr4, SiiF14. It 
bears the same relation to these elements as antimony to those 
of the nitrogen group. This fact is plainly visible in the 
atomic weights :— 
N = 14 
C =12 
P = 31 
Si =28 
As = 75 
Sb = 122 
Sn = 118 
The element of the carbon group corresponding to the 
arsenic of the nitrogen group, if it indeed does exist, is unknown. 
Like antimony, but to a greater degree, tin exhibits a 
decidedly metallic character, and forms the transition to the 
real metals. According to its physical properties tin is a true 
metal: with hydrogen (unlike the metalloids) it does not 
unite. Therefore we will treat it, together with its higher 
analogue, lead, with the metals. 
In the preceding pages we have considered four groups of 
elements, which comprise all the so-called metalloids (with the 
exception of boron). In each group the last members, pos- 
sessing the highest atomic weights, exhibit already distinct 
metallic properties, especially when in a free condition. This 
is plainly the case with tin, antimony and arsenic. Tellurium 
and selenium, also (in the crystalline modification) possess 
marked metallic appearance; finally, iodine has a metallic 
lustre. With addition of metallic character diminishes, on 
the other hand, the affinity for hydrogen; such compounds 
of iodine, tellurium, antimony and arsenic are very unstable 
and readily decompose into their constituents; finally, tin and 
bismuth do not combine with hydrogen. 
The remarkable relations between the atomic weights of the 
elements of the four groups in which, at the same time, the 
deportment of the elements finds its expression, are perceived 
from the following table:— ATOM AND MOLECULE. 
157 
C = 12. 
Si = 28. 
Sn = 118. 
N = 14. 
P = 31. 
As = 75- 
Sb = 122. 
0 = 16. 
S = 32. 
Se= 79. 
Te = 125. 
FI = 19. 
Cl= 35.4. 
Br= 79.7. 
I = 126.5. 
These relations will find more consideration in the exami- 
nation of the Periodic System of the Elements. 
ATOM AND MOLECULE. 
In the study of the hydrogen derivatives of the elements of 
the four previously mentioned groups, we arrived at the es- 
tablishment of the following formulas :— 
CH4. 
SiH4. 
NH3. 
ph3. 
AsH3. 
SbHs. 
0H2. 
sh2. 
SeH2. 
TeH2. 
F1H. 
C1H. 
BrtL 
IH. 
As these possess a fundamental importance, and serve for 
the derivation of further, more important conclusions and 
generalizations, we again introduce together, here, all the 
actual relations and considerations in connection, upon the 
basis of their deduction. 
The fact, proven by the analysis and synthesis of chemical 
bodies, that the elements combine with each other according 
to constant multiple proportions, explains itself most simply 
by the supposition of elementary atoms, which can com- 
bine among themselves in different quantity (page 59). 
The existence of isomeric bodies can only be concluded from 
the atomic constitution of matter. To determine the number 
of atoms in a compound, and, therefore, also, the true relative 
atomic weights of the elements, the experimental combining 
weights, however, offer no footing (page 79); the question 
can only be solved upon a basis of other actual relations. 
The specific gravity of bodies in the gaseous and vapor 
form, and the volume relations according to which the gases 
combine, serve best for this task. 
The physical properties of the gases and vapors lead to the 
supposition that they consist of very small discrete particles, 
molecules, which are separated from each other by relatively 
large but like distances, and that, therefore, in equal volumes 
of all gases an equal number of molecules is contained. 
Hence molecules are the smallest particles of matter—masses 
which occur separated in gases. Their relative weights, there- 
fore, are directly expressed by their relative gas densities 158 
INORGANIC CHEMISTRY. 
(page 70). Molecules of compound bodies are composed 
of atoms; therefore, the relative atomic weights may be 
derived from the molecular weights. The atomic weight is the 
smallest quantity of an element which is contained in the molecular 
weight of any one compound. Although the atomic weights of 
the above considered elements have been deduced from a small 
number of compounds, yet all further investigations have 
confirmed the observation that these elements are not con- 
tained in a smaller quantity in any molecule. Thus oxygen 
is present in all molecules, in at least 16 parts by weight; 
nitrogen not less than 14 ; carbon not less than in 12; if the 
smallest quantity of hydrogen contained in a molecule, i. e., 
an atom, equals 1. 
From a comparison of the densities of the elements with 
those of their derivatives we concluded (p. 68)—upon the 
basis, that in equal volumes of gases the same number of 
molecules is contained—that the molecules of the elements 
consist of two or more atoms. Thus the molecules of hydro- 
gen, the halogens, oxygen, nitrogen, are composed of two atoms ; 
the sulphur molecules at 500° of six atoms; the molecules of 
phosphorus and arsenic of four atoms :— 
h2 ci2 o2 n2 S2 P4 As4. 
The molecular quantities of all bodies in the gaseous state 
occupy equal volumes. Referring the densities (specific 
gravities) to hydrogen as unit (=1), the molecular weights 
are double the densities, and conversely, the latter are half 
the molecular weights (p. 69). The existence of allotropic 
modifications of the elements confirms (p. 7'8) the hypo- 
thesis that the molecules of the elements are composed of 
several atoms. In the gaseous condition the allotropy is 
known only in the case of oxygen and sulphur. Ordinary 
oxygen has two atoms, ozone three; the sulphur molecule, 
at 1000°, contains two atoms ; at 500°, howrever, six atoms. 
Although we possess no means of determining the molecular 
size of the solid elements, yet there are many indications that 
the free elements occurring in several modifications, like P, 
As, C, and Si, consist of larger complex atoms. Further, the 
energetic action of the elements in the moment of their 
formation (see p. 69) argues for their complexity when free. 
The phenomena of the so called status-nascens find, in the 
majority of cases, their explanation in the thermo-chemical processes 
accompanying them. If, e. g., in the action of zinc upon nitric acid, ATOM AND MOLECULE. 
159 
where, in accordance with the analogy with other acids, hydrogen 
must be set free, the same does not occur, but the nitric ac d is 
reduced to ammonia and nitrogen oxides—while free hydrogen does 
not exert any action on nitric acid—then it is the heat disengaged 
by the formation of zinc and ammonium nitrate, which, in accordance 
with the principle of greatest heat liberation, represents the true 
cause of the reaction. By this, however, the supposition that the free 
atoms act more energetically than the molecules is not refuted ; it 
finds an additional confirmation in thermo-chemistry. 
The determination of the density is the simplest means 
of ascertaining the molecular quantity; but we have 
another, a pure chemical procedure, leading to the same 
end. As molecules are the smallest quantities which can 
exist in a free condition, it is very probable the same 
quantities appear in the chemical reactions. Indeed, the 
study of the latter brings us the same molecular quanti- 
ties as are derived from the densities. For example, the 
compound CH3C1 results when chlorine acts upon marsh 
gas; we hence infer that in the molecule of the latter hydro- 
carbon, four atoms of H are present:— 
CH4 + Cl2 = CHgCl -1- HC1. 
From ethylene—C2H4—we obtain C2H3C1; its composition, 
therefore, cannot be expressed by the simpler formula CH2. 
With iodine ammonia yields NIII2; with hydrogen chloride 
NH4C1:— 
NH3 + 21, = NHI2 + 2HI 
NH3 + HC1 = NH4C1. 
Both reactions lead to the formula NH,. In the same 
manner the compound nature of the elementary molecules is 
disclosed ; in all accurately determined reactions we perceive 
that hydrogen, the halogens, oxygen, nitrogen, invariably act 
or separate out with two atoms, as is observed from the fol- 
lowing equations :— 
Mn02 + 4HC1 = MnCl2 + 2H20 + Cl2, 
NH4N02 = 2H20 + N2, 
Ag20 + H202 = Ag2 + H20 -j- 02. 
This chemical method for obtaining the molecular value is, 
although less simple, so much more general than the one 
depending on the determination of the densities, inasmuch 
as it finds application in the case of the non-volatile bodies. 
Thus, for example, the composition of hydrogen peroxide is 
expressed by the formula HO ; the methods of its production 160 
INORGANIC CHEMISTRY. 
and chemical rearrangements point, however, with great 
probability, to the doubled molecular formula H202 
Hence the most varied relations carry us to the same con- 
clusions as to the existence and the magnitude of the atoms 
and molecules. The atomic-molecular formulas used to day 
in chemistry are only the expression of actual relations; they 
are entirely independent of speculative abstractions, but 
forcibly call out these. 
From the above derived molecular formulas of the hydro- 
gen compounds, arose further important generalizations. We 
have the four following groups of hydrogen compounds (forms 
of combination or types) :— 
The Quantivalence of the Elements. Chemical Structure, 
CH4 — NH3 — 0H2 — F1H 
SiH4 — PH3 — SH2 — C1H 
— AsH3 —SeH2 — BrH 
— SoH3 — TeH2 — 1H 
In each group, step by step, with increasing atomic weight, 
the affinity of the elements for hydrogen diminishes. Yet 
the number of hydrogen atoms which are combined with 
an atom of the other elements is constant for each group. 
Hence, to each element, in its relation to hydrogen, we must 
ascribe a particular function of affinity, called quantivalence, 
valence or atomicity. The elements of the fluorine group are 
univalent or monatomic; the elements of the oxygen group 
diatomic or bivalent; nitrogen and its analogues, triatomic; 
carbon and silicon, finally, are tetratomic elements—if we 
accept the quantivalence of hydrogen as the unit of valence. 
The halogens, combining with one atom of hydrogen, possess 
one affinity unit; oxygen combines with two hydrogen atoms, 
and possesses, therefore, two affinity units, etc. 
The quantivalence of the elements is frequently designated 
by the lines or Roman numerals placed above the symbols :— 
i 
Cl 
ii 
0 
in 
N 
IV 
C 
The mutual union of affinity units is indicated by a line 
connecting them:— 
II 
I 
N 
/ \ 
H H 
H 
I 
H—C—H 
I 
H 
H—Cl 
H-O-H 
Hydrogen 
chloride. 
Water 
Ammonia. 
Methane. QUANT1 VALENCE. 
161 
In these formulas the atoms of oxygen, of nitrogen, of 
carbon—indeed, of all the monatomic elements—represent, as 
it were, the nuclei to which the hydrogen atoms attach them- 
selves ; their valence units are, as it were, the points of 
attachment for the valence unit of hydrogen. 
In these molecules the hydrogen atoms can now be replaced 
or substituted by other elements (p. 151). According to this 
the monatomic halogen atoms each replace a hydrogen atom. 
H 
ii / 
O 
* \ 
H 
H 
ii / 
O 
\ 
Cl 
ci 
II / 
o 
\ 
Cl 
H 
in / 
N-I 
\ 
I 
Cl 
III / 
Sb-Cl 
\ 
Cl 
W ater. 
Hypochlorous 
acid. 
Chloriae 
oxide. 
Nitrogen 
iodide. 
Antimony 
trichloride. 
Instead of the lines we can more conveniently employ 
brackets :— 
fH 
IV 1 ni 
Cjc! 
I Cl 
Cl 
IV J Cl 
c Cl 
L Cl 
fH 
IV j Cl 
Si] Cl 
L Cl 
Chloroform. 
Carbon 
tetrachloride. 
Silicon 
chloroform. 
By the replacement of hydrogen by the monatomic metal 
potassium, we arrive at the compounds— 
K—Cl 
H 
/ 
O 
\ 
K 
fH 
N] K 
(K 
fH 
C | H 
K 
Pot. chloride. 
Di-pot. amide. 
Potas. methyl. 
The diatomic elements, like oxygen, sulphur, replace two 
hydrogen atoms in the compounds of that element:— 
Pot. hydrate. 
Sb=o 
H 
%0 
C\ o11 
£ / 0” 
Si \ 0n 
Antimony 
oxychloride. 
Carbon 
dioxide. 
Silicon 
dioxide. 
Methylene 
oxide. 
Finally, triatomic nitrogen can replace three atoms of hy- 
drogen, or of the halogen :— 
IV f N111 
C\H 
or, 
I IV 
H—C=N 
iv r Nm 
C\ci 
or, 
Hydrogen cyanide. 
I iv III 
Cl—C=N 
Chlorcyanogen. 162 
INORGANIC CHEMISTRY. 
In all these compounds the quantivalence peculiar to the 
elements appears. 
Like valence is designated witli the word equivalence. 1 atom of 
chlorine is equivalent to an atom of hydrogen: 35.4 parts, by weight, 
of chlorine are, then, equivalent to 1 part, by weight, of hydrogen. 1 
atom oxygen is equivalent to 2 atoms of hydrogen ; consequently, 2 
parts, by weight, of hydrogen, are equivalent to 16 parts, by weight, 
of O, or 1H to 8 parts of oxygen. Further, 1 atom of N, or 14 parts, 
is equivalent to 3 atoms or 3 parts hydrogen; 1 part hydrogen is, 
therefore, equivalent to 134 = 3.66 parts nitrogen, etc. These quanti- 
ties, equal to 1 part, by weight, of H are termed equivalent weights, 
and were formerly employed instead of the atomic weights. As is per- 
ceptible from the preceding, the equivalent weights are parts of the 
atomic weights corresponding to the valence units of the atoms. 
If, consequently, the valence of the elements, in relation to 
hydrogen (as also to other elements), has a definite value, the 
question naturally arises—what will result, if an atom of hy- 
drogen be withdrawn from the saturated molecules, e. g., water, 
H20, ammonia, NH3 or methane—CH4. The groups or resi- 
dues— 
— 0— H 
Hydroxyl 
II 
— S —H 
Sulphydrate 
III 
— nh2 
Amid 
IV 
- c - h3 
Methyl 
can plainly not exist in a free condition, as one affinity-unit 
of the elements, having higher atomicity is not saturated. 
When free these groups, therefore (like the elementary atoms), 
unite with the free affinities and enter into complicated com- 
pounds. Thus result, for example, the bodies— 
HO — OH 
Hydrogen 
peroxide 
II II 
HS —SH 
Hydrogen 
persulphide 
III III 
h2p- ph2 
Liquid 
hydrogen 
phosphine 
IV IV 
h3c-ch3 
Dimethyl 
or 
ethane. 
To such combination carbon is particularly inclined. 
Removing an atom of H from dimethyl or ethane (C2H6) the 
so-called group ethyl remains—- 
c2h5 
or 
IV IV 
CHS CH2 
in which one carbon affinity is unsaturated ; this can again 
unite with the methyl group CH3. The resulting compound 
is— 
By the continuation of this process, as it were, of a chain- 
like union of the carbon atoms, we obtain a whole series of 
hydrocarbons (C4H10, C5H]2, etc.), with the general formula 
OIi-2n +2 (compare page 144). 
c3h8 
or 
IV IV IV 
H3C — CH2 — CH3. CHEMICAL STRUCTURE. 
163 
Not only similar residues or groups combine, but also 
dissimilar:— 
hi ii 
H.N — OH 
IV II 
H3C —OH 
IV in 
h3c — nh2. 
Such combinations are generally effected by reactions of 
double decomposition. Thus methyl hydroxide (wood spirit) 
results from the action of methyl iodide (CHjI) upon silver 
hydrate (AgOH):— 
Hydroxyl amine 
Methyl hydroxide 
Methylamine. 
CH3I + AgOH = Agl + CH3OH 
methylamine by the action of methyl iodide and ammonia 
Dimethyl is produced when sodium acts upon methyl 
iodide:— 
CII3I + NH3 = CH3NH2 + HI 
2CH3I + Na2 = 2NaI + C2H6. 
By the withdrawal of sodium the methyl groups are liber- 
ated, and these then combine with each other. 
Further the atoms, with several atomicities, can unite with 
two and three affinities (double and triple union) :— 
III HI 
N = N 
\ / 
O 
IV IV 
H — C = CH2 
IV IV 
HC = CH 
Kthylene 
(p. H4.) 
Acetylene 
(p. 145.) 
Hyponi trous 
oxide, 
By complete mutual union result the free elements:— 
P = P 
I I 
P = P 
H —H 0 = 0 1ST — N 
Hydrogen. 
Oxygen. 
Nitrogen. 
This manner of joining or combination of the atoms in the 
molecule, not with their entire mass, but with single affinities, 
is designated chemical constitution or chemical structure of com- 
pounds; the formulas representing them are called constitution 
or structural formulas. Naturally the actual position of atoms 
in space (of which we have no knowledge) is not indicated by 
the chemical structure. The fundamental principle of chemical 
structure consists in this, that the affinity of anatom unites itself 
with the affinity of another atom. The following circumstances, 
however, complicate these simple relations: Among the 
elements of the nitrogen group we saw that P and N combine 
with 3 and 5 atoms of chlorine and the halogens ; that sulphur, 
selenium and tellurium take up 2 and 4 atoms of chlorine and 
bromine; that iodine unites with 1 and 3 atoms of chlorine 164 
INORGANIC CHEMISTRY. 
and five of fluorine. Only the tetratomic elements, carbon 
and silicon, are capable of combining with four hydrogen, 
and also not more than four atoms of the halogens:— 
IV 
CC14 — 
III 
PC13 
PC15 
II 
SC12 
IV 
SC14 
IC1 
III 
IC13 
V 
IF1S 
Hence, it appears that the elements (excepting carbon) in 
their relation to chlorine (and to the halogens) do not express 
such a constant balance as they do to hydrogen. P and its 
analogues appear tri- and pent-atomic ; the elements of the 
sulphur group, di- and tetr-atomic; iodine finally appears 
mon-, tri- and pent-atomic. 
This higher affinity of the metalloids shows itself more dis- 
tinctly and universally in the more stable oxygen derivatives. 
In the elements of the four previously considered groups we 
are acquainted with the following oxygen compounds, in 
which the members of each group afford perfectly analogous 
compounds :— 
IV 
co2 
III 
N2y03 
- n!o5 
II 
SC12 
IV 
- so2 
VI 
so3 
C120 
III 
- ci203 
- i2o5 
VII 
I207 
Here the affinity of iodine (and of the halogens) reaches 
seven, and the elements of the sulphur group six affinities. The 
elements of the nitrogen group are, both in their relation to 
chlorine and oxygen, not more than pentatomic. Carbon 
finally extends, both to H as to Cl and 0, not more than four 
affinities. These relations are more apparent, as we will see, 
in the hydroxyl derivatives of the oxides and the acids. Hence 
we conclude, that the equivalence is not an absolute property 
belonging per se, to and for the elements solely (like the atomic 
weights) but that it appears as a relative function of the mutual 
action of the various elements. In general we can distinguish 
the hydrogen valence and the halogen or oxygen valence. For 
all elements the hydrogen valence is constant: for Cl = 1, for 
O = 2, for N = 8, for C = 4. As regards oxygen, the quan- 
tivalence of the most of the elements appears to be variable; and 
indeed it varies, as visible from the above given formulas, for 165 
MOLECULAR COMPOUNDS. 
the Cl group from 1 to 3 to 5 and to 7 ; for the elements of 
the S group from 2 to 4 to 6; for P and its analogues from 3 
to 5. The regular increase of maximum valence from C = 4 
to Cl = 7 argues the actual occurrence of a change of valence 
whether we ascribe a variable valence to the elements or 
accept the maximum as the true measure, and regard the 
lower compounds as unsaturated, is entiAly immaterial, as 
we possess no conception upon the nature of valence. Later 
we will discover that this alteration of valence finds full expres- 
sion and generalization in the periodic system of the elements, 
which is based upon the grouping of the elements according 
to their atomic weights. 
In the manner just represented the idea of quantivalence is viewed 
purely from an empirical standpoint. Another opinion prevails, 
which, denying the alterability, regards the valence as an absolute con- 
stant property of the elementary atoms. According to this idea, the 
true valence, or atomicity, is only derivable from the hydrogen com- 
pounds; the halogens are absolutely monatomic, the elements of the 
0 group diatomic, and those of the N group triatomic, etc. To carry 
out the constant atomicity, for all the compounds different 
suppositions are made. To explain the oxygen compounds a linking 
union of the diatomic oxygen atoms, in the same manner as the union 
of the C atoms in carbon compounds, is assumed, visible from the fol- 
lowing:— 
i ii ii ii i 
Cl 0 0 0 Cl 
Chlorine trioxide. 
I II II II II I 
Cl 0 0 0 0 H 
Perchloric acid. 
II ,0 
s< I 
xo 
/O. 
H/ \ 
S 0 
ii /0-0-H 
S\ 
x0-0-H 
II III II II II III II 
0 = N—0 — 0 — O — N = 0 
Sulphur dioxide. 
Sulphur trioxide. 
Sulphuric acid. 
II III II II II 
0 = N— 0 — 0 — H 
Hence it would appear that the oxygen atom? can unite without end, 
in a chain-like manner, similar to the C atoms in the carbon com- 
pounds. However, there actually occurs a maximum combining affin- 
ity of different groups for oxygen, which varies regularly from group to 
group (for Cl — 7, for S = 6, for N = 5, for C = 4). Consider it as 
we may, on the supposition of constant affinity, the reason for the dif- 
ferent number of linking 0 atoms must be based on the nature of the 
o'her element. The circumstance, then, that the highest oxides or 
their hydrates (HC104. H2S04, HN'Oj) are more stable than the 
lower (p. 176), argues decidedly against the chain-like union of the 
oxygen atoms, as occurs in the unsaturated peroxides. 
To explain other compounds according to the constant atomicity 
theory, a difference between atomic and molecular compounds is noted. 
Nitrogen pentoxide 
Nitric acid 166 
INORGANIC CHEMISTRY. 
The former are such as can be explained by constant atomicity. All 
others are regarded as molecular compounds, resulting from the addi- 
tion of two or more entire molecules, upon the basis of newly supposed 
molecular affinities. Thus the compounds PC15, SC14, IC13 are viewed 
as addition products from atomic combinations with chlorine mole- 
cules :— 
PCI,. Cl, 
SC12, Cl2, 
ICl, Cl, 
Phosphorus 
pentachloride 
Sulphur 
dichloride 
Iodine 
trichloride. 
It was thought that a proof that these bodies were differently con- 
stituted from the real atomic compounds, was seen in 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 mole- 
cules PC15, SC14 is only gradual, increasing with the temperature, and 
that they do exist undecomposed, at lower temperature, in a vapor 
form (compare p. 130). True atomic compounds, like H2S04 and 
HN03 frequently separate into simpler molecules, in their conversion 
into gas (compare Sulphuric Acid). 
Especially demonstrating for the pentavalence of phospho- 
rus is the, at ordinary temperatures, gaseous pentafluoride, 
PF15, also, that iodine affords a volatile pentafluoride, IF15. 
It is noteworthy that, generally, the metalloids with the 
lower halogens (fluorine and chlorine) yield more stable and 
higher compounds than with bromine and iodine, which 
possess a higher atomic weight (p. 44). The idea that mo- 
lecular compounds can also exist in vapor form, would mark 
their distinction from the atomic compounds as purely arbi- 
trary—not present in the nature of the bodies themselves. 
Other gaseous compounds also exist, which in no light can be 
regarded as molecular. Thus, the usually tetratomic or hexa- 
tomic tungsten (WC14, WOCl4) forms a gaseous pentachloride, 
WC15, and molybdenum, perfectly analogous to tungsten, 
yields a pentachloride, MoC15. Further, the pentatomic van- 
adium (VdOCl3) gives rise to a gaseous tetrachloride, VdCl4. 
The salts of ammonia, according to the theory of constant 
atomicity, are not regarded as ammonium compounds 
v 
(NH4C1, NO,NH4, S(NH4)4 
y 
y 
(page 120), but as addition products of ammonia with acids 
—NH3, HC1, NH3, N03H, (NH3)2, S04H2—whereby the 
analogy of the same with metallic salts appears enigmatical. 
Further, the properties of many compounds, like 
y 
y 
iv 
iv 
v 
P0C13, (C2H5)3PO, (CH3)3S.OH, (CH3)2SO, (CH3)4N.OH, MOLECULAR COMPOUNDS. 
167 
and others, can scarcely be interpreted by the constant val- 
ence theory. The existence, also, of potassium permanganate, 
Mn04K, is not compatible with a constant di- or tetravalence 
of manganese. 
Indeed, up to the present, the acceptance of molecular additions 
could not be entirely dispensed with, especially for the so-called water 
of crystallization compounds ; the proposed effort, however, success- 
fully continues to derive all such compounds on the basis of higher 
quantivalence of the elements. While, of course, the constant valence 
theory comprises only the so-called atomic compounds, the extended 
valence idea draws all others into the circle of generalizations. 
Before all, we must be mindful that the nature of chemical union and 
the cause of the valence of the atoms is entirely unknown to us, and 
that, therefore, neither the idea of a variable, nor yet of a constant 
quantivalence constitutes a final explanation. So long as we do not 
possess an hypothesis upon the real causes, our task can only confine 
itself to collecting the varying combination relations of atoms into 
single points of view. The supposition of a constant valence, 
w’hich formerly was given preference as the simpler, has shown itself as 
insufficient. Because it maintains real differences, where such are not 
perceptible, it departs from the ground of induction. The idea of a 
variable valence, considering all facts, is not prejudicial. Indeed, it 
finds a pregnant analogy in the deportment of hydrocarbon radicals. 
By elimination of hydrogen from the border molecules (as C2H6) arise 
radicals or groups of increasing valence (C2H5, C2II4, C2H3, C2H2, 
), just as the valence 'increases from fluorine to carbon 
(FI =19; 0 = 16; N =14 ; C = 12). The group C2H2, however, is 
di-and tetravalent; the group C2H3, mono- and trivalent. Every- 
thing, however, indicates that the chemical elements do not represent 
final individuals, but are composed of one or several primary sub- 
stances. 
The principles of chemical structure, above presented, show 
themselves most distinctly, and with regularity, in carbon. 
The constitution of the innumerable varieties of carbon com- 
pounds explains itself by the tetratomic nature of the carbon 
atoms, and their ability to combine themselves with each 
other by a simple affinity. In other, the so-called inorganic 
compounds, the valence and structure relations are more com- 
plicated, and are far less investigated, but even in them so 
many regularities show themselves that thereby the actual 
material is greatly simplified, and is made more comprehen- 
sible. The theory of valence and structure is the first at- 
tempt to refer the facts underlying the law of multiple pro- 
portions to the functions of the elementary atoms. As this 
doctrine only comprises actual relations, it cannot be negated, 
but only further developed. 168 
INORGANIC CHEMISTRY. 
The platform of the theory of atomicity and structure was 
laid down by A. Kekule, (1857-1859). It constitutes a rep- 
resentation and further development of the type theory of 
Gerhardt, at the basis of which lies, unrecognized, the idea of 
the different valence of the atoms. 
OXYGEN COMPOUNDS OF THE METALLOIDS. 
Almost all the oxygen derivatives of the metalloids are of 
an acid-forming nature ; with water they yield acids:— 
I207 + H20 — 2HT04 ; S03 + H20 = H2S04 
Iodine 
heptoxide. 
P2O5 ~j~ 3H20 — 2H3P04 
Periodic 
acid. 
Sulphur 
trioxide. 
Sulphuric 
acid. 
Conversely, these oxides may be formed by the removal of 
Avater from the acids; therefore, they are ordinarily termed 
anhydrides of the corresponding acids; I207—anhydride of 
periodic acid; S04 as sulphuric-anhydride; P205 as phos- 
phoric anhydride, etc. 
By the rejolacement of the hydrogen of the acids by metals 
salts arise. According to the number of hydrogen atoms re- 
placable by metals, acids are- distinguished as 
dibasic, tribasic or monohydric, dihydric, trihydric, etc. 
Phosphorus 
peutoxide. 
Phosphoric 
acid. 
Like the metalloids the metals are of various atomicities ; 
the monatomic metals (Na, K, Ag,) replace each one hydrogen 
atom; those of greater atomicity replace more. Therefore 
the metals of higher atomicity can unite several residues. 
From this the importance e. g., of the following chemical 
formulas is explained:— 
S04Ca 
Calcium sulphate. 
k2so4 
Pot. sulphate 
S04KNa 
Pot. sod. sulphate 
no3k 
Pot. nitrate 
(N03)2Cu 
Copper nitrate 
hi 
(N03)sBi 
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 indifferently written at the beginning 
or end of the chemical formulas of the salts; in the second 
case we mean to indicate that the metal atoms are in union 
with the oxygen. OXYGEN COMPOUNDS OF THE HALOGENS. 
OXYGEN COMPOUNDS OF THE HALOGENS. 
Excepting fluorine, all the halogens combine with oxygen 
to form anhydrides and acids entirely analogously constituted, 
although not all the members are known for each halogen. 
The acids have one atom of hydrogen replacable by metals, 
and hence are monobasic. Chlorine forms the following 
anhydrides and acids:— 
Anhj'd rides. 
C120 
C1203 
(C120J 
(Cla07) 
Acids. 
HC10 —Hypochlorous acid. 
HC102—Chlorous acid. 
HClOj—Chloric acid. 
HCIO4—Perchloric acid. 
The anhydrides C1205 and C1207 (corresponding to the 
compounds I205 and I207) are unknown. The compound 
Cl2Ot exists and must be regarded as a mixed anhydride of 
chlorous and chloric acid. 
The following formulas express the chemical structure of 
these compounds: — 
I II I 
Cl-O—Cl 
Hypochlorous 
anhydride. 
I II 
Cl—0-H. 
Hypochlorous 
acid. 
Ill III 
OCl-O—CIO 
III 
(CIO)-OH. 
Chlorous anhydride. 
Chlorous acid. 
o2ci—0—cIo2 
Chloric anhydride. 
(C102)—OH. 
Chloric acid 
VII VII 
03C1—0—C103 
(C103)—OH. 
Perchloric anhydride. 
III V 
OCl-O—C102 
Perchloric acid. 
In the acids we assume the presence of the monotomic 
group OH (hydroxyl or water residue) in which hydrogen, by 
action of metals or bases, can be replaced by metals. The 
acid groups (C102 or C103) combined with hydroyxl are called 
acid residues or radicals. In the anhydrides two acid radicals 
are united by an oxygen atom ; by the action of water they 
decompose into 2 molecules of acid :— 
Chlorous-chloric anhydride. 
Ill III 
OCL 11 CIO—OH 
111 /O -f- 0< = hi 
OCK 1 XH CIO—OH 170 
INORGANIC CHEMISTRY. 
The salts of perchloric acid are called perchlorates; those of 
chloric acid, chlorates ; those of chlorous acid, chlorites ; those 
of hypochlorous acid, hypochlorites. 
Hypochlorous Oxide—C120—Hypochlorous anhydride, is 
produced by conducting dry chlorine gas, in the cold, over 
precipitated and dried mercuric oxide:— 
Hg0 + 2Cla = HgCl2 + Cl20. 
The disengaged gas is condensed in a bent glass tube, cooled 
by a freezing mixture. 
Hypochlorous oxide is a red brown liquid, resembling chlo- 
rine, boiling at +20°, and passing into a yellow vapor. The 
vapor density is 43.5 (H = 1), corresponding to the formula, 
C120 = 87. The oxide is very unstable, and in the course of 
a few hours decomposes, yielding chlorine and oxygen. It 
has strong oxidizing and bleaching properties. Dissolves in 
water to hypochlorous acid:— 
Hydrochloric acid decomposes it into water and chlorine:— 
C120 + H20 = 2 HOC1. 
Hypochlorous Acid—HCIO—is only known in aqueous 
solution. It is obtained by conducting chlorine into water 
containing suspended mercuric oxide. The concentrated 
solution is yellow in color and is decomposed by light. It 
oxidizes and bleaches energetically. The bleaching action of 
this acid, 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:— 
C120 + 2 HC1 = H20 + 2 Cl2. 
Cl2 +H20=2HC1 + 0 
2C10H = 2HC1 + 02. 
The acid itself is very feeble and incapable of decomposing 
carbonates. Its salts (Bleaching powder, see Chloride of 
Lime) are formed by the action of chlorine, in the cold, upon 
strong bases:— 
2NaOH + Cl a = NaCl + NaOCl + H20. 
Upon heating their solutions with dilute nitric acid the free 
acid distills over. 
On shaking the aqueous solution of hypochlorous acid with mercury, 
there is produced a white precipitate of HgO, HgCl2, soluble in hydro- 
chloric acid (salts of hypochlorous acid form HgO). This behavior CHLORINE TETROXIDE. 
171 
serves to distinguish hypochlorous acid from chlorine, which under 
like circumstances forms Hg2Cl2, insoluble in hydrochloric acid 
(Reaction of Wolter). 
Chlorine Trioxide—C1203—Chlorous anhydride. This 
results from the deoxidation of chloric acid, as e. g., when a 
mixture of KC103, nitric acid and reducing substances like 
arsenic trioxide is warmed— 
a yellowish green gas escapes, which can be condensed at 
—20°. Chlorine trioxide is a reddish-brown liquid, boiling 
about 0° and decomposing rapidly. Its vapors, when heated 
to 50°, explode with violence. In water the oxide is soluble, 
forming chlorous acid C102H or CIO.OH, unknown in a free 
condition. With alkalies it yields salts called chlorites. 
Chlorine Tetroxide—C1204—is the mixed anhydride of 
chloric and chlorous acid, as water and the alkalies, decom- 
posing it into these two acids:— 
2HC103 = C1203 + H20 + 02, 
gg }o +H 20 = CIO.OH + C102.0H 
Chlorous acid. 
Chloric acid. 
It forms when sulphuric acid acts upon potassium chlorate 
(KC103); by slight heating, a dark yellow gas escapes, which, 
at —20° condenses to a reddish-yellow liquid, boiling at -j-9°. 
The oxide is an energetic oxidizing agent, very unstable, and, 
in daylight, readily decomposes, with violent explosion. 
Hence, we must avoid mixing sulphuric acid with potassium 
chlorate. The formation of it can be effected in a perfectly 
harmless way, and its powerful oxidizing action be illustrated, 
by throwing some potassium chlorate and a few pieces of yel- 
low phosphorus into water, contained in a measuring glass, 
then, by means of a pipette, allowing sulphuric acid to touch 
the bottom of the tube, drop by drop. By the action of the 
tetroxide, which is disengaged, the phosphorus will burn 
under water with a brilliant light. 
When concentrated sulphuric acid is added to a mixture of 
potassium chlorate and sugar, violent combustion occurs. 
The vapor density of the oxide is 33.7 (H = l); the gaseous mole- 
cules, therefore, possess the formula C102 (=67.5). It is very prob- 
able, that, at lower temperature, the molecules have a double formula, 
C1204 ; this is confirmed by the manner in which water decomposes 
it, and by its perfect analogy with nitrogen tetroxide, for which a 
dissociation of the molecules N204 into N02 has been proven. 172 
INORGANIC CHEMISTRY. 
Chloric Acid—HC103, or C102.OH—is obtained by decom- 
posing an aqueous solution of barium chlorate, by means of 
sulphuric acid:— 
(C103)2Ba + S04H2 = BaSO4 + 2 HC103 
Barium chlorate. 
Barium sulphate. 
The barium sulphate separates as a white insoluble powder, 
and can then be filtered off from the aqueous solution of 
chloric acid. This is then concentrated, under an air-pump, 
until the specific gravity is 1.28 ; when about 40 °Jc chloric 
acid is present; it is oily and, when warmed to 40°, decom- 
poses into chlorine, oxygen and perchloric acid, HC104. The 
concentrated aqueous solution oxidizes strongly; sulphur, 
phosphorus, alcohol and paper are inflamed by it. Hydro- 
chloric acid eliminates chlorine from the acid and its salts : 
HClOg + 5 HC13 = 3H20 + 3C12 
The chlorates are produced, together with chlorides, by the 
action of chlorine, in the heat, upon many bases (compare 
Potassium chlorate):— 
Perchloric Acid—HC104, or C10:i0H. Of all the oxygen 
derivatives of chlorine this is the most stable. As previously 
stated, it forms by the decomposition of chloric acid ; more 
easily obtained from its salts. Upon heating potassium chlo- 
rate to fusion, oxygen escapes and potassium perchlorate 
results:— 
6 KOH + 3 CIa = 5 HC1 + KC103 + 3 H20. 
Upon warming the perchlorate with four parts sulphuric 
acid, perchloric acid distills over :— 
2KC103 = 2KC104 + KC1 + 02. 
2C104K + H2S04 = K2S04 + 2HC104. 
Pure, the acid is a mobile, colorless liquid, fuming strongly 
in the air ; its specific gravity is 1.78 at 15°. It boils at 
110°. It cannot be preserved, since after a few days it 
decomposes with violent explosion. In contact with phos- 
phorus, paper, carbon and other organic substances, it also ex- 
plodes. Upon the skin it produces painful wounds. Dissolves 
in water with hissing and with one molecule of the solvent 
forms the crystalline hydrate HC104 + H20, fusing at 50°; 
the crystals fume in the air and gradually deliquesce. The 
second hydrate—HC104 -f 2ILO—is a thick, oily liquid, 
resembling sulphuric acid, and boiling unchanged at 208°. 
Also obtained by evaporation of the aqueous solutions of IODIC ACID. 
173 
perchloric and chloric acids. When the crystalline hydrate 
is distilled it breaks up into anhydrous perchloric acid and 
the second hydrate :— 
2C104H.Ha0=C104H + C104H.2H20. 
Bromine yields the following oxygen compounds 
HBrO Hypobromous acid. 
HBr03 Bromie acid. 
HBr04 Perbromic acid. 
The corresponding anhydrides are not known. The acids 
are perfectly analogous to the corresponding chlorine com- 
pounds. 
Hypobromous Acid — HBO—is formed when bromine 
water acts upon mercuric oxide ; the aqueous solution can be 
distilled in vacuo, and possesses all the properties of its chlo- 
rine analogue. 
Bromic Acid—Br03H. Bromates are formed by the 
action of bromine, in the heat, upon the aqueous solution of 
the alkalies or of barium hydrate ; from the barium salts the 
acid may be obtained in aqueous solution by liberation by 
means of sulphuric acid. More practically is the acid 
procured by the action of bromine upon silver bromate or 
oxidation of bromine by means of hypochlorous acid :— 
The aqueous solution may be concentrated in vacuo until its 
quantity reaches 50.6 % BrO:tH and then closely corresponds 
to the formula Br03H -f- 7H20. When heated it breaks 
up into bromine, oxygen and water. 
Perbromic Acid—BrO,H—is said to be formed in the action 
of bromine vapor upon perchloric acid :— 
5C120 + Br2 + H20 = 2Br03H + 10C1. 
and perfectly similar to the latter. 
C104H + Br = Br04H + Cl 
Iodine forms the following anhydrides and acids. 
I^05 — HI0s — Iodic Acid. 
I207 — HI04 — Periodic Acid. 
Iodic Acid—HI03. Its salts (iodates) are formed similarly 
to those of chloric and bromic acids, by dissolving iodine in a 
hot solution of potassium or sodium hydrate:— 
6K0H + 8I2 = 5KI + IK03 + 3H20. 174 
INORGANIC CHEMISTRY. 
The free acid can be obtained by the oxidation of iodine with 
strong nitric acid, or by the aid of chlorine; further, by the 
action of iodine upon chloric or bromic acids, whereby 
the iodine directly eliminates the chlorine and bromine :— 
2HC103 + I2 = 2HI03 +CI2. 
Upon evaporating the aqueous solution the free iodic acid 
crystallizes in colorless rhombic tablets of specific gravity 
4.63. The solution oxidizes strongly. Heated to 1703 iodic 
acid decomposes into water and iodic anhydride:— 
2HI0S =I205 + H20. 
Decomposed, like chloric acid, by hydrochloric :— 
2I03H + 10HC1 = I2 + 5C12 + 6H20. 
By reducing agents (H2S, S02, HI) it is reduced to iodine. 
Similar decompositions are sustained by periodic acid. 
Iodic Anhydride—1205—forms a white crystalline powder, 
which dissolves in water to iodic acid. At 300° it decom- 
poses into iodine and oxygen. 
Periodic Acid—HI04. Produced by the action of iodine 
upon perchloric acid :— 
Upon the evaporation of the aqueous solution, the acid crys- 
tallizes out with two molecules of water (HI04, 2H20—com- 
pare Sod. periodate). In the air the crystals deliquesce, fuse 
at 130°, and at a higher temperature decompose into water 
and periodic anhydride, the latter readily breaking up into 
oxygen and iodic anhydride :— 
2HC104 + I2 =2HI04 + Cl2 
2(HI04) + 2H20 = I205 + 02 + 5H20. 
The existence of the hydrates of periodic and perchloric acids, as 
also of many others (see Sulphuric and Nitric acids), which formerly 
we were accustomed to view as molecular compounds (p. 165), are 
interpreted at present by the acceptance of hydroxyl groups, directly 
combined with the element of higher equivalence :— 
VII 
CIO4H -(- H20 = C102(0H)3— trihydrate or tryhydric acid 
VII . 
CIO4H + 2H20 = CIO (OH)s—pentahydrate or pentahydric acid 
VII 
C104H + 3H20 = Cl (0H)7—heptahydrate or heptahydric acid 
The extreme hydrates (C1(0H)7 and I(OH)7, in which all seven 
affinities of the halogen atom are attached to hydroxl groups, are not 
known, but probably exist in aqueous solution. As they give up water PERIODIC ACID. 
175 
and one atom of 0, becoming simultaneously united with two bonds to 
the halogen, they yield the lower hydrates—even to the monohydrate 
C1030H. Just as in the monohydrate, perchloric acid is monobasic 
in the polyhydrates, since but one hydrogen atom is replaced by metals : 
On the other hand, periodic acid (I030H is not only monobasic, but 
as a pentahydrate (IO(OH)5) can, similar to the polybasic acids, also 
furnish polymetallic salts, as— 
C106H5 + KOH = C104K + 3H20. 
i(°H)3 
10 ( (ONa)2 
vn f(°H)# 
IOt(OAg), 
VII 
10(0, Na)5 
VII 
I0(0Ag)5 
Salts also exist which are derived from condensed polyiodic acids, 
as— 
.(OH)* 
I0\ 
>° 
I0\ 
x(0H)4. 
—Diperiodic acid, etc. 
(Compare disulphuric, dichromic acid, etc.) 
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 in their highest combinations are, in fact, heptads. 
The oxygen compounds of the halogens in some respects 
display a character exactly opposite to the hydrogen deriva- 
tives. While the affinity of the halogens with reference to 
hydrogen diminishes with increasing atomic weight from FI 
to I (see page 56), the degrees of affinity for oxygen is the 
exact reverse. Fluorine is not capable of combining with 
oxygen ; the chlorine and bromine compounds are very un- 
stable and in free condition generally not known ; the iodine 
derivatives, on the contrary, are the most stable. In accord 
with this is the fact that in the higher oxygen compounds 
chlorine is liberated directly by bromine, and bromine by 
iodine, while in the hydrogen and metallic compounds of the 
halogens the direct reverse—that iodine and bromine are 
removed by chlorine—is the case. 
Further, the oxygen compounds exhibit the remarkable 
peculiarity, that their stability increases with the addition of 
oxygen. The lowest acids HC10, IIBrO, HC102, are very 
unstable, even in their salts ; they possess a very slight acid 
character, and are, too, separated from their salts by carbon 
dioxide. The most energetic and most stable are the highest 176 
INORGANIC CHEMISTRY. 
acids, IIC104, HBr03, HI03, in w'hich the higher valence of 
the halogens appears. The corresponding oxygen compounds 
of the sulphur and nitrogen groups are perfectly similar—a 
property scarcely to be connected with the supposition of a 
linking grouping of the oxygen atoms (according to the 
constant atomicity theory, see p. 165). 
The peculiar behavior of the oxygen compounds of the 
halogens, their variable stability and decomposition, as also 
their modes of formation, find nearer explanation in their 
thermo-chemical relations. All oxide compounds of chlorine 
and bromine are endothermic, i. e., in their production from 
the elements heat is rendered latent (compare p. 57). They 
do not result, therefore, by direct union of the elements; fur- 
ther, they are little stable, decompose readily with elimination 
of oxygen, and then strongly oxidize. The heat, appearing 
in the formation of chlorine monoxide, and of the hypothet- 
ical pentoxides, C1205 and Br205 (in their production from the 
elements and solution in water), corresponds to the symbols:— 
(C12,0 — gas) = — 18040; (Cl2,05,Aq) = — 20480; 
(Br2,05,Aq) — — 43500. 
In the formation of iodine pentoxide and iodic acid, heat is 
liberated :— 
This explains its stability in comparison with the chlorine 
and bromine compounds, also the direct production of iodic 
acid by the oxidation of iodine. When the pentoxides are 
compared with each other, it is seen that, in the formation of 
bromine pentoxide, Br205, the most heat is rendered latent— 
the affinity of bromine for oxygen, consequently, is the lowest, 
that of iodine the greatest. The same is also evident from the 
heat produced from acids, in dilute aqueous solution, or of 
the potassium salts in solid condition:— 
(I2,05) = + 44,860; (I,03,H) = + 57,880. 
(Cl,0,H,Aq) — 28,940; (Br.03H,Aq) = 12,420; (I.03,H,Aq) = 
55,710; (Cl,OsK) = 94,600; (Br,03,K) = 87,000 ; 
(I,03,K) = 128,400. 
From this is understood, that chlorine and bromine are 
separated by iodine, with formation of iodic agid, from chloric 
and bromic acids, while bromine does not act upon chloric 
acid. Later, we will see that also in the group of sulphur 
and of phosphorus, the middle members, selenium and arsenic, 
exhibit a less liberation of heat in their oxygen compounds— 
their affinity, therefore, is slighter. OXYGEN COMPOUNDS OF SULPHUR. 
177 
2. OXYGEN COMPOUNDS OF THE ELEMENTS OF 
THE SULPHUR GROUP. 
The elements sulphur, selenium and tellurium combine 
with two atoms of II, and yield oxygen acids, which also con- 
tain 2 H atoms :— 
In these acids 1 and 2 atoms of H can be replaced by metals ; 
hence they are dibasic. By replacement of one atom of H 
are formed the so-called acid or primary salts, while, by the 
replacement of both hydrogen atoms are obtained the neu- 
tral or secondary salts :— 
H2S S02H2 h2so3 h2so4. 
so4kh 
so4k2. 
Acid potassium sulphate. 
Neutral potassium sulphate. 
1. OXYGEN COMPOUNDS OF SULPHUR. 
(S02H3) 
so2 
Hyposulphurous acid. 
so3h2 
Sulphurous auhydride. 
Sulphuric anhydride. 
so3 
Sulphurous acid. 
S04H2 
Sulphuric acid. 
In addition to these, so to speak, normal compounds, exist 
still others, more complicated, which will be studied later. 
The structure of these derivatives may be expressed in the 
following formulas:— 
II 
o = s = o 
Sulphur Dioxide. 
VI 
o=s = o 
II 
o 
IV /OH 
0 = S\ 
OH* 
Sulphurous acid. 
„ V. /OH 
0 = s( 
II XOH 
0 
Sulphur Trioxide. 
Sulphuric acid. 
Sulphur Dioxide, S02, or sulphurous anhydride, is formed 
by burning sulphur or sulphides in the air:—S = S02. 
* The structure of sulphurous acid must probably be expressed by 
the formula, II—S02—OH, according to which 1 atom of H is con- 
nected with sulphur, but the other is contained as hydroxyl. This 
appears from the carbon derivatives of sulphurous acid. Probably, in 
compounds, both structural cases exist, as two isomeric series of neutral 
ethers of the acid are known. 178 
INORGANIC CHEMISTRY. 
The combustion may also be effected by the action of 
metallic oxides (copper oxide, manganese peroxide) which 
give up their oxygen readily. It is most conveniently pre- 
pared for laboratories by heating sulphuric acid with mercury 
or copper:— 
The acid is similarly decomposed by heating it with carbon 
2H2S04 + Cu = CuS04 + S02 + 2H20. 
Copper Sulphate. 
2S04H2 + C = 2S02 + C02 + 2EI20. 
By this method we get a mixture of carbon and sulphur diox- 
ides, which are separated with difficulty. Owing to its solu- 
bility in water, sulphur dioxide must be collected over mer- 
cury. 
It is a colorless gas, with a suffocating odor. Its density is 
32 (H = I), corresponding to the molecular formula S02 = 
64. It condenses at —15°, or at ordinary temperatures 
under a pressure of two atmospheres, to a colorless liquid, of 
specific gravity 1.45, which solidifies at —76° and boils at 
—10°. Upon evaporation, the liquid sulphur dioxide absorbs 
much heat; on pouring liquid S02 upon mercury, in a clay 
crucible and accelerating the evaporation by blowing air upon 
it, the metal will solidify. Water dissolves, with liberation 
of heat, 50 volumes of sulphur dioxide gas, which 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. In dry con- 
dition the gases combine; if their mixture be conducted over 
feebly heated platinum sponge * sulphur trioxide results :— 
2S0Z + 02 = 2S03. 
2 vols. 2 vols. 
In aqueous solution the dioxide slowly absorbs O from the 
air, forming sulphuric acid :— 
so2 +h2o + 0 = h2so4. 
The oxidation of the aqueous sulphur dioxide to sulphuric 
acid proceeds more rapidly by the action of Cl, Br and I:— 
In consequence of the affinity of the halogen for hydrogen 
and of sulphurous acid for oxygen the decomposition of a 
S03H2 + H20 + Cl2 = S04H2 + 2HC1. 
* Instead of platinum sponge, platinized asbestos may be applied ; 
this is obtained by immersing asbestos in a platinic chloride solution, 
then in ammonium chloride, and afterwards drying and igniting. SULPHUROUS ACID. 
179 
molecule of water is here effected. 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. Upon this property rests its applica- 
tion in the bleaching of wools and silks, which are strongly 
attacked by the ordinary chlorine bleaching agents (p. 43). 
By stronger reducing agents the dioxide may be deoxidized, 
thus by H2S sulphur is separated out:— 
S02 + 2H2 S= 2H20 + 2S. 
If, however, both gases are strongly diluted by other neutral 
gases, the action is but very slow. 
A mixture of equal volumes S02 and Cl2 unite in direct sunlight 
to thionyl chlorine S02 Cl2 (p. 188). When sulphur dioxide acts 
upon warmed phosphoric chloride, there is formed with the oxychlo- 
ride the compound SOCl2 : — 
S02 + PC15 = P0C13 + S0C12. 
Chlorthionyl--SOCl2—may be viewed as sulphur dioxide in 
which one atom of 0 is replaced by two atoms of chlorine- It is a 
colorless liquid with a sharp odor, and boils at 78°. Water decom- 
poses it into hydrogen chloride and sulphurous acid : — 
.OH 
SOCl2 + H20 = SO ( + 2HC1. 
xOH 
Sulphurous Acid—H2S03—is not known in free condition, 
but is probably present in the aqueous solution of S02. On 
cooling the concentrated solution to 0°, colorless cubical crystals 
separate with the composition (S02 -|- 15H20) or (S03H2 + 
14H20). Upon longer standing of the aqueous solution, 
especially in sunlight, sulphur separates with formation 
of sulphuric acid :— 
3S02 + 2H,0 = 2S04H2 + S. 
Sulphurous acid is dibasic and forms two series of salts; the 
primary (KHS03) and secondary (K2S03). 
Sulphites. These are obtained by saturating solutions of 
bases with S02. When sulphurous acid is separated out of its 180 
INORGANIC CHEMISTRY. 
salts by stronger acids it decomposes into its anhydride and 
water:— 
Na2S03 + 2HCL = 2NaCl + S02 + H20. 
Hydrosulphurous Acid—H2S02. On adding zinc to the 
aqueous solution of sulphurous acid the metal dissolves without libera- 
tion of hydrogen. A yellow solution is obtained, which decolorizes 
indigo and litmus solutions energetically. Schiitzenberger has 
shown that this property is due to the hydrosulphurous acid 
contained in the solution, formed there by the action of the H set free 
by the zinc upon a second molecule ofS03H2 :— 
H2S03 + 2Zn = S03Zn -f- H2, and 
H2S03 + H2 = S02H2. 
The pure aqueous solution is obtained by the decomposition of its 
salts. Its solution has an orange yellow color, reduces powerfully, 
bleaches and soon decomposes with separation of sulphur. The 
bleaching action of this lowest oxygen compound of sulphur reminds 
us rf a similar behavior of the lower oxygen derivatives of chlorine 
and bromine. 
The salts are more stable than their acid. The sodium salt is 
obtained by the action of zinc filings upon a solution of primary sodium 
sulphite, and has the formula HNaS02 ; the second atom of hydrogen 
(as in the case of hypophosphorous acid (H2P02) cannot be replaced 
by metals. Hence the acid is monobasic, and its structure may be 
IV 
expressed by the formula HSO, OH. The solutions of the salts absorb 
oxygen from the air and are converted into sulphites:— 
HNaS02 + 0 = HNaS03. 
Two peculiar oxides of sulphur, which, however, do not afford any 
corresponding acids and salts, but resemble the peroxides more, are 
sulphur sesquioxide and sulphur heptoxide. 
Sulphur Sesquioxide—S203. Obtained by solution of flowers 
of sulphur in anhydrous sulphuric anhydride ; separates out in blue 
drops, which solidify to a malachite-like mass. It decomposes grad- 
ually, more rapidly on warming, into S02 and sulphur. It is violently 
broken up by water, with formation of sulphur, S02, S04H2 and 
polythionic acids. In concentrated sulphuric acid it dissolves with a 
blue color. 
Sulphur Heptoxide—S207—is produced by the action of the 
silent discharge of an electric stream of great tension upon a mixture 
of SO 2 and oxygen, and separates in oily drops, which solidify to a 
crystalline mass at 0°. Upon standing, especially upon warming, it 
gradually decomposes into S03 and oxygen:— 
S207 = 2S03 + 0. SULPHURIC ACID. 
181 
It fumes strongly in the air, and with water decomposes into sul- 
phuric acid and oxygen 
Its solution in concentrated sulphuric acid is tolerably stable. The 
same arises, too, in the electrolysis of sulphuric acid, and by addition of 
H202 to strongly cooled sulphuric acid. 
S207 -f 2H20 = 2S04H2 4- 0. 
Sulphur Trioxide—SO*—or sulphuric anhydride, is pro- 
duced, as previously described, by the union of S02 and oxy- 
gen, aided by platinum black; or when S02 and air are 
conducted over glowing oxide of iron (Wohler). It is most 
conveniently obtained by warming fuming (Nordhausen) 
sulphuric acid (p. 186); the escaping white fumes are con- 
densed in a chilled receiver. Sulphur trioxide exists in two 
different (polymeric) modifications. In the one fofm, obtained 
by cooling the vapors, there is produced a white, matted, 
silky mass which, after fusion, crystallizes in long, colorless 
prisms ; it fuses at 16° and boils at about 46°. The vapor 
density agrees with formula SO*. By keeping it below 25° it 
passes in another so-called solid modification, which does not 
fuse until above 50°, and passes into the liquid variety. 
According to later investigations of Weber both modifications are 
not pure anhydride, but hydrous. The pure anhydride he obtained 
from the ordinary asbestos like variety, by repeated, careful distilla- 
tion in a closed tube. It is a readily mobile liquid, of specific gravity 
1.940 at 16°, which solidifies to long, transparent, needles, like salt- 
petre. The crystals fuse at 14.8° and boil at 46.2°. By the addition of 
a small quantity of moisture the transparent crystals pass into the 
asbestos-like needles of the ordinary anhydride. 
Sulphuric oxide fumes strongly in the air, and attracts 
moisture with avidity. When thrown on water it dissolves 
with hissing, to form sulphuric acid (SOs -f- H20 = H2S04. 
When the vapors are led through heated tubes they are 
decomposed into S02 and oxygen. 
Sulphuric Acid—H2S04. This acid has long been known 
and is extensively applied in technology, etc. Besides the 
reaction first mentioned, it arises in the oxidation of sulphur 
by nitric acid. Formerly it was obtained by heating ferrous 
sulphate (FeS04); at present, however, it is almost exclusively 
manufactured in large quantities, after the so-called English 
lead chamber process. This method is based upon the con- 
version of S02 into S04H2. Sulphur or pyrite (FeS2) is 182 
INORGANIC CHEMISTRY. 
roasted in ovens, and the disengaged S02 immediately con- 
ducted, together with air, into a series of large leaden chambers 
in which it is frequently brought in contact with nitric acid 
and steam. By the combined action of these substances 
(sulphur dioxide, nitric acid, oxygen of the air and water) 
sulphuric acid is formed in the chambers and collected upon 
the floor of the same. 
The lead chamber process is very complicated, being influenced by 
the quantity of the reacting substances and the temperature, and is even 
not yet entirely explained. It is most simply represented as follows: 
in the presence of water, the nitric acid oxidizes the S02 to sulphuric 
acid, whereby it is reduced to nitrogen oxide or nitrogen dioxide:— 
3S02 + 2HN03 + 2H20 = 3H2S04 + 2N0. 
The oxygen of the air (which entered the chambers simultaneously 
with the S02) and the steam convert the NO again into nitric acid:— 
2N0 + 30 + H20 = 2HN03, 
and this converts a fresh portion of S02 into sulphuric acid. In this 
manner apparently one and the same quantity of nitric acid, by suffi- 
cient air access and water, changes an unlimited amount of S02 into sul- 
phuric acid ; the nitrogen oxide (and other oxides of nitrogen) acts, as 
it were, as oxygen bearer. In truth, in this process, in addition to the 
NO, small quantities of N20 and N are formed from the nitric acfd, 
and these not being oxidized by the air, escape, with excess of the 
latter, out of the chambers. For the continuation of the process, the 
regular addition of a definite amount of nitric acid is required. 
In practice, by means of the escaping nitrogen and excess of air, 
the active nitrogen oxides (N203 and N02) are carried along and 
withdrawn from the action. To avoid further loss of saltpetre thus 
occasioned, the escaping brown gases are aspirated by the so-called 
Gay-Lussac tower. The same is constructed from lead sheets, and filled 
with pieces of coke, over which concentrated sulphuric acid constantly 
trickles. The latter completely absorbs the nitrogen oxides N203 
and N02, with formation of nitrosylsulphuric acid (see p. 199). From 
the acid collected at the bottom of the tower—the so-called nitroso 
acids, the nitrogen oxides can be regained and made useful in the pro- 
duction of sulphuric acid in the chambers. This takes place, at present, 
in the so-called Glover tower, which, made of lead plates and fire-proof 
bricks, is 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 combus 
tion of the sulphur ovens stream against it. In this way the hot gases 
are cooled to the required temperature (70-80), from the acid chamber 
water evaporates, and, at the same time, the nitrogen oxides are set 
free (see p. 200) and carried into the lead chambers. Hence, the 
Glover tower.serves. not only for complete utilization of the nitrogen 
oxides, but also for the concentration of the chamber acid. 183 
SULPHURIC ACID. 
The chamber process may be illustrated by the following laboratory 
experiment: A large glass flask (Fig. 84) A replaces the lead cham- 
ber ; in its neck are introduced, by means of a cork, several glass 
tubes, serving to conduct the various gases. In a S02 is developed 
by heating a mixture of H2S04 and Hg or copper strips. The flask b 
contains some dilute nitric acid and copper turnings, from which is 
Fio. 84 
evolved (NO). In c water is boiled, to afford steam. Through d air 
enters, while the excess of gases escapes through e. By the meeting 
of (NO) with the air red fumes of nitrogen dioxide (N02) and nitrogen 
trioxide (N203) arise, and these in presence of water change the 
sulphur dioxide to sulphuric acid :— 
and 
S02 + N02 + H20 = H2S04 + NO 
The regenerated nitrogen oxide yields N02 with the oxygen of the air, 
and this converts another portion of S02 into sulphuric acid. In time 
aqueous sulphuric acid collects upon the bottom of the vessel. If at 
first only S02, NO and air enter without the steam, there is produced 
(by aid of the moisture of the air) the compound S02 { OH2 so' 
called nitrosulphonic acid) which covers the walls of the vessel with a 
white crystalline sublimate (comp. p. 200). These crystals, known as 
S02 + N203 + H20 = H2S04 + 2N0. 184 
INORGANIC CHEMISTRY. 
lead chamber crystals, also form in the technical manufacture of 
sulphuric acid, when an insufficient quantity of water is conducted into 
the chambers. Water decomposes them into sulphuric acid and 
nitrogen oxide. 
The acid collecting in the chambers (chamber acid) possesses, 
when the operation has been properly conducted, the specific 
gravity 1.5 (50° according to Beaume); it contains about 
60% H2S04 and 40% H20. For concentration the chamber 
acid is at first warmed in open pans until the specific gravity 
reaches 1.72 (603) BeaumS). The lead vats are strongly 
attacked by further evaporation, hence the acid is finally 
heated in glass vessels, or better, platinum retorts, until the 
residual liquid has acquired the specific gravity, 1.83 (66° 
Beaume). It is now entered up on trade under the name crude 
sulphuric acid (Acidum sulfuricum crudum). It still contains 
about 8% water and traces of lead and arsenic. 
By the distillation of the crude English acid an aqueous 
solution at first distills over (£ distillate) and at 330° almost 
pure H2S04 (Acidum sulfuricum purum or destillatum). 
This has the specific gravity 1.854 at 0° or 1.842 at 12°, and 
contains about 1.5% water. On cooling this to —35° white 
crystals separate, which by repeated recrystallization fuse at 
-f 10.5° ; this is the anhydrous acid, H2S04. Upon heating 
this white fumes of S03 escape at 403; the liquid begins to 
boil at 290°, and at 330° the acid, with 1.5% H20, again 
distills over. 
From these data it is obvious that sulphuric acid, even at gentle 
warmth, sustains a partial decomposition (dissociation) into S03 and 
H20, which, upon cooling, again unite to sulphuric acid. The disso- 
ciation is complete at a boiling temperature, as seen from the vapor 
density, which has been found to be 24.5. The normal vapor density, 
corresponding to the molecular formula H2S04 = 98, must be equal 
to 46 (9/); the empirically found formula, half as large, is explained 
by the decomposition of the molecule of H2S04 into the molecules 
S03 and H20. 
so4h2 = so3 + h2o 
1 volume. 1 volume. 1 volume. 
Concentrated sulphuric acid is a thick, oily liquid. On 
cooling, the ordinary English sulphuric acid (containing 8% 
water), to 0°, large six-sided prisms of the hydrate S04H2 + 
II20 separate; these fuse at -f 8.5°, and give up water at 
205°. The second hydrate, S04H.2 + 2H30, corresponding to 
the maximum contraction, has the specific gravity 1.62, and 
yields water at 195°. The concentrated acid possesses an ex- SULPHURIC ACID. 
185 
tremely great affinity for water, and absorbs aqueous vapor 
energetically, hence applied in the drying of gases qnd dessica- 
tors. It unites with water with great evolution of heat, and, for 
this reason, it is practically recommended, in mixing, to pour 
the acid in a thin stream into the water, and not the reverse, 
as otherwise explosive phenomena occur. In mixing sul- 
phuric acid with water, a contraction of the mixture takes 
place, whose maximum corresponds to the hydrate S04H2 -4- 
2H20. 
The existence of the hydrate of sulphuric acid, like that of periodic 
acid, finds explanation in the supposition of hydroxyl groups:— 
S04TI2 + 2 H20 =S(OH)6 Hexahydroxylsulphuric acid 
S041I2 + H20 — SfOH), Tetra “ “ “ 
S04H2 =S02(0H)2 Normal sulphuric acid. 
The tetra- as well as the hexahydroxylsulphuric acid form, by the 
action of bases, only salts of the normal dibasic acid ; salts in which 
more H atoms are replaced by metals, as with periodic acid, are not 
known. 
The affinity of sulphuric acid for water is so great that the 
former withdraws the hydrogen and oxygen from many sub- 
stances, with the production of water. In addition to carbon, 
many organic compounds contain hydrogen and oxygen in 
the proportion in which these elements yield water. This ex- 
plains the charring action of H2S04 upon wood, sugar and 
paper. When sulphuric acid acts upon alcohol (C2H60), 
ethylene C2H4 (p. 144), results. 
By conducting H2S04 over red hot porous bodies it is de- 
composed into sulphur dioxide, water and oxygen:— 
Upon this is based a method for manufacturing oxygen tech- 
nically ; the sulphur dioxide is absorbed by water and again 
changed to H2S04. Heated with S, P, C, and some metals 
(Hg, Cu), the acid is reduced to dioxide (see above). Nearly 
all the metals are dissolved by it, forming salts; only lead, 
platinum, and a few others are scarcely at all attacked. 
It is a very strong acid, and on warming separates most 
other acids from their salts ; upon this depends its application 
in the manufacture of hydrochloric and nitric acids. The 
barium salt (BaS04) is characterized by its insolubility in 
water and acids; therefore sulphuric acid added to solutions 
of barium compounds produces a white pulverulent precipi- 
tate, which can serve to detect small quantities of the acid. 
h2so4 = so2 + h2o + o. 186 
INORGANIC CHEMISTRY. 
Pyrosulphuric or Disulphuric Acid—H2S2 07. On with- 
drawing one molecule of water from two of the acid there 
results the compound S2 07H2, whose formation and structure 
may be represented by the following formula :— 
SO /OH so /oh 
b°2\OH tt _ __ 2\OH 
q0 /OH /OH 
bU2\OH bU2\OH 
As this contains two hydroxyl groups it is a dibasic acid ; 
besides, however, it reveals, from its formation, an anhydride 
character. Later we will observe that almost all polybasic 
acids as phosphoric acid PO (OH)3, silicic acid SiO (OH)2, 
chromic acid Cr02 (OH)2 are capable, by condensation of 
several molecules, with elimination of water, of forming like 
derivatives which bear the name Poly- or Pyro-acids. 
The disulphuric acid is contained in the so-called fuming 
or Nordliausen sulphuric acid (Acidum sulfuricum fumans), 
which is obtained by heating dehydrated ferrous sulphate— 
green vitriol (Fe£04). It is a thick, oily, strongly fuming 
liquid, of specific gravity, 1.85-1.9. On cooling this, large, 
colorless crystals of H2S207 separate ; these fuse at 35°. On 
warming it breaks up into sulphuric acid and sulphur trioxide, 
which volatilizes:— 
s2o7h2 = so4ii2 + so3. 
Conversely, the disulphuric acid dissolves S03 and can 
become sulphuric acid. 
The production of fuming sulphuric acid depends on this : 
it may be regarded as a solution of S03 (or S207H2) in excess 
of sulphuric acid. 
Technically, fuming sulphuric acid is obtained from pyrites (FeS2)— 
(at present only in Bohemia). By weathering of the pyrites in the air fer- 
rous sulphate and ferric oxide arise. The first can be dissolved out with 
water. The solution is evaporated and roasted in a reverberatory 
furnace, whereby the ferrous salt is changed to ferric salts. The latter 
are then distilled from earthen retorts, when sulphuric acid and the 
trioxide pass over and are collected in the receivers:— 
Fe2(S04), = Fe203 + 3SOs 
Fe2 { (OHh = + s°a + S04H2. SULPHURIC ACID. 
187 
The residue, consisting of red ferric oxide, finds application as 
colcothar (caput mortuum) in polishing and as a paint. 
Recently, the method of Winkler has been employed to obtain solid 
fuming sulphuric acid H2S207. The mixture of S02 + O, obtained 
by the heating of English sulphuric acid (p. 184) and absorption of 
steam produced at same time by sulphuric acid in a coke tower, is 
conducted over glowing platinized asbestos (p. 178) and the resulting 
SO3 taken up by concentrated sulphuric acid. 
With water disulphuric acid yields sulphuric acid. Its salts 
may be obtained by heating the primary salts of the latter 
acid:— 
2S0.tHK = h2o + k2s2o7. 
Primary pot. 
sulphate 
pot. disulphate. 
By further warming, these decompose into S03, and sul- 
phates: S207K2 = S04K2 + S03; hence these reactions may 
serve for the formation of S03. 
Sulphuric Acid Chlor-anhydride. Under the name of 
halogen anhydrides we understand the derivatives resulting from the 
replacement of OH in hydroxides by halogens. Conversely, the chlor- 
anhydrides, by the action of water, pass into the corresponding acids :— 
OH 
S02 { £} + 2H20 = S02+ 2 HC1. 
XOH 
The ordinary method for the preparation of the chloranhydrides 
consists in permitting PC15 to act on the acids. Sulphuric acid has two 
hydroxyl groups ; therefore it can furnish two chloranhydrides. The 
first, S02C Qpj — sulphuryl hydroxy-chloride or chlorsulphonic 
acid—results when 1 molecule PC15 acts upon 1 molecule H2S04 :— 
Cl 
s°2 { OH + PC15 = SO. + POCI3 + HC1. 
OH 
The resulting P0C13 acts further upon 2 molecules S04H2, with 
formation of phosphorous and chlorsulphonic acids. 
It is formed, too, by the direct union of S03 with HC1. It is most 
practically prepared by conducting chlorine gas through S04H2 (16 
parts) and allowing PC13 (7 parts) to gradually drop in. Or HC1 gas 
is led into solid fuming sulphuric acid (S207H2), so long as absorption 
occurs, and then distilled (Otto). 188 
INORGANIC CHEMISTRY. 
Chlorsulphuric acid is a colorless, strongly fuming liquid, of specific 
gravity 1.776 at 18°, which boils at 152°. The salt S02 | results 
from the union of S03 with HC1. 
The second chlor,anhydride S02C12, or sulphuryl-chloride* forms 
when PC15 acts upon S03 ; by heating S03HC1 to 180°: 2S03HC1 == 
S02C12 + S04H2; and also by the direct union of S02 with Cl2 in 
sunlight. A colorless, suffocating, strongly fuming liquid of specific 
gravity, 1.66 results. It boils at 77°. Water decomposes both these 
anhydrides into sulphuric aud hydrochloric acids. 
Thionyl chloride—SOCl2—(p. 179), maybe regarded as a chloranhy- 
dride of sulphurous acid. The chloranhydride (S205C12) of disulphuric 
acid is known. It boils at 146°. 
Amid. Derivatives of Sulphuric Acid. When NH, acts on 
S03, the compound S032NH3 arises, which is to be regarded as am- 
monium sulphaminate —S02^q^2jJj It is a white powder, which may- 
be crystallized from water : the solution is not precipitated by barium 
/ATT 
salts. When ammonia acts upon S0 q the ammonium salt of 
/ar\ OFT 
disulphimid-acid gQ2’opp forms; this is tribasic, as all three 
H atoms can be replaced by metals. 
POLYTHIONIC ACIDS. 
By this name (from deiov, sulphur,) is understood the com- 
plex acids of sulphur, containing 2 or more atoms of the 
latter. The following are known :— 
S203H2 — Thiosulphuric acid. 
S206H2 — Dithionic Acid. 
S306H, — Tri “ 
S406H'2 — Tetra “ 
S506H2 — Penta“ “ 
The general chemical character of these acids is represented 
most simply and distinctly in the following structural formu- 
las. We suppose that in them one or two monatomic groups, 
S03H or —S02 — OH, are contained, in which one affinity 
of sulphur is unsaturated. This is known as the sulpho 
group; it is also present in organic sulpho-acids, and corres- 
ponds to the acid forming carbon group, COOH, known as 
* The group combined with 20H in H2S04 is known as sulphuryl. POLYTHIONIC ACIDS. 
189 
carboxyl. From this group, (written in another form,) are 
derived the above observed acids:— 
h,so2,oh - H0,S02,0H °\So!-OH 
The following structural formulas express the polythionic 
acids:— 
Sulphurous acid. 
Sulphuric acid. 
Disulpburic acid. 
/SO3H 
\SO3H 
Dithionic acid. 
0/SO3H 
u\so3h 
Trith ionic acid. 
hs,so3h 
Thiosulphuric acid. 
q /SO3H 
&2\S03H 
Tetrathiouic acid. 
q /S03H 
&3\S03H 
Pentathionic acid. 
The last three acids may be viewed as derivatives of the 
hydrogen sulphides, SH2, S2H2, and S3H2, in which both H 
atoms are replaced by two monatomic sulpho groups. In 
thiosulphuric acid, only lH-atom is replaced by sulpho; the 
dithionic acid, on the other hand, results by the direct union 
of two sulpho groups, with their free affinities. 
/ OTT 
Thiosulphuric acid, H2S203 = SO,^gjj ’ generally known 
as hyposulphurous acid, can be considered sulphuric acid in 
which the oxygen of an hydroxyl group is replaced by sul- 
phur. It is not known in a free condition, since as soon as it 
is freed from its salts by stronger acids, it decomposes instan- 
taneously into S02, S and H20:— 
S203Na2 + 2HC1 = 2NaCl + S02 + S + H20. 
Its salts, called hyposulphites, are of practical importance 
(compare sodium hyposulphite). They form by the direct 
addition of sulphur to sulphites :— 
similar to the formation of sulphates by the addition of O to 
the sulphites. 
Particularly interesting is the formation of thiosulphuric 
acid by the action of iodine upon a mixture of sodium 
sulphite and sodium sulphide :— 
Na2S03 + S = Na2S203; 
f NaS02.0Na S02.0Na 
+ la = I +2NaI. 
I NaSNa SNa 
Sodium hyposulphite. 190 
INORGANIC CHEMISTRY. 
Conversely sodium hyposulphite is split up by sodium amal- 
gam into H03Na2 and Na2S. 
Dithionic Acid—II2S206—is only known in aqueous solution. 
By concentration in vacuo and by heating it decomposes into sulphuric 
acid and sulphur dioxide. Its manganese salt results from the action 
of sulphur dioxide upon Mn02 suspended in water. 
Mn02 + 2S02 = MnS206. 
Barium hydrate converts this into the barium salt, and from this aided 
by sulphuric acid the free dithionic acid is obtained. 
Trithionic Acid—H2S306 —is not known in free condition. 
Its salts are produced when an aqueous solution of primary potassium 
sulphate is warmed with flowers of sulphur :— 
Separated from its salts by other acids it decomposes into H2S04, 
SO 2 and S. Of great interest is its production by the action of 
iodine upon a mixture of sodium sulphite and hyposulphite :— 
6HKS03 + 2S = 2K2S306 + K2S203 + 3H20. 
S02.0Na 
NaS02.0Na , T 4-2NaI 
NaS,S02.0Na + ~ • + 
S02.0Na 
Tetrathionic Acid—H2S406. When iodine acts upon 
solutions of hyposulphite its salts are produced. 
KS.SO3K s.so3k 
+!,= I +2KI 
KS.SO3K S.SO3K 
Potassium tetrathionate. 
The free acid decomposes, upon concentration, into H2S04, S02 
and 2S. 
Pentathionic Acid—H2S506—results when H2S is conducted 
into an aqueous solution of sulphur dioxide :— 
5S02 ~j~ 5H2S — H2S506 -f- 4H20 -f- 5S ; 
further, by the action of S2C12 upon barium hyposulphite : 
S.S03 S.SO3 
/ \ / \ 
S2C12 + Ba Ba = S Ba + BaCl2 + S. 
\ X \ / 
S.S03 S.S03 
According to Spring’s late investigations, pentathionic acid does 
not exist. 
The polythionic acids are distinguished from sulphuric acid 
by the solubility of their barium salts. SELENIC AND TELLURIC ACID. 
191 
2. Oxygen Derivatives of Selenium and Tellurium. 
Se02 
Selenium dioxide. 
Se03H2 
Selenious acid. 
Se03 
Selenium trioxide. 
Se04H2 
Seleuic acid. 
Selenium Dioxide—Se02—or selenious anhydride, is pro- 
duced when selenium burns in the air or in oxygen, forming 
long white needles, which sublime at about 3203 without fus- 
ing. In water it readily dissolves to selenious acid, H2Se03. 
The latter is also formed by dissolving the metal in concen- 
trated nitric acid. When the solution is evaporated it crys- 
tallizes in large, colorless prisms, which decompose, on heating, 
into the anhydride and water. Sulphurous oxide reduces 
selenious acid, with separation of free selenium :— 
H2Se03 + 2S02 + H20 = 2H2S04 + Se. 
Selenic Acid—H2Se04—is obtained by conducting chlorine 
gas into an aqueous solution of selenious acid. 
H2Se03 + H20 + Cl a = H2Se04 + 2HC1 
The solution may be concentrated until it attains a specific gravity 
of 2.96, when it becomes an oily liquid, similar to sulphuric acid, and 
containing 95 per cent. H2Se04. If the solution be heated above 
280°, the acid breaks up into Se02,0 and H20. The anhydride of 
the acid is unknown. 
The salts of selenic acid are known as selenates, those of selenious 
acid as selenites. 
The derivatives of tellurium are analogous to those of selenium. 
The dioxide—Te02—results from burning tellurium, and forms a white 
crystalline mass, fusing at a red heat and subliming. In water it is 
almost insoluble. 
Tellurous Acid—H,TeO 3—is produced when the metal is dis- 
solved in concentrated nitric acid. Water will precipitate it from such 
a solution as a white amorphous powder. On warming, it readily 
breaks up into Te02 and water. 
Telluric Acid—H2Te04. Upon fusing tellurium or its dioxide 
with saltpetre, potassium tellurate is formed, from which, by means of 
sulphuric acid, telluric acid is obtained. From aqueous solutions it 
crystallizes in large, colorless prisms, with 2 molecules of H20 
(H2Te04 -f- 20), which are expelled at 100° C., and the acid 
becomes a white powder. Water dissolves the latter with difficulty 
and exhibits only a slightly acid reaction. Carefully heated, the acid 
breaks up into water and the trioxide Te03. which is a yellow 
mass insoluble in H20, and upon further application of heat decom- 
poses into Te02 and oxygen. 192 
INORGANIC CHEMISTRY. 
The affinity of the elements of the oxygen group for the 
halogens appears, the reverse of that of the halogens with 
hydrogen, to gradually increase with rise of atomic weight 
from oxygen to tellurium successively; OCl2 is very unstable 
and is formed with heat absorption, SC12 and SC14, only exist 
at lower temperatures, while SeCl4, TeCl2 and TeCl4 even 
exist as gases. On the contrary, the thermal relations in the 
formation of the oxygen compounds indicate that in them the 
affinity of selenium is the least, as is the case with bromine in 
the halogen group (p. 176). This follows from the heat dis- 
engagement in the formation of the acids R03H2, and also the 
higher acids, R04H2 (from the elements and water) :— 
(S,02,Aq) = 78770 
(S,03,Aq) = 142400 
(Se,02,Aq) = 56790 
(Se,03, Aq) = 77240 
(Te,02,H20) = 81990 
(Te,03,Aq) = 107040 
The same is seen in the formation of the anhydrous 
dioxides:— 
Of course in all these compounds the affinity of selenium 
to oxygen is the least, and in it the reduction of selenious 
by sulphurous acid, and the slight stability of selenic acid 
find explanation. 
(S,02—gas) = 71070 (Se,02—solid) = 57700. 
3. Oxygen Derivatives of the Elements of the Nitrogen Group. 
The halogens combine with one atom of hydrogen and also 
alford oxygen acids containing one atom of the former. The 
elements of the sulphur group contain two atoms of H in the 
hydrogen derivatives and oxygen acids. Corresponding to 
this the elements of the N group combining with 3 atoms of 
H, form acids which, too, contain 3 atoms. 
HCl 
HC104 
Perchloric acid. 
HC103 
Chloric acid. 
h2s 
h2so4 
Sulphuric acid. 
h2so3 
Sulphurous acid. 
ph3 
po4h3 
Phosphoric acid. 
po3h3 
Phosphorous acid. 
The acids containing three atoms of H, designated normal 
or Ortho-Acids (as H3P04, H3As04, H3As03) can yield mono- 
basic acids by the removal of one molecule of water. Such 
derivatives, having one atom of H, are called meta-acids :— 
h3po4 
Orthophosphoric acid. 
H3As03 
Orthoarsenious acid. 
hpo3 
Metaphosphoric acid. 
HAs02 
Metaarseuious acid. NITRIC ACID. 
193 
These meta-acids of phosphorus and arsenic are less stable 
than the ortho-acids and pass into the latter by the absorption 
of water. The ortho-acids of N, on the other hand, are less 
stable and only exist in some salts. The ordinary acids and 
salts of N belong to the meta-series and contain one atom of H 
(or metal): 
h3no4 
Orthonitric acid. 
H3NO3 
Orthouitrous acid. 
HNOs 
Ord. Nitric acid. 
hno2 
Ord. Nitrous acid. 
1. OXYGEN DERIVATIVES OF NITROGEN 
n203 
Nitrogen trioxide. 
hno2 
Nitrous acid. 
n2o 
Hyponitrous oxide. 
N205 
Nitrogen pentoxide. 
hno3 
Nitric acid. 
HNO 
Hyponitrous acid. 
Besides these compounds there exist further: Nitrogen 
tetroxide (N204), the mixed anhydride of nitrous and nitric 
acids, and two oxides, nitrogen dioxide (NOa) and nitrogen 
oxide (NO), which do not yield acids. 
The following formulas express the structure of these com- 
pounds :— 
nen 
Nitrogen. 
Ill III 
N = N 
\/ 
O 
Ill III 
0 = N — 0 — NO 
Nitrogen trioxide. 
o2n —0 —no2 
Nitrogen pentoxide. 
The salts of nitric acid are called nitrates; those of nitrous, 
nitrites. 
Hyponitrous 
oxide. 
Ill 
ON — OH 
Nitrous acid. 
02N — OH 
Nitric acid. 
Ill 
ON —H 
Hyponitrous. 
acid 
Ill V 
ON — 0 — N02 
Nitrous-Nitric 
anhydride. 
Nitric Acid—HN03—This acid occurs in nature only in 
the form of salts,—potassium, sodium and calcium saltpetre 
(compare these)—which have resulted from the decay of nitro- 
genous organic substances in the presence of strong bases (the 
alkalies). Sometimes present in the air as ammonium salt. 
The free acid is formed in very slight quantity by continu- 
ously conducting the electric sparks through moist air. 
To prepare nitric acid heat potassium or sodium nitrate with 
sulphuric acid, by which operation nitric acid will distill over 
and sodium sulphate remain :— 
2NaN03 + H2S04 = Na2S04 + 2HN03 and 
NaN03 + H2S04 = HNaS04 + HN03. 194 
INORGANIC CHEMISTRY. 
The process may be conducted in the distillation apparatus 
figured (page 44). The quantity, by weight, of sodium nitrate 
and sulphuric acid corresponding to the second equation 
must be employed, since, with less acid to complete the reac- 
tion, a higher temperature is requisite, and, in consequence, 
the formed nitric acid will be partially decomposed. 
Pure anhydrous nitric acid is a colorless liquid of specific 
gravity 1.54 at 0°, fuming in the air and at —40°, solidifying 
to a crystalline mass. At medium temperatures it undergoes 
a partial decomposition (like H2S04) into water, oxygen and 
nitrogen dioxide N02, which dissolves in the acid, with a 
yellowish-brown color ; the colorless acid changes upon stand- 
ing, becoming yellow rapidly in sunlight. At 86° the acid 
commences boiling, with partial decomposition; the first 
portions are colored yellow by the dissolved nitrogen dioxide, 
while subsequently, some aqueous acid distills over. Nitric acid 
is completely decomposed into nitrogen dioxide, oxygen and 
water, when its vapors are conducted through red hot tubes:— 
2HN03 = 2N02 + H20 -f 0. 
The acid mixes in all proportions with water. Upon dis- 
tilling the dilute aqueous solution, at first only pure water 
passes over; the boiling temperature gradually rises, and, at 
121°, a solution goes over, containing 68%, HN03, and indi- 
cating a specific gravity of 1.414 at 15°. This is the ordinary 
concentrated nitric acid of trade. On distilling this, together 
with 5 parts sulphuric acid, almost anhydrous acid is obtained, 
which is freed of contained N02, by conducting a stream of air 
through it. 
Generally, the anhydrides of acids distill at temperatures lower than 
the acids themselves (S03 is more volatile than H2S04). The higher 
boiling point of the aqueous acid, in relation to the anhydrous, is 
probably explained by the fact that, in the solution, the hydrate, 
HN03 -f H20, i. e., the normal nitric acid (N04H3) = NO(OH)3, 
compare page 193, is present. The solution boiling at 121°, however, 
contains more water than corresponds to this hydrate (just as distilled 
sulphuric acid contains water), so that it can be regarded as a mixture 
of the trihydrate (NO(OH)3) and pentahydrate (N(OH)5). 
Nitric acid is a very powerful acid, oxidizing or dissolving 
almost all metals (gold and platinum excepted). Nearly all 
the metalloids, like sulphur, phosphorus and carbon, are con- 
verted by it into their corresponding acids. Especially does 
the acid act as a strong oxidizing agent, destroys organic color- 
ing substances and readily decolorizes a solution of indigo. In NITRIC ACID. 
195 
so doing, the nitric acid itself is deoxidized to the lower oxida- 
tion products of nitrogen (NO and N02). Some substances 
even reduce the acid to ammonia. Thus, for example, in 
bringing zinc into nitric acid the metal is dissolved without 
the liberation of hydrogen. The latter at once, in statu nas- 
cendi, acts upon the excess of acid and reduces it to ammonia, 
which forms an ammonium salt with the acid ; hence, in solu- 
tion, we have ammonium nitrate in addition to the zinc 
nitrate — 
2PINO3 + Zn = Zn(N03)2 + H, and 
2HN03 + 4H2 = N03NH4 -f 3H20. 
If the aqueous nitric acid be more dilute (containing more 
than 10 °Jo N03H), the same is reduced by zinc and other 
metals, not to ammonia, but to the nitrogen oxides N20, NO, 
N203 and N204. 
The reduction of nitric acid to ammonia by nascent hydrogen 
occurs more easily in alkaline solution. On treating such a 
solution of nitrates with zinc or aluminum filings, all the N 
of the nitric acid is converted into ammonia :— 
HN03 + 4Ha = NH3 + 3H20. 
By the action of tin upon nitric acid there is formed 
hydroxylamin, together with ammonia (p. 121). 
Nitric Acid—N020H—as also its hydrates NO(OH)3 and 
N(OH)4 (p. 194), form almost exclusively only salts with 1 
aeq. of the metals of the form N03Me; these are called nitrates, 
and are all soluble in water. 
Red Fuming Nitric Acid (Acidum Nitricum Fumans) is the 
name given a nitric acid containing much nitrogen dioxide in 
solution. It is obtained by the distillation of 2 molecules 
HN03 with 1 molecule sulphuric acid (see above), or, simpler, 
by the distillation of commercial nitric acid with much 
sulphuric acid. Generally it has the specific gravity 1.5-1.52 
and solidifies at — 40° to a crystalline mass. It possesses 
greater oxidizing power than the colorless nitric acid. 
A mixture of 1 volume nitric acid and 3 volumes concen- 
trated 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 
the mixture depends upon the presence of free chlorine and 
both chlorine derivatives (N02C1 and NOC1), which may be 
considered the chloranhydrides of nitric and nitrous acids. 196 
INORGANIC CHEMISTRY. 
Nitroxyl Chloride—N02C1—the chloranhydride of nitric acid, 
results from the union of N02 with chlorine, and also, accordingto the 
ordinary method of forming chloranhydrides (see p. 187), by the action 
of PCI 5 or POCl3 upon nitric acid, or better, its silver salt:— 
3N020Ag + POCl3 = PO(OAg)3 + 3N02C1. 
Silver nitrate. Silver phosphate. 
It is a yellowish liquid, boiling at -(- 5°. With water it breaks up 
into nitric and hydrochloric acids. 
Nitrosyl chloride—NOC1—is produced when hydrochloric acid 
acts upon nitric acid, and also by the union of NO (2 vols.) with 
chlorine (1 volume). It is a reddish-yellow gas which below 0° con- 
denses to a liquid. With water it forms nitrous and hydrochloric 
acids:— 
It may, therefore, be looked upon as the chloranhydride of nitrous 
acid—NO.OH. 
N0C1 + H20 = HN02 + HC1. 
Nitrogen Pentoxide—N205—nitric anhydride, arises when 
phosphoric anhydride acts on nitric acid :— 
further by conducting nitroxyl chloride over silver nitrate:— 
Ag.O.NO, + N02C1 = } 0 + AgCl. 
2HN03 + P205 = N205 + 2HP08 
It forms colorless, rhombic prisms, fusing at 30° and boiling 
with partial decomposition at 47°. It is very unstable, 
decomposing readily into N204 and O and sometimes explod- 
ing spontaneously. With water it yields HN03 and evolves 
much heat by the union:— 
*J°2}0 + H20 = 2N020H. 
Nitrogen Trioxide, N203, nitrous anhydride, is formed by 
the direct union of nitrogen oxide (4 vols.) with oxygen (1 
vol.) about—18° :— 
4 NO+ 02=2 N203. 
4 vols. 1 vol. 
And by mixing liquid nitrogen tetroxide, N204, with a little 
water:— 
2 gg2 } 0 + H20 = gg } 0 + 2 N02,0H; 
further by introduction of nitrogen oxide into liquid nitrogen 
tetroxide:— 
and of nitrogen oxide into nitric acid :— 
N204 + 2 NO = 2 N203; 
2 N03H 4- 4 N0= 3 N203 + H20. NITROGEN TETROXIDE. 
197 
It is a dark blue liquid, boiling at 0° with partial decompo- 
sition into NO and N02; both gases combine again by cool- 
ing to N203. The trioxide mixes with a little cold water, 
forming, probably, nitrous acid (HN02); by more water, and 
when warm, it is decomposed to nitric acid and nitrogen oxide 
gas:— 
3 N02H = HN03 + 2 NO 4- H20 
Nitrous Acid, HN02, is not known in a free state. Its salts 
are obtained by glowing the nitrates:— 
kno3 = kno2 + 0. 
The withdrawal of oxygen is rendered easier if oxidizable 
metals, e. g., lead, be added to the fusion. * 
On adding sulphuric acid to the nitrites, brown vapors are 
disengaged ; these consist of N02 and NO. We may suppose 
that the nitrous acid, at first liberated, is broken up into water, 
and the trioxide, which, as we have above seen, decomposes 
into N02 and NO very readily. Similar reddish-brown vapors 
are obtained if nitric acid of specific gravity 1.3-1.35 be 
permitted to act upon starch or arsenious oxide (As203). 
On cooling, these vapors condense to a liquid, which at 
medium temperatures is green, and probably consists of N204 
and N203. When the green Hquid is warmed, vapors escape, 
which, on cooling, condense to a blue liquid, consisting princi- 
pal lv of N203. 
The nitrous acid separated out in the solution and its 
decomposition products—N02 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 acidi- 
fied red solution of potassium permanganate is decolorized 
by the addition of nitrites. 
In very dilute solution the action proceeds according to the 
following equation :— 
5N02H + 2Mn04K + 3S04H2 = 5N03H + S04K2 + 
2S04Mn + 3H20. 
Nitrogen Tetroxide—N204 or nitrogen dioxide N02 (formerly- 
called hyponitric acid) really constitutes two compounds. The 
former only exists at low temperatures ; on warming, it suffers 
a gradual decomposition into the simpler molecules N02 which 
upon cooling recombine to N204. We meet here the interest- 198 
INORGANIC CHEMISTRY. 
ing case of dissociation, ocurring even at a medium tempera- 
ture. N204 is colorless, while N02 is colored red-brown: it 
appears, therefore, with increasing temperature the color, 
corresponding to the increasing dissociation of the complex 
molecules N204, gradually becomes darker. 
The theoretical vapor density of N204 (molecular weight = 92) 
equals 46, while that of N02 (46) = 23. At the boiling temperature 
(26°) of the liquid compound the experimental vapor density has been 
found = 38 ; from this, by calculation, we find that at this temperature 
20 per cent, of the molecules N204 decompose into the molecules 
N02. Hence we conclude that the dissociation of the compound 
N204 commences already in the liquid state, which is confirmed by the 
yellow coloration appearing at 0°. Sulphuric acid, as we saw (p. 184), 
exhibit* a similar dissociation in liquid condition. With rising tem- 
perature the density of the vapor steadily diminishes, becomes con- 
stant finally at 150° and equals 23. Then all the molecules (N204) 
are decomposed into the simpler molecules N02; at the same 
time, at this temperature the dark coloration of the vapors attains its 
maximum. 
Nitrogen tetroxide is formed by the union of two volumes of 
nitrogen oxide with one volume of oxygen :— 
2N02 -|- 02 —N204. 
2 vols. 4 vols. 
We can get it more conveniently by heating dry lead nitrate, 
which decomposes according to the following equation 
The escaping vapors condense in the cooled receiver to 
liquid N204. 
The varying molecular composition of nitrogen tetroxide at 
lower and higher temperatures manifests itself also in its 
chemical reaction. We saw that by the action of a little cold 
water, the tetroxide is decomposed into nitrogen trioxide and 
nitric acid (p. 197). With excess of cold water, and also with 
an aqueous solution of alkalies, it forms nitric and nitrous 
acids—that is, their salts :— 
(N0,)2Pb = Pb + 0 + 2N0v 
N02x 
>0 + H20 = N020H -f NO, OH. 
NO x 
Both reactions plainly indicate that the liquid tetroxide 
represents the mixed oxide of nitric and nitrous acid; 
similarly, the compound C1204 constitutes the mixed oxide 
of chloric and chlorous acid (p. 171). NITROSYLSULPHURIC ACID. 
199 
Warm w’ater converts the tetroxide into dioxide N02, which 
in turn yields nitric acid and nitrogen oxide — 
2N02 + H20 = 2HN03 + NO. 
The tetroxide and dioxide have strong oxidizing properties ; 
many substances burn in their vapors ; iodine is set free from 
the soluble metallic iodides by them. 
.0, NO. 
IMitrosylsulphuric Acid, SOaNH = S02 This com- 
xOH. 
pound, also termed nitrosulphonic acid, which, as intermediate 
product in the technical manufacture of sulphuric (see p. 182), 
is important, and also for the analytical determination of the 
nitrogen oxides, is produced by conducting nitrogen trioxide 
and tetroxide into concentrated sulphuric acid:— 
2S°!\oii + N*°* = 2SO‘\oh!° + H*°- 
SO»\OH + NO.) 0 = SOKoH ° + N0*’0H- 
Nitrogen monoxide—NO—is not absorbed by pure sulphuric 
acid, but will be if the same contain nitric acid:— 
O.NO 
/ 
3S04H2 -f N03H + 2NO — 3S02 + 2H20. 
\ 
OH 
Further, the nitrosylsulphuric acid results from the action 
of sulphurous oxide, nitrogen tetroxide and little water:— 
O.NO 
/ 
2S02 + N204 + 0 + H20 = 2S02 
\ 
OH. 
Most easily obtained by conducting sulphur dioxide into 
strongly cooled anhydrous nitric acid :— 
O.NO 
/ 
S02+NO3H = S02 ; 
\ 
OH. 
there results a thick magma, which may be dried upon 
porous earthen plates under the desiccator. 200 
INORGANIC CHEMISTRY. 
Nitrosylsulphuric acid forms a leafy or granular crystalline, color- 
less mass (chamber acid crystals p. 184), which fuse about 73° and 
decompose into the trioxide, sulphuric acid and nitrogen trioxide. 
In moist air it deliquesces, and with water breaks up into sulphuric and 
nitrous acids :— 
O.NO OH 
/ / 
S02 + H2 0 = S02 + NO.OH. 
\ \ 
OH OH 
and the latter further partly into nitric acid and nitrogen oxide. 
In concentrated sulphuric acid nitrosylsulphuric acid dissolves with- 
out any change ; the solution called nitroso acid, produced also in the 
sulphuric acid manufacture in the Gray-Lussac tower, is very stable and 
may be distilled without decomposition. When diluted with water it 
remains at first unaltered, until the specific gravity of the solution 
reaches 1.55-1.50 (51-48° B); then escape, especially on warming, all 
the nitrogen oxides. When the nitroso acid is poured in much water, 
the nitrososulphuric acid breaks up into (like the pure acid, see above) 
sulphuric and nitrous acids; the latter partially further into N03H and 
2NO. Therefore in titrating nitrous acid with permanganate of 
potassium (Mn04K. See p. 197) we only get the results corresponding 
to nitrososulphuric acid, if the latter be poured into the permanga- 
nate (Lunge.) 
All the nitrogen oxides and acids are separated as nitrogen oxide 
(NO) by shaking the nitroso-acid with mercury—a procedure serving 
equally well for estimating the amount of nitroso acid by means of the 
nitrometer. All nitrogen oxides are expelled from the nitroao acid by 
dilution with water and application of heat (see above.) More easily 
and completely (even upon concentration to 58°B = 1.679 specific 
gravity) occurs the action of sulphur dioxide : — 
O.NO OH 
/ / 
2SOz + 2HzO + S02 =3S02 + 2NO. 
\ \ 
OH OH 
Upon this depends the denitrating action of the Glover tower—see p. 
182. 
The anhydride of nitrosulphonic—S2N209 = 0 { 2 0 NO 's 
produced in the heating of the latter (together with S04H2 and 
N203—see p. 199). In pure condition it is obtained by saturation of 
sulphur trioxide with nitric oxide :— 
3S0) + 2N0 = 0 {|8*:8:no + 80»- NITRIC OXIDE. 
201 
It is a crystalline, colorless mass, fusing at 217° and boiling without 
decomposition about 3b0°. Much water decomposes it, the same as 
nitrosylsulphuric acid. 
The chloranhjdride of nitrosulphonic acid, SN04C1 = SO 
is formed by the union of sulphur trioxide with nitrosylchloride :— 
S03 + N0C1 = S02<0-N0 
White leaflets, decomposed by heat into its components, and with water 
breaks up into sulphuric, hydrochloric and nitrous acids. 
Nitric Oxide—NO. When different metals are dissolved 
in somewhat diluted nitric acid this oxide is formed, inasmuch 
as the hydrogen in statu nascendi at first liberated reduces 
another portion of the acid. It is most conveniently obtained 
by pouring dilute nitric acid (specific gravity 1.2) upon cop- 
per filings:— 
3Cu + 8HN03 = 3Cu(NOs)2 + 4H20 + 2N0 
The action begins in the cold. 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 N02. Therefore all the air must be 
expelled from the generating vessel by NO and the gas col- 
lected over water when the interior of the apparatus has 
become colorless. 
Nitric oxide is a colorless gas, of specific gravity 15 
(H = 1) or 1.039 (air = 1), which can be condensed by cold 
and pressure. It is slightly soluble in water, dissolving, how- 
ever, very readily in an aqueous solution of ferrous salts, and 
imparting to the liquid a reddish-brown color; heat expels it 
from the same. Nitric oxide is readily soluble in nitric acid. 
According to its concentration it is colored brown, yellow, 
green and blue, as nitrogen trioxide is formed :— 
2HN03 + 4N0 = 3N203 + H20. 
Potassium permanganate oxidizes it, like nitrous acid 
(p. 197), to nitric acid :— 
10NO+6MnO4K+9SO4H2=10NO3H+3SO3K2 + 6SO4Mn + 4H2O 
As nitric oxide contains 57% oxygen, it is capable of 
sustaining the combustion of some substances, but to bring 
about the previous separation of oxygen from the nitrogen re- 202 
INORGANIC CHEMISTRY. 
quires energetic reaction. Hence, phosphorus continues to 
burn in this gas, ivhile a sulphur flame, developing only a 
slight heat, is extinguished; glowing charcoal does the same, 
while an energetically burning splinter will continue to burn 
in it. On shaking a few drops of readily volatile carbon di- 
sulphide into a cylinder filled with NO, and bringing a flame 
to the mouth of the vessel, the carbon disulphide vapors will 
quietly burn in the gas, giving a bright luminous flame, emit- 
ting strong actinic rays; in this combustion, the C and S of 
the CS2 unite with the oxygen of the nitric oxide. 
On (determining the quantity of heat disengaged in the combustion 
of phosphorus, carbon or other substances in NO gas, it will be discov- 
ered that the same is greater (about 21600 calories), 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 necessary for the 
separation of NO into N and 0, than for the separation of the mole- 
cules of combined oxygen atoms—an additional proof that the mole- 
cules of free oxygen (as of other elements) consist of atoms. 
With oxygen, NO at once forms brown vapors of nitrogen 
dioxide :— 
2N0 + Q2 = 2N02 
2 vol. 1 vol. 2 vol. 
With less oxygen nitrogen trioxide is produced, (p. 196). 
Just as with oxygen so with chlorine, NO combines to nitro- 
sylchloride NOC1 (p. 196), and the compound, NOCl2, which 
is, as yet, little investigated ; with bromine, it unites to form 
NOBr3. At a red heat NO becomes N02 and N. With 
hydrogen and moderate heat, it forms water and nitrogen : 
NO -f- H2 = N -f- H20 ; a mixture of both gases burns with 
a white flame. On conducting NO and II together over 
platinum sponge, water and ammonia are produced:— 
2N0 + 5H2 = 2NH3 + 2H20. 
The volumetric analysis of nitric oxide gas may be easily 
executed as follows : Fill 
a bent glass tube (Fig. 85) 
over mercury with NO 
gas; introduce in the same 
a piece of sodium and heat 
the latter with a lamp. 
The sodium combines with 
the oxygen, and free nitro- 
Fig. 85. NITROUS OXIDE, 
203 
gen separates ; the volume of the latter always equals half 
the volume of the employed nitric oxide gas, which would 
follow from the formula NO :— 
The molecular formula of the oxide is NO = 30, as its vapor density 
is 15 (H = 1). NO, N02 and chlorine dioxide, C102 (p. 171), pre- 
sent an apparent anomaly as regards the common laws regulating the 
valence of the elements. Ordinarily, the quantivalence changes from 
an odd number to an odd, and from an even to an even number 
(p. 165). Nitrogen usually is pentatomic and triatomic ; in the cited 
compounds it appears di- and tetra-tomic. This abnormal behavior 
of N finds a partial explanation in the position it occupies in the peri- 
odic system of the elements. 
2N0 = N2 + 02 
2 vols. 1 vol 1 vol. 
Nitrous Oxide—Ilyponitrous Oxide—N20—is formed when 
zinc or tin acts upon dilute nitric acid. It may be best 
obtained by heating ammonium nitrate, which at about 170° 
breaks up directly into water and nitrous oxide:— 
NH4N03 = N20 + 2H20. 
This compound is a colorless gas, of sweetish taste and 
slight odor. Its density is 22 (H = 1), or 1.52 (air = 1), 
corresponding to the molecular formula N20 — \4. In cold 
water it is tolerably soluble (1 volume H20 dissolves at 0° 
1.305 volumes N20); therefore it must be collected over 
water or mercury; cooled to 88°, or under a pressure 
of 30 atmospheres at 0°, it condenses to a colorless liquid of 
specific gravity 0937. By evaporation of the liquid in the 
air its temperature diminishes to — 100°, and solidifies to a 
crystalline, snowy mass. If the aqueous nitrous oxide be 
evaporated under an air pump its temperature falls to 
— 140° ; the lowest which has been attained. Although 
this oxide contains less oxygen than nitric oxide, it sup- 
ports the combustion of many bodies more readily than the 
latter, because more easily decomposed into oxygen and nitro- 
gen. A glimmering chip inflames in it, as in oxygen ; phos- 
phorus burns with a bright luminous flame, while one of 
sulphur is extinguished. The liquid nitrous oxide behaves 
like the gas; a glowing coal thrown on its surface burns 
with a bright light. A mixture of equal volumes of nitrous 
oxide and hydrogen explodes like detonating gas, only less 
violently :— 
N,0 + H, = N2 + H20. 
1 TOl. 1 VOL 1 VOl. 204 
INORGANIC CHEMISTRY. 
From oxygen, to which it is very similar, nitrous oxide may 
be distinguished by its not affording brown vapors with nitric 
oxide, as is the case with the former. It is not capable of 
combining with oxygen. Conducted through a glowing tube 
it is converted into nitrogen and oxygen. Inhaled in slight 
quantity it has an exhilarating effect, therefore termed laugh- 
ing 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 a 
volume of N20 an equal volume of nitrogen will be separated 
—corresponding to the molecular formula:— 
N20 + K2 = N2 + k2o. 
1 vol. 1 vol. 
Hyponitrous Acid—NOH = O = NH. As nitrous and nitric 
acids correspond to nitrogen trioxide and pentoxide, so may hypo- 
nitrous oxide be regarded as the anhydride of the recently dis- 
covered hyponitx-ous acid :— 
although the latter is not formed by the hydration of N20. On the 
contrary, hyponitrous acid, by removal of water by sulphuric acid, 
yields nitrous oxide. 
The acid is got in aqueous solution by decomposition of its silver 
salt AgNO by means of hydrochloric acid. Dissolved in water it is 
colorless, reacts strongly acid and is tolerable stable. It liberates 
iodine from potassium iodide and reduces a permanganate solution. 
With silver nitrate the silver salt is again precipitated. 
The silver salt AgNO is obtained by action of sodium amalgam 
upon a solution of potassium nitrite, neutralization with acetic acid and 
precipitation by silver nitrate. It is a lie ht yellow, amorphous powder, 
not altered by the light. Heated above 110° it decomposes with 
explosion. Soluble in dilute sulphuric acid, and upon neutralization 
with ammonium hydrate it is reprecipitated. By concentrated sulphuric 
acid it, as well as the solution of the free acid, is decomposed, with 
liberation of N20. 
N20 + H20 = 2NOH, 
In their thermo-chemical deportment is found the explanation of 
the varying deportment of the oxygen compounds of nitrogen, their 
little stability, and oxidizing properties, as well as their mode of forma- 
tion. All nitrogen oxides are endothermic compounds, i. e., they are 
produced from their elements with heat absorption (compare p. 176) 
corresponding to the symbols ;— 
(N20) = —20600 2(N,0) = —43200 (N2,03) = —22200 
2(N,02) = —6200 (N2,05 = gas) = —1200. OXYGEN COMPOUNDS OF PHOSPHORUS. 
205 
Proceeding from nitric oxide (NO), we observe from the above 
numbers that the formation of the higher oxides from it occurs with 
heat disengagement and, indeed, the latter becomes regularly less for 
each further combined oxygen atom :— 
2(NO,0) = + 2700 (N203,0) = + 17000 (2N02,0) =-f 4000. 
A similar decrease in heat disengagement, according to multiple 
proportion, is also observed in other compounds. The great absorp- 
tion of heat in nitrous oxide (N20 = 21000), with which also its 
difficult condensability agrees, is to be interpreted by the fact that the 
affinity of N for O is weaker than that of the oxygen atoms in the 
molecule of oxygen (compare p. 202). 
Heat disengagement, on the contrary, occurs in the production of 
nitric acid from its elements :— 
(N,03H—liquid) = 41600 (N,03,H,Aq) =48800. 
This explains the relative stability of nitric acid. 
2. OXYUEN COMPOUNDS OF PHOSPHORUS. 
P 2 0 3 
Phosphorus 
trioxide. 
po2h3 
Hypophosphorous 
acid. 
po3h3 
Phosphorous 
acid. 
p205 
Phosphorus 
pentoxide. 
po4h3 
Orthophosphoric 
acid. 
From orthophosphoric acid the two following anhydride 
acids (compare p. 192) are derived:— 
HP03—Metaphosphoric acid. 
H4P207—Pyrophosphoric acid 
The structure of these compounds is expressed by the fol- 
lowing formulas:— 
H2PO — OH 
hpo/oh 
mu\OH 
v /OH 
PO—OH 
\OH 
In hypophosphorous acid two atoms of hydrogen are directly 
combined with pentatomic phosphorus, while the third atom 
forms an hydroxyl group with oxygen. The former, by the 
action of bases, is easily replaced, and, therefore, hypophos- 
phorous acid is a monobasic add. Phosphorous acid contains 
one atom of H united to P and two hydroxyl groups; therefore, 
it is dibasic. Finally, phosphoric acid has three hydroxyl 
groups, and forms three series of salts. By the elimination 
of one molecule of H20 from H3P04, metaphosphoric acid re- 
Hypophosphorous 
acid. 
Phosphorous 
acid. 
Phosphoric 
acid. 206 
INORGANIC CHEMISTRY. 
suits—an anhydride, which, at the same time, is a monobasic 
acid, as it contains one hydroxyl group :— 
y 
P02 — OH—Metaphosphoric acid. 
On removing one molecule of II20 from two molecules 
Ii2P04, pyro- or diphosphoric acid is formed, (see p. 186) :— 
V /OH /OH 
PO—OH PO—OH 
v >oh-h’° = >° 
PO-OH PO—OH 
\0H \0H 
Pyrophosphoric acid contains four hydroxyl groups, hence 
is tetrabasic. 
Finally, if, from two molecules of phosphorous acid, or phos- 
phoric acid, all the II atoms be removed, in the form of water, 
two perfect anhydrides are produced :— 
2 Molecules Phosphoric Acid. 
1 Molecule Diphosphoric acid. 
Ill III 
OP - 0 —PO 
and 
02P —0P02 
The salts of phosphoric acid are termed phosphates; those 
of phosphorous acid, phosphites, and of hypophosphorous acid, 
hypophosphites. 
Phosphorous 
anhydride. 
Phosphoric 
anhydride. 
Hypophosphorous Acid—H:1P02. Hydrogen phosphide 
escapes when a concentrated solution of sodium or potassium 
hydrate is warmed with yellow phosphorus, leaving behind in 
solution, a salt of hypo-phosphorous acid. The free acid 
may he separated from the barium salt by means of sulphuric 
acid ; the insoluble barium sulphate being filtered off* from 
the aqueous solution of the acid, and the latter concentrated 
under the air-pump. Hypophosphorous acid is a colorless, 
thick liquid, with a strong acid reaction. Below 0° it some- 
times solidifies to large white leaflets, which fuse at + 17.4°. 
Heat converts it, with much foaming, into hydrogen phos- 
phide and phosphoric acid :— 
2P0 2H3 = PH3 + PO4H3. 
It readily absorbs oxygen, becoming phosphoric acid, hence 
acts as a powerful reducing agent. It reduces sulphuric acid 
to sulphur dioxide, and even to sulphur. From their solu- 
tions it precipitates many of the metals; from copper sulphate 
it separates the hydride—Cu2H2. PHOSPHORUS TRIOXIDE. 
207 
The acid is monobasic. Its salts are easily soluble in 
water and absorb oxygen from the air, thus becoming phos- 
phates. Heated in dry condition they set free the hydride 
of P and are converted into pyrophosphates; some also 
yield metallic phosphides. 
Phosphorous Acid—H,P03—is formed at the same time with 
phosphoric acid in the slow oxidation of P in the air. The 
decomposition of the trichloride by water gives it more con- 
veniently :— 
PC13 + 3H20 = PO3H3 + 3HC1. 
By evaporation of this solution under the air pump the 
phosphorous acid becomes crystalline. The crystals are readily 
soluble in water and deliquesce in the air. It fuses at 70° 
and decomposes on further heating into PH, and phosphoric 
acid:— 
In the air the acid absorbs oxygen and changes to phos- 
phoric acid. Plence it reduces strongly and precipitates the 
free metals from many of their solutions. In presence of 
water the halogens oxidize it to phosphoric acid. It is a 
dibasic acid, forming two series of salts, in which 1 and 2 
atoms* of H are replaced by metals. 
In the air the phosphites do not oxidize, excepting by 
means of oxidizing agents. When heated they generally 
decompose into hydrogen and pyrophosphates. 
4P03H3 — PH3 -p 3P04H2. 
Phosphorus Trioxide — P203— Phosphorous anhydride, 
results from passing a slow, dry air current over gently 
heated phosphorus; further, by the action of the trichloride 
upon phosphorous acid :— 
It forms white, voluminous flocks, which readily sublime 
and possess a garlic-like odor. Water dissolves the oxide, 
forming phosphorous acid. From the air it attracts oxygen 
and moisture very energetically, yielding phosphoric acid. 
P03H3 + PC13 = P203 + 3HC1. 
* Therefore, the structural formula HPO (OH)2 is assigned to this 
acid. There appears to exist another phosphorous acid, at least in 
compounds, to which the formula P(OH)3 belongs. 208 
INORGANIC CHEMISTRY. 
Phosphoric Acid — P04Hs, or Orthophosphoric acid, is 
formed when the pentoxide is dissolved in hot water, and by the 
decomposition of the penta- or oxy-chloride (POCl)3 by water 
(see p. 130). It may be obtained by decomposing bone ash 
(P04)2Ca3 with sulphuric acid, or, better, by oxidizing yellow 
phosphorus with nitric acid. The aqueous solution is evapo- 
rated to dryness in a pewter 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 called acid (P04H2K), neutral (P04HK2) 
and basic (P04K3). As this designation does not entirely 
correspond with the behavior of the salts to litmus, it is more 
rational to term them primary, secondary and tertiary; or to 
speak of them according to the number of hydrogen atoms re- 
placed by metals, as e. g., monopotassium phosphate (II2KP04), 
dipotassium (K2HP04) and tripotassium ( K:!P04) phosphate. 
The tertiary phosphates, excepting the salts of the alkalies, 
are insoluble in water. With a silver nitrate (AgN03) solu- 
tion soluble phosphates give a yellow precipitate of tri-silver 
phosphate P04Ag3. 
Pyrophosphoric Acid — H4P207—(structure p. 206) is 
formed by the continuous heating of orthophosphoric acid to 
200-300°, until a portion of it in ammonium.hydrate does not 
yield a yellow but pure white precipitate with silver nitrate. 
The sodium salt is easily obtained by heating di-sodium phos- 
phate :— 
The acid presents a white crystalline appearance, and is 
readily soluble in water. In solution at ordinary temperatures 
it slowly, by warming rapidly, takes up water, and like all 
anhydrides, passes into the corresponding acid—orthophos- 
phoric acid. It is tetrabasic. Its salts are very stable and 
are not altered by boiling with water ; warmed with acids they 
become salts of the ortho-acid. The soluble salts give a white 
precipitate, Ag4P207, with silver nitrate. 
2Na2HP04 = Na4P207 + H20. 
Metaphosphoric Acid—HP03 or P02.0H—results from 
heating the ortho- or pyro- acid to 400°. It can be more con- 
veniently obtained by dissolving the pentoxide in cold water : 
A glassy, transparent mass (Acidum phosphoricum glaciale) 
which fuses on heating and volatilizes at higher temperature 
P205 + H20 = 2HP0S. PHOSPHORUS PENTOXIDI 
209 
without suffering any change. It deliquesces in the air and 
dissolves with ease in water. (The commercial glassy phos- 
phoric acid contains sodium and magnesium phosphate and 
dissolves with difficulty in water.) The solution coagulates 
albumen; this is a characteristic method of distinguishing 
the meta- from the ortho- and pyro- acids. In aqueous 
solution the acid changes gradually, by boiling rapidly, into 
the ortho- acid:— 
hpo3 + h2o = h3po4 
It is a monobasic acid. Its salts, the metaphosphates, are 
readily obtained by the ignition of the primary salts of the 
ortho- acid :— 
Boiling the aqueous solutions of these salts converts them 
into the ortho-primary salts. With silver nitrate the soluble 
metaphosphates give a white precipitate, AgP03. 
NaH2P04 = NaPOs + H20. 
In addition to the ordinary salts of metaphosphoric acid, various 
modifications of the same exist: these are derived from the polymeric 
ineta-acids. II2P206, H3P309, H4P4O12, etc. These acids arise from 
the corresponding polyphosphoric acids, which are formed by the 
union of n molecules of the ortho acid, with the separation of n — 1 
molecules of water (p. 18(i), just as the mtta acid is formed from the 
ortho. Boiling their solutions converts all the above into primary 
orthophosphates. 
Phosphorus Pentoxide—P205 or Phosphoric anhydride, is 
formed by burning phosphorus in a current of oxygen or dry 
air. The following procedure serves for the preparation of it 
(Fig. 86):— 
In the glass balloon A a piece of P placed in an iron dish attached to 
ab is burned. The necessary amount of air is drawn through the 
vessel by means of an aspirator; to dry it perfectly it passes fii>t 
through the bent tube containing pieces of pumice stone moistened 
with sulphuric acid. After the phosphorus has been consumed, fresh 
pieces of it are introduced into the little dish through ab, and the upper 
end of the tube closed with a cork. The P205 formed collects partly in 
A and partly in the receiver. 
Phosphorus pentoxide forms a white, voluminous, floccu- 
lent mass, which is fixed. It attracts moisture energetically 
and deliquesces in "With hissing it dissolves in cold 
water to metaphospftbric apid. Owing tp its great affinity for 
water it serves as aji >gent for drying gases, and also for the 
withdrawal of Water* from many substances. 210 
INORGANIC CHEMISTRY. 
Fig. 86. 
Chlor-Anhydrides of the Acids of Phosphorus. The 
halogen derivatives of P considered on page 129 may be viewed as the 
halogen anhydrides of phosphorous and phosphoric acids (p. 187). The 
compounds PC13, PBr3 and PI3 are derived from phosphorous acid, 
because with water they yield the latter acid : — 
PC13 + 3H2 0 = H3 P03 + 3HCI. 
The compounds POCl3, POBr3 are the halogen anhydrides of 
phosphoric acid : — 
while PC15 and PBr5 correspond to the normal hydroxide P (OH)5, 
which has not been obtained in a free condition. 
Analogous to the o-xjmhloride—POCl3—is the compound PSC13. 
It is obtained by the action of PC15 upon hydrogen sulphide and some 
metallic sulphides: — 
P0C13 + 3H20 = PO (OH), + 3HC1; 
a reaction very similar to that occurring in the formation of phosphorus 
oxychloride. Phosphorus sulpkochloride—PSC13—is a colorless 
liquid, fuming in the air and boiling at 124°. Water decomposes it 
into phosphoric and hydrochloric acids and hydrogen sulphide. 
PC15 + H2S = PS Cl 3 + 2HC1, 
COMPOUNDS OF PHOSPHORUS WITH SULPHUR. 
With sulphur phosphorus forms a number of compounds 
which are obtained by direct fusion of P with S. As the 
union of ordinary P with S occurs usually with violent explo- ARSENIC TRIOXIDE. 
211 
sion, to produce these compounds red phosphorus should be 
employed. 
The compounds P2S3 and P2S5, analogously constituted to 
P203 and P206, are solid crystalline substances, melting at 
higher temperatures and subliming without decomposition. 
Water changes them to hydrogen sulphide and the corres- 
ponding acids, phosphorous and phosphoric. They combine 
with metallic sulphides to compounds {e. g., PS4K3) which 
are analogously constituted to the salts of phosphoric acid 
(see sulpho-salts of arsenic). 
At ordinary temperatures, P2S and P4S are liquids, which 
inflame readily in the air. 
Resides the preceding, other phosphorus derivatives containing N 
exist. These have been little studied and, at present, offer little 
interest. Such compounds are PN2H (phospham), PNO, PNC12. 
On allowing ammonia to act upon POCl3, there arise, by replacement 
of Cl by the group NH2 (amido) the so-called amid derivatives: 
P0C12NH2, P0C1(NH2)2 and PO(NH2)3. 
8. OXYGEN DERIVATIVES OF ARSENIC. 
As203 
Arsenic trioxide. 
As03H3 
Arsenious acid. 
As 2 0 5 
Arsenic peutoxide. 
As04H3 
Arseuic acid. 
Arsenic Trioxide, As20.,, or Arsenious anhydride, 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. On a large scale it is obtained 
metallurgically, as a bye-product in the roasting of ores con- 
taining arsenic. The trioxide which is thus formed volatil- 
izes 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 form of a transparent, amorphous, glassy 
mass (arsenic glass), the specific gravity of which equals 3.69. 
Upon preservation, this variety gradually becomes non-trans- 
parent and porcelanous, acquires a crystalline structure, and its 
specific gravity increases to 3.74. Upon dissolving this oxide 
in hot hydrochloric acid, it crystallizes on cooling, in shining, 
regular octahedra. At the same time, the interesting phe- 
nomenon is observed, that the solution of the’ glassy variety 
upon crystallizing phosphoresces strongly in the dark, while 212 
INORGANIC CHEMISTRY. 
the porcelanous does not exhibit this property. Arsenic 
trioxide crystallizes in similar forms of the regular system, 
when its vapors are rapidly cooled, but upon cooling slowly, 
it assumes the shape of rhombic prisms; therefore, it is dimor- 
phous. Heated in the air, it sublimes above 218°, without 
fusing; by higher pressure, however (in sealed tubes), it melts 
to a liquid which, upon solidifying, is glassy. 
The vapors of As203 have the vapor density 198 (H = 1). Corres- 
ponding to formula As203 ( = 198) the vapor density should be J-f5- 
= 99. From the experimentally determined vapor density, which is 
doubly as great, it follows that the gaseous molecules of the trioxide 
possess the double formula As406. Before we noticed that the mole- 
cule of free arsenic also consists of four atoms (As4); in arsenic tri- 
oxide this complex arsenic group, consequently, is retained ; while in 
arsine (AsH3), and arsenious chloride (AsCl3) the molecules contain 
but 1 atom of arsenic. 
The trioxide dissolves with difficulty in water; the solu- 
tion possesses a sweetish, unpleasant, metallic taste, exhibits 
but feeble acid reaction and is extremely poisonous. Very 
soluble in acids and forms salts, probably, with them ; at least, 
on boiling a solution of As203 in strong hydrochloric acid 
arsenious chloride, AsC13, volatilizes. From this and its feeble 
acid nature we perceive an indication of the basic character 
of the trioxide corresponding to the already partially metallic 
nature of arsenic (see p. 139). 
Nascent hydrogen converts the trioxide into arsine (AsH3) ; 
heated with charcoal it is reduced to the metallic state. 
Upon heating As203 in a narrow glass tube with C, the 
reduced arsenic deposits as a metallic mirror on the sides. 
Oxidizing agents convert it into arsenic acid. 
Arsenious Acid—H3As03—corresponding to As203, is not 
known in a free condition. It probably exists in the aqueous 
solution, but upon evaporation the anhydride separates out. 
In its salts (arsenites) it is tribasic and affords mostly tertiary 
derivatives : Ag3As03, Mg3(As03)2. The alkali salts, soluble 
in water, absorb oxygen from the air and serve as powerful 
reducing agents, they themselves becoming arsenates. 
Other salts derived from the meta-arsenious acid, HAs02, 
also exist. 
Arsenic Acid—H;,As04—is obtained by the oxidation of 
arsenic or its Trioxide with concentrated nitric acid or by 
means of chlorine. Upon evaporating the solution rhombic ARSENIC TRISULPHIDE. 
213 
crystals of the formula H3As04 -j- iH20 separate out; these 
deliquesce on exposure. They melt at 100°, lose their water 
of crystallization and yield orthoarsenic acid H3As04, which 
heated to 140-180° passes into pyroarsenic acid—H4As207:— 
2H3As04 = As207H4 -j- H20. 
At 200° this again loses water and becomes Meta-arsenic 
acid—HAs03. With water the last two acids become ortho 
again; hence the latter is perfectly analogous to phosphoric 
acid. 
At a red heat the meta-arsenic acid loses all its water and 
becomes Arsenic Pentoxide—As205, a white, glassy mass. 
Very strong ignition breaks this up into As203 and 02; in 
contact with water it gradually changes to arsenic acid. 
Orthoarsenic acid is readily soluble, and is a strong tribasic 
acid. Its salts—the arsenates—are very similar to the phos- 
phates and are isomorphous with them. With the soluble salts 
silver nitrate gives a reddish-brown precipitate of trisilver- 
arseniate, Ag3As04, 
COMPOUNDS OF AKSEXIC WITH SULPHUR. 
Like phosphorus, arsenic, upon fusion with sulphur, yields 
several compounds. In these derivatives is shown the metallic 
nature of arsenic, because they, according to the common 
method of forming the metallic sulphides, can be obtained by 
the action of hydrogen sulphide upon the oxygen derivatives 
of arsenic:— 
As203 “I- 3H2S = As2S3 -)- 3H20 
Arsenic Trisulphide—As2S3—is precipitated from solutions 
of arsenious acid or its salts by hydrogen sulphide, as a lemon- 
yellow amorphous powder. Also from solutions of arsenic 
acid we get the trisulphide (at same time mixed with S), as it 
is at first reduced to arsenious acid:— 
This compound is readily prepared by fusing As203 with 
sulphur. In nature it occurs as auripigment, in the form of 
a brilliant, leafy, crystalline mass of gold-yellow color, of 
specific gravity 3.4. On fusing artificially prepared arsenic 
trisulphide it solidifies to a similar yellow mass, the specific 
gravity of which equals, however, 2.7. In water and acids 
the trisulphide is insoluble, dissolves readily in ammonium 
hydrate and the alkalies. 
As205 + 2H2S = As203 + 2H20 -f- 2S. 214 
INORGANIC CHEMISTRY. 
Arsenic Pentasulphide—As2S5—separates from the solution 
of sodium sulph. arseniate, Na3AsS4 (see below), upon the 
addition of acids, as a bright yellow powder. 
The Arsenic Disulphide—AsS2—also exists. It occurs in 
nature as Realgar, forming beautiful, ruby-red crystals, of 
specific gravity 3.5. These, as powder, find application as a 
pigment. It is prepared artificially by fusing As with S. 
Arsenic Sulpho-Salts.—Owing to the similarity of sulphur to 
oxygen we may anticipate for arsenic (as also for other elements) the 
existence of sulphur acids corresponding to the oxygen acids, e. g., 
sulpharsenious acid, H3AsS3, and sulpharsenic acid, H3AsS4. How- 
ever, these acids are unknown in a free state, although their salts, 
knovn as sulphides and sulpho-salts, are found, and they, too, corres- 
pond perfectly with the oxygen salts. Just as the latter arise by the 
union of metallic oxides with acid oxides, so are formed the sulpho- 
salts by the combination of alkaline sulphides with the oxides of sul- 
phur derivatives:— 
As2S3 + 3K2S = 2K3AsS3 
Tripotassium sulpharsenite. 
As2S5 -j- 8K2S = 2K3AsS4 
Tripotassium sulpharseniate 
For the preparation of these sulphosalts, arsenic sulphide is dissolved 
in the aqueous solution of potassium or sodium sulphide, or hydrogen 
sulphide is conducted through the alkaline solution of the oxygen 
salts:— 
The sulphosalts of the alkalies and ammonium are easily soluble in 
water and, on evaporating the solution, generally separate in crystals. 
Acids decompose them, arsenic sulphide separating out and hydrogen 
sulphide becoming free:— 
K3As04 + 4H2S = K3AsS; + 4H20. 
2K3 AsS4 + 6HC1 = As2S5 + 6KC1 + 3H2S. 
Antimony, carbon, tin, gold, platinum and some other metals form 
sulphosalts similar to those of arsenic (and also phosphorus). 
The oxygen derivatives of antimony are in constitution 
analogous to those of arsenic : Sb203 and Sb205. In them is 
expressed more distinctly the metallic nature of antimony, 
which we observed appearing also in the halogen derivatives. 
The lowest oxygen compound does not possess acid, but almost 
solely basic properties; only forms salts with acids, hence 
called Antimony oxide. The normal hydrate II3Sb03, corres- 
4. OXYGEN COMPOUNDS OF ANTIMONY. ANTIMONY OXIDE. 
215 
ponding to arsenious acid, H3As03, is not known. An hydrate 
Sb0.2tl or SbO.OH, analogous to meta-arsenious acid, does 
exist; it deports itself like a base. 
The higher oxidation product, the pentoxide Sb205, on the 
contrary, has an acid nature and forms salts with the bases. 
The hydrate, Sb04H3 or ortho-antimonic acid and its salts 
have not been obtained. The known salts are derived from 
pyro-antimonic acid, II4Sb207 and meta-antimonic acid, 
HSb03: these exist in a free condition. 
Antimony Oxide—Sb203—is obtained by burning the 
metal in the air or by oxidizing it with dilute HN03. By 
sublimation it may be obtained in twro different crystal 
systems, in regular octahedra and in rhombic prisms. 
Arsenic trioxide also crystallizes in the same forms ; therefore 
the two compounds are isomorphous. On adding sodium 
carbonate to the solution of the trichloride a white precipitate 
of antimony hydrate or antimonious acid, IISb02, separates 
out:— 
2SbCl3 + 2Na2C03 + HaO = 2SbO.OH + 6NaCl + 3C02. 
Boiling changes the hydrate to oxide. The latter and the 
hydrate are soluble in sodium and potassium hydrate, very 
probably, forming salts (NaSb02) which decompose upon evap- 
orating the solution. In this behavior is perceived also the 
acid nature of antimony hydrate, therefore it has received the 
name of antimonious acid. 
The acid forms salts with acids which are derived either 
from the normal hydrate, H3Sb03, or from the hydrate, HSb02= 
SbO.OH. In the salts of the first kind we have 3 hydrogen 
atoms of the hydrate replaced by acid radicals, or, what is the 
same, a triatomic antimony atom displacing 3 atoms of hydro- 
gen of the acids:— 
Sb03(N02)3 or (N03)3Sb. 
Antimony nitrate. 
In the second variety of antimony salts derived from the 
hydrate, SbO.OH, hydrogen is replaced by a monatomic acid 
residue, or the hydrogen of the acid is substituted by the mon- 
atomic group, SbO, known as antimonyl:— 
Sb0.0.N02 or N03.Sb0. 
Antimonyl nitrate. 
Antimony Sulphate—(S04)3Sb2—which separates when a 
solution of the oxide in sulphuric acid is cooled. 
Of these salts may be mentioned the following :— 216 
INORGANIC CHEMISTRY. 
Antimonyl Sulphate—S04(SbO)2—is formed when antimony 
oxide is dissolved in somewhat dilute sulphuric acid and on cooling 
crystallizes in fine needles. Water decomposes both, forming basic 
salts ; hence the basic nature of antimony oxide is slight. 
Antimonic Acid —HSb03—or more correctly Metantimonic 
acid, is formed upon warming antimony with concentrated 
nitric acid, and is a white powder, almost insoluble in 
water and in nitric acid, but reddens blue litmus paper. It 
is a weak monobasic acid, the salts of which are mostly 
insoluble. 
If antimony pentachloride be mixed with much water 
Pyroantimonic Acid, H4Sb207, separates as a white powder. 
Its salts are formed by the fusion of antiinonic acid or meta- 
antimonates with potassium or sodium hydrates :— 
2KSb03 + 2K0H = K4Sb207 + H20 
Hydrochloric acid precipitates pyroantimonic acid from the 
solutions of these salts. 
By gentle ignition the meta and pyro-acids yield Antimony 
Pentoxide, Sb205, a yellow, amorphous mass, soluble in hy- 
drochloric acid. By heating the oxygen compounds for some 
time with air access, they are converted into the oxide, Sb204, 
which can be viewed as antimonyl antimoniate, Sb03.SbO, 
or as a mixed anhydride 2 j O. It is a white powder, 
becoming yellow when heated and non-volatile. 
COMPOUNDS OF ANTIMONY WITH SULPHUR. 
These are perfectly analogous to the S compounds of 
arsenic and form sulphosalts wTith alkaline sulphides corres- 
ponding to the oxygen salts. Acids precipitate antimony 
sulphide from the sulphosalts. 
Antimony Trisulphide—Sb2S3—is found in nature as 
stibnite, in radiating crystalline masses of dark gray color and 
metallic lustre: specific gravity = 4.7. When heated it 
melts and sublimes. The artificial sulphide obtained by pre- 
cipitating a solution of the oxide with hydrogen sulphide is an 
amorphous red powder. Fused it solidifies to a mass exactly 
like stibnite. The sulphide dissolves in concentrated HC1, 
upon warming, to form antimony trichloride. 
The compound Sb2S20, occurring in nature as red stibnite, SODIUM SULPHANTIMONIATE. 
217 
can be artificially prepared, and serves as a beautiful red color, 
under the name of antimony cinnabar. Kermes minerale, 
employed in medicine, is formed by boiling antimony sulphide 
with a sodium carbonate solution, and'is a mixture of Sb2S3 
and Sb203. 
Antimony Pentasulphide—Sb2S5—or gold sulphur (sulfur 
auratum) is precipitated by H2S from acid solutions of anti- 
monic acid ; more conveniently obtained by the precipita- 
tion of sodium sulphantimoniate Na3SbS4 with hydrochloric 
acid:— 
2Na3SbS4 +6HC1 = Sb2S5 + 6NaCl + 3H2S 
It is an orange-red powder, like the trisulphide; it decom- 
poses on being heated into Sb2S3 and S2. In strong hydro- 
chloric acid it dissolves with separation of sulphur and hydro- 
gen sulphide to antimony trichloride. 
Sodium Sulphantimoniate—Na3SbS4 (Schlippe’s salt) re- 
sults from boiling pulverized Sb2S3 with sulphur and sodium 
hydrate. Upon concentrating the solution, it crystallizes in 
large, yellow tetrahedra containing 9 molecules II20 (SbS4- 
Na3 -f- 9H20) : exposed to the air it becomes covered with a 
brown layer of Sb2S5. It serves principally for the prepara- 
tion of the officinal gold sulphur. 
The affinity of the elements of the nitrogen group to hydro- 
gen diminishes successively with increase of atomic weight, 
agreeing with the addition of metallic character, while the 
affinity to chlorine, concluding from the thermo-chemical 
relations, in general increases (compare p. 139). However, 
the heat disengagement in the formation of AsC13 is somewhat 
less than that in the case of PC13, from which the non-existence 
of the compounds AsX5 (see p. 134) finds a partial explana- 
tion. The slight affinity of arsenic expresses itself yet more 
distinctly in the oxygen compounds, just as in the case of the 
halogen and oxygen group (p. 176), the arsenic corresponding 
to bromine and selenium— 
Br = 79.7 Se = 79 As = 75. 
shows, in the formation of its compounds, a less development 
of heat:— 
(N,04,Hs, Aq) = 117400 
(P,04,H3,Aq) = 305300 
(As,04,H3,Aq) = 215200 
(N205.Aq) = 29800 
(P2.05) = 363800 
(As203) = 219400 218 
INORGANIC CHEMISTRY. 
Phosphoric acid is, therefore, more stable and more energetic than 
nitric and arsenic acids ; nitric acid oxidizes phosphorus and arsenic 
to phosphoric and arsenic acids. The latter acid is readily reduced 
to arsenious acid. 
4. Oxygen Derivatives of the Elements of the Carbon Group. 
According to analogy with the hydroxyl derivatives of the 
elements of the three first groups :— 
cio3,oh. 
S02(0H)2. 
PO(OH)3, 
we may conclude the existence, for the tetratomic elements 
—carbon, silicon and tin, of the following normal hydrox- 
ides, corresponding to the halogen compounds, CCl4,SiCl4, and 
SnCl4:— 
IV 
C(0H)4, 
Normal 
Carbonic acid. 
Si(OH)4. 
Normal 
Silicic acid. 
Sn(0II)4 
Normal 
Stannic acid. 
These normal hydrates or acids have but little stability, 
and exist mostly only in some derivatives. By the separation 
of a molecule of water, they pass into 
Or 
co3h2 
IV 
CO(OH)2 
Carbonic acid. 
Si03H2 
SiO(OH)2 
Silicic acid. 
Sn03H2 
SnO.(OH)2 
Stannic acid. 
These hydroxyl derivatives deport themselves toward the 
normal just as the meta-acids of the elements of the N-groups 
do to the ortho-acids (see p. 192). They constitute the ordinary 
acids of the tetratomic elements, carbon, silicon and tin, and 
are, as they contain 2 hydroxyl groups, 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, 0 = 16, FI = 19. 
Fluorine and oxygen do not afford any oxygen acids. The 
normal nitrogen acid, NO(OH)3, is very unstable, and passes 
into the meta-acid, HN03. Corresponding to this is, also, the 
normal carbonic acid (C(OII)4) not capable of existing. 
Indeed, the meta or ordinary carbonic acid, H2C03, is also 
very unstable and decomposes, when separated from its salts, 
at once into water and carbon dioxide, C02. Even silicic and 
stannic acids break up readily into water and their oxides:— 
co2 
Carbon dioxide 
Si02 
Silicon dioxide. 
Sn02 
Stannic oxide. OXYGEN COMPOUNDS OF CARBON. 
219 
Carbon Dioxide—C02—or carbonic anhydride (generally 
called carbonic acid). It is produced when carbon or its 
compounds are burned in air or oxygen. Found free in the 
air (in 100 volumes, upon average, 0.05 volumes C02), in 
many mineral springs (acid springs), and escapes in large 
quantities from the earth in many volcanic districts. It is 
prepared on a large scale by burning coke ; in the laboratory 
it may be most conveniently obtained by the decomposition 
of calcium carbonate (marble or chalk) with dilute hydro- 
chloric acid :— 
1. OXYGEN COMPOUNDS OF CARBON. 
CaC03 + 2HC1 = CaCl2 + C02 + H30. 
Calcium carbonate. Calcium chloride. 
Carbon dioxide is a colorless gas, of sweetish odor and taste. 
Its gas density equals 1.524 (air = 1), or 22 (H = 1), corres- 
ponding to the molecular formula, C02 = 44. Owing to its 
weight, the gas may be collected by displacement of air, and 
may be poured from one vessel into another filled with air. 
Under a pressure of 36 atmospheres (at 0° C.), carbon dioxide 
condenses to a mobile, colorless liquid, not miscible with water, 
and boiling at —78°. The specific gravity of the liquid car- 
bon dioxide is 0.99 at —10°, at 0° 0.94. Hence, it expands 
more equally than even all gases, although, ordinarily, the 
coefficient of expansion of liquids is less than that of gases. 
Similar deportment is observed in the case of other bodies 
compressible, under pressure, to liquids. 
Above 32.5° carbon dioxide cannot be condensed, by any pressure, 
to a liquid, although it may be reduced to a smaller volume than that 
which the liquid C02 would equal In the same way all other coercible 
gases show a critical point in temperature at which they are no longer 
able to be condensed to liquids. * * * * That the so-called per- 
manent gases (as H, N, NO, 0) could not formerly be condensed was 
due to the fact that they were compressed at temperatures lying above 
their critical points. 
When liquid carbon dioxide stands exposed to the air so 
much heat is absorbed by the evaporation of a portion of it 
that the remainder solidifies to a white, snow-like mass. Solid 
carbon dioxide is a bad conductor of heat and evaporates 
but slowly. It can, therefore, although its temperature is 
-78°, be taken up in the hand and even introduced into the 
mouth without danger, as it is always encircled by a gaseous 
layer, and thus not in immediate contact with the skin; upon 
pressing it hard, however, between the fingers it causes a pain- 220 
INORGANIC CHEMISTRY. 
ful blister. By the evaporation of solid carbon dioxide under 
the air pump its temperature is lowered to -130° C. (compare 
p. 192). 
Water dissolves its volume of the gas at 14° ; at 0° it takes 
up 1.79 vols. For every pressure this proportion remains con- 
stant, i. e., at every pressure the same volume of gas is 
absorbed. As gases are condensed in proportion to the 
pressure, the quantity of absorbed gas is also proportional to 
the former. Hence 1 volume H20 absorbs (at 14°) and two 
atmospheres 2 volumes, at 3 atmospheres 3 volumes, etc., of 
carbon dioxide—measured at ordinary pressure. The quan- 
tity of gas absorbed at high pressure escapes when the latter 
is diminished with effervescence of the liquid; upon this 
depends the sparkling of soda water and champagne, which 
aresaturated with C02under higher pressure. Every naturally 
occurring water, especially spring water, holds C02 in solution, 
which imparts to it a refreshing taste. 
As the product of a complete combustion carbon diox- 
ide is not combustible, and is unable to support the combus- 
tion of most bodies; a glimmering chip is immediately 
extinguished in it. In similar manner it is non-respirable. 
Although it is not poisonous, in the true sense of the word, yet 
the admixture of a few per cent, of C02 to the air makes it 
suffocating, as it retards the separation of the same gas from 
the lungs. 
By continued action of the electric sparks it is decom- 
posed into carbonous oxide (CO) and oxygen ; upon heating 
to 1300° it suffers a partial decomposition (dissociation) into 
CO and O. Also decomposed when conducted over heated K 
or Na, with separation of carbon; the potassium combines with 
oxygen to form potassium dioxide:— 
C02 + 2K2 = C + 2K20, 
which, with excess of C02, forms potassium carbonate (K2C03). 
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 C02 it follows that 
in one volume of it an equal volume of O is contained. 
We may satisfy ourselves of this by burning C in a definite 
volume of O ; after cooling there is obtained an equal vol- 
ume of carbon dioxide :— 
c + o2 = co2. 
1 vol. 1 vol. CARBON DIOXIDE. 
221 
The experiment is most practically executed by aid of the 
apparatus of Hofmann pictured in Fig. 87. The spherical 
shaped expansion of the eudiometer limb of the V 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 the carbon to be burned lies, 
while the other wire terminates in a thin piece 
of platinum, which is in contact with the car- 
bon. For the performance of the experiment 
the air is expelled from the globe limb by 
means of a rapid current of oxygen, the stopper 
placed in air tight, the mercury level noted 
and the copper wires connected with the poles 
of an induction stream of 3-4 Bunsen elements, 
which induces the burning of the carbon. As 
the volume of the enclosed gas is greatly ex- 
panded by the heat developed, it is advisable, 
in order to avoid the jumping out of the stopper, 
to previously, by running out mercury, reduce 
the pressure of the gas to two-thirds. A prac- 
tical modification of this apparatus consists in 
having the usual cylindrical eudiometer limb 
provided with two side tubes, in which carbon 
electrodes can be inserted. 
Fig. 87. 
The Physiological Importance of Carbon Dioxide. The gas is 
present in the atmosphere and is inhaled by the plants. The chloro- 
phyl grains in the green parts of the plant decompose carbon dioxide in 
sunlight, with a partial separation of oxygen ; from the residue, by the 
mutual action of water and ammonia is formed the limited quantity of 
carbon compounds peculiar to plants. The respiration and life pro- 
cess of animals are essentially the reverse of the above. These absorb 
through the lungs the oxygen of the air, which, influenced by the blood 
corpuscles, oxidizes the substances present in the blood, and in this 
manner shapes the life process. The final products of the oxidation 
are carbon dioxide and water, which are exhaled. The absorption of 
0 by animals and its separation by plants, as also the reverse course 
of CO2, are about alike, so that the quantities of O and C02 in the 
air show no appreciable alteration. 
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 latter, 
the red disappears, in consequence of the evaporation of the 
carbon dioxide. 222 
INORGANIC CHEMISTRY. 
AVe may then regard it as probable that free carbonic 
acid H2C03 is contained in the aqueous solution, which, how- 
ever, decomposes readily into the dioxide C02 and water. 
The salts of carbonic acid are formed by the action of carbon 
dioxide upon the bases:— 
Carbon dioxide is therefore easily absorbed by potassium 
and sodium hydrate. On conducting it through a solution of 
calcium or barium hydrate there is produced a white precipi- 
tate of barium or calcium carbonate—CaC03. 
Carbonic acid is dibasic, forming primary (acid) and 
secondary (neutral) salts, C03HK and C03K2, called also 
carbonates. As the acidity of carbonic acid is only slight the 
secondary salts formed from strong bases exhibit a basic 
reaction. Most acids expel the weak carbonic acid from its 
salts, with decomposition into carbon dioxide and water. 
Carbon Monoxide—CO—produced in the imperfect com- 
bustion of carbon by insufficient access of air, and when 
carbon dioxide is conducted over red hot coals : C02 4- C = 
2CO. The forms of apparatus described, p. 221 serve for the 
demonstration of this volume relation. The monoxide is 
produced further, by glowing carbon with different metallic 
oxides, e. g., zinc oxide: ZnO 4* C = Zn 4- CO. For its 
preparation oxalic acid is warmed with sulphuric acid ; the 
latter withdraws water from the former and the residue 
breaks up into carbon dioxide and monoxide :— 
2K0H + C02 = C03K2 + H20. 
Potassium carbonate. 
c2o4h2 = co2 + co + h2o. 
The disengaged mixture of gases is conducted through an 
aqueous solution rtf sodium hydrate, by which the C02 is ab- 
sorbed, the monoxide passing through unaltered. Pure 
monoxide may be prepared by heating yellow prussiate of 
potassium (see Iron) with 9 parts H2S04. The resulting 
gas is colorless and odorless, and can only be condensed with 
difficulty. Its specific gravity is 14 (H = 1), corresponding 
to the molecular formula CO = 28. It is almost insoluble 
in water, but is dissolved readily by an ammoniacal solution 
of cuprous chloride (CuCl). Ignited, it burns in the air with 
a faintly luminous flame, to carbon dioxide. AVith air or 
oxygen it affords an explosive mixture:— 
2C0 + 02 = 2C02 
2 vols. 1 vol. 2 vols. CARBON DISULPHIDE 
In consequence of its oxidation it is capable of reducing 
most metallic oxides at a red heat— 
Burning bodies are extinguished by it. Inhaled, it acts very 
poisonously, even in slight quantity, as it expels the oxygen 
of the blood. The carbon vapor developed in heated ovens 
closed too soon, is carbon monoxide. As an unsaturated 
compound, this oxide, like ethylene, unites directly with 2 
atoms of chlorine, to yield carbon oxychloride, or phosgene 
gas COCl2:— 
CuO -f CO = Cu -f-C02 
CO + Cl 2 = COCl2 
1 vol. 1 vol. 1 vol. 
This is obtained by bringing together equal volumes of CO 
and CL in direct sunlight, or, better, by conducting CO into 
SbCl5. It is a colorless, suffocating gas, of specific gravity 49.4 
(H = 1), agreeing with the molecular formula COCl2 = 98.8. 
Water decomposes it into hydrogen chloride and carbon 
dioxide:— 
COCl2 + H20 — C02 -f- 2HC1. 
COMPOUNDS OF CARBON WITH SULPHUR. 
Carbon Disulphide. CS2 is formed like the dioxide, by 
the direct union of carbon and sulphur, if vapors of the 
latter are led over glowing carbon; the escaping disulphide 
vapors are condensed in a cooled receiver. It forms a 
colorless, mobile liquid, of characteristic odor, and refracts 
light strongly. Its specific gravity equals 1.27. It is very 
volatile and burns with a blue flame, to carbon dioxide and 
sulphurous acid. The mixture of the vapors with oxygen 
explodes violently when ignited. 
CS2 + 302 - C02 + 2S02 
1 vol. 3 voU. 1 vol. 2 vols. 
In nitrous oxide the vapors burn with a bright, blinding 
flame. On blowing a strong current of air upon carbon di- 
sulphide in a porcelain capsule (which conducts heat poorly), 
so much warmth is absorbed by the evaporation, that the re- 
sidual liquid solidifies to a white, snow-like mass. Carbon 
disulphide is insoluble in water ; it mixes, in every proportion, 
with alcohol and ether. It dissolves iodine with a violet red 
color, and serves as an excellent solvent for sulphur, phos- 
phorus, caoutchouc and the fatty oils. 224 
INORGANIC CHEMISTRV 
Carbon disulphide may be viewed as the anhydride of 
sulphocarbonic add—H2CS3. The salts of these acids are 
obtained by the solution of CS2 in alkaline sulphides (see 
sulphosalts p. 214) :— 
CSa + KaS = K2CS,. 
On adding hydrochloric acid to the solutions of these salts 
the sulphocarbonic acid separates as a reddish-brown oil. This 
decomposes readily. 
The sulphur compound corresponding to CO is not known : * 
there exists, however, one containing both oxygen and sulphur 
— Carbon oxysulphide, COS. It is produced (in small quantity) 
when a mixture of sulphur vapors and carbon monoxide 
gas is passed through red hot tubes and by heating carbon 
disulphide with sulphuric oxide:— 
Most readily obtained from potassium sulpbocyanide— 
CN. SK—see organic chemistry—by the action of dilute sul- 
phuric acid. It is a colorless gas, with an odor reminding 
one of carbon dioxide and hydrogen sulphide. It is present 
in some sulphur springs. It is very readily inflammable and 
burns with a blue flame :— 
CS2 + 3S03 = COS + 4S02. 
2C0S + 302 = 2C02 -f- 2S02. 
2 vols. 3 vols. 2 vols. 2 vols. 
It is decomposed at a red heat into CO and sulphur. Solu- 
ble in an equal volume of water, decomposing gradually into 
the dioxide and hydrogen sulphide :— 
cos + h2o = co2 + sh2 
Cyanogen Compounds. Of the innumerable compounds 
of C treated in organic chemistry, we will here mention only 
those of cyanogen, as they are of importance in inorganic 
chemistry. 
Nitrogenous carbon compounds heated with potassium 
hydrate yield potassium cyanide—CNK—which with iron 
forms the so-called yellow prussiate of potassium, K4Fe (CN)6. 
From these two compounds all the other cyanogen derivatives 
may be prepared. They all contain the group CN, called 
* Upon standing in sunlight CS2 is said to break up into S and CS 
—a chestnut-brown powder of specific gravity 1.66. HYDROGEN CYANIDE. 
225 
cyanogen. In it we have a triatomic nitrogen atom combined 
with a tetratomic carbon atom ; the fourth affinity of the latter 
III IV 
is not saturated: NEE C—: it is similar, therefore, to the groups 
OH, NH2, CH3, and is a monatomic radical. In chemical 
behavior the cyanogen group is very similar to the halogens 
chlorine and bromine; with the metals it forms metallic 
cyanides (KCN, AgCN) very similar to the haloid salts. 
Hydrogen cyanide is evolved when the cyanides are warmed 
with sulphuric acid :— 
2KCN + H2S04 = K2S04 + 2HCN 
Hydrogen Cyanide, HCN, is a colorless, mobile liquid, of 
peculiar odor, and boiling at 27°. It is an acid, forming salts 
with metals and bases, and is known as Hydrocyanic or Prussic 
acid. Both it and its salts are very powerful poisons. If 
the CN group is separated from its salts it doubles itself, form- 
ing dicyanogen or free cyanogen, C2N2 (N=C—C=N), because, 
like the other monatomic groups (as CH3, see p. 162), it can- 
not exist in a free condition. 
The heat occurring in the formation of the above cited simplest car- 
bon compounds corresponds with the symbols:— 
(C,0) = 28800. 
(C,0,S) = 1400. 
(C.H4) = 19500. 
(C0,0) = 68000. 
(C,S2) =—12600. 
(C2,H4) = — 4700. 
(C,02) = 96900. 
(C,N,H) =— 28*00. 
(C2,H2) = —48300. 
If an element combine with another according to multiple propor- 
tions, there occurs, in the union of the first atom, generally, a greater 
disengagement of heat than with the following atom (compare nitrogen 
oxides, p. 204). On the contrary, it is seen that the union of the 
second atom of oxygen with carbon (CO,0), sets free 68000 calories; 
that of the first atom (C,0), however, only 28800 calories. This can 
only be explained by the fact that, for the vaporization and disaggrega- 
tion of the solid carbon molecules, heat becomes latent (absorbed). 
If we assume that the direct union of the first atom also disengaged 
68000 calories, it would follow from this that, in the dissociation of 12 
weight parts carbon in gaseous free atoms, 39200 calories were ab- 
sorbed. From this, the heat absorption, in the production of CS2, 
CNH, C2H2, would be explained, while otherwise, in every direct 
chemical union, heat is invariably disengaged. 
Comparing the elements of the carbon group with each other, we 
discover that the heat disengagement is greatest with the compounds 
of silicon: — 
(C,C14) = 
(C,02) =96900. 
(Si, Cl4) = 157600. 
(Si,02) = 219000. 
(Sn,CI4) = 129200. 
(Sn,02,H20) = 133500. 
From these numbers we observe that tin dioxide, but not that of 
silicon can be reduced by carbon. 226 
INORGANIC CHEMISTRY. 
2. OXYGEN COMPOUNDS OF SILICON. 
Silicon Dioxide, Si02 (Silica), is widely distributed in na- 
ture as rock-crystal, quartz, sand, • etc. It is artificially 
obtained as a white, amorphous powder, of specific gravity 2.2, 
by the combustion of amorphous silicon, or by the ignition of 
silicic acids. In crystalline form, it only occurs in nature in 
figures of the hexagonal system, with the specific gravity, 2.6; 
these crystals are colorless, or colored by impurities. In the 
oxy-hydrogen flame it fuses to a transparent glass. 
Silicon dioxide is insoluble in water and all acids ; only by 
hydrofluoric acid is it decomposed with the formation of 
silicon fluoride (SiFl4) and water (p. 155). Strong ignition 
with carbon or potassium reduces it to metallic silicon. 
Boiled with potassium or sodium hydrate the artificially 
prepared dioxide dissolves; some of the naturally occurring 
amorphous varieties are also soluble, but not the crystallized 
dioxide. By fusion with the hydroxides or carbonates of the 
alkalies all varieties of silicic acid yield a glassy mass (water 
glass) soluble in water and containing silicates (K4Si04 or 
K2Si03). 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 
normal silicic acid, II4Si04:— 
Washed with water and dried in the air it becomes a white 
amorphous powder having the composition H2Si03. The 
Na4Si04 + 4HC1 = 4NaCl + H4Si04 
Fig. 88. 
freshly precipitated acid is somewhat soluble in water, more 
readily in dilute hydrochloric acid. On adding to the excess 
of dilute hydrochloric acid a solution of sodium silicate, 
the separated silicic acid remains dissolved. From the hydro- SILICON DIOXIDE. 
227 
chloric acid and sodium chloride solution we can obtain by 
dialysis a perfectly pure aqueous solution of silicic acid. 
Proceed as directed in following lines. Pour the hydrochloric 
acid solution into a wide cylindrical vessel, whose lower 
opening is covered with animal bladder or parchment paper, 
and the vessel (dialyser) then suspended in another con- 
taining pure water, Fig. 88. Osmosis now sets in. The 
sodium chloride and hydrochloric acid particles pass through 
the parchment paper into the outer water, while on the other 
hand, water particles pass from the outer vessel into the 
dialyser; to silicic acid the parchment paper is not permeable. 
This alternate diffusion of the different particles occurs until 
the outer and inner liquid show the same quantity of diffusible 
substances. Upon introducing the dialyser into a fresh 
portion of water, the dialysis commences anew. After 
repeated renewal of the external water there is found, 
finally, in the dialyser a perfectly pure silicic acid solution, 
free from sodium chloride and hydrochloric acid. The 
solution may be concentrated by evaporation ; then it readily 
passes into a gelatinous mass. The same occurs instanta- 
neously in dilute solutions if a trace of sodium carbonate be 
added or carbon dioxide be led into it. 
Like sodium chloride, all crystallizable soluble substances diffuse 
through parchment. These are known as crystalloids, to distinguish 
them from the non-diffusible colloids. To the latter belong gum, 
gelatine, albumen, starch, glue (eolla, hence the name colloid), and 
especially the most substances which occur principally in vegetable and 
animal organisms. Like silicic acid these colloids exist in liquid 
soluble and solid gelatinous condition. By dialysis many other sub- 
stances (like ferric and aluminum oxides) which ordinarily are insol- 
uble, can be brought into aqueous solution. 
We have already seen that acids like sulphuric, phosphoric, 
and arsenic, by the union of several molecules and the elimi- 
nation of water, are capable of forming anhydro- or poly-acids. 
Silicic acid is particularly inclined to that kind of condensa- 
tion. It forms a large number of poly-silicic adds, derived from 
the normal and ordinary acid, according to the common for- 
mula :— 
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 % H20 by ignition, 
represent such poly-acids. The natural silicates are the salts 
mSi(OH)4 — m1I20. 228 
INORGANIC CHEMISTRY. 
of such acids. The majority are derived from the acids: 
H2Si205, H4Si308, H2Si307, H4Si409 and others. Only a few 
silicates are obtained from the normal acid, e. g., chrysolite— 
Mg2Si04. 
Corresponding to CS2 is 
Silicon Disulphide, SiS2, which may he made by heating 
amorphous silicon with sulphur, or by conducting sulphur vapors over 
an ignited mass of silica and carbon. It sublimes in shining, silky 
needles, which water changes to silicic acid and hydrogen sulphide. 
To the group of carbon and silicon belongs also Tin (see 
p. 156), which forms perfectly analogous oxygen compounds. 
However, in them is plainly visible the metallic character of 
tin, and the basic of the oxides. The tin hydrates, Sn(OH)4 and 
SnO(OH)2, are weak acids, which with the alkalies yield only 
feebly stable, basic reacting metallic salts. The basic char- 
acter exhibits itself more in the lower stage of oxidation— 
stannous oxide, SnO, and hydrate Sn(OH)2. Hence, we will 
consider tin with the metals. 
Titanium. 
Ti = 48. 
Zirconium. 
Zr = 90. 
Thorium. 
Th = 234. 
Just as vanadium, niobium and tantalum attach themselves 
to the elements of the phosphorus group, so stand, in like rela- 
tion to the silicon group, the three elements, titanium, zirco- 
nium and thorium :— 
P= 31 
As = 75 
Sb = 121 
V = 51 
Nb = 94 
T = 182 
Si = 28 
Sn = 178 
Ti = 48 
Zr= 90 
Th = 234 
In all their deportment, they strongly resemble tin ; they 
possess, however, in their derivatives, a more metallic charac- 
ter. They are tetratomic, forming compounds of the form 
MeX4, in which X represents monatomic elements and groups ; 
those of the form MeX2, corresponding to the stannous deriv- 
atives, are unknown. The hydrates, Me(OH)4 and MeO(OH)2, 
have a stronger basic nature than stannic acid and form 
stable salts with acids; the basicity increases successively 
with the atomic weights, in the order, Ti Zr Tli. Corres- 
ponding to this, the acidity of the hydrates, i. e., their capa- 
bility of exchanging H for metals, gradually diminishes. 
Thorium hydroxide, Th(OH)4, is not able to form metallic 
salts. TITANIUM. 
229 
TITANIUM. 
Ti — 48. 
By ignition the hydrates yield white amorphous 
Occurs in nature as titanium dioxide (rutile, anatase, 
brookite) and in titanates (perofskite Ti03 Ca, Menaccanite 
FeTiOs). Free titanium is a gray, metallic powder, obtained 
by heating potassium fluoride (K2TiFl6) with potassium. It 
burns when heated in the air, and decomposes water on 
boiling. Dissolves in dilute hydrochloric and sulphuric acids, 
with evolution of hydrogen. 
Titanium Chloride—TiCl4—is formed, like silicon chloride, 
by conducting chlorine over a glowing mixture of the dioxide 
and carbon. A colorless liquid, of specific gravity 1.76, fum- 
ing strongly in the air, and boiling at 136°. The vapor density 
equals 95 (H = 1), corresponding to the molecular formula 
TiCl4 = 190. Not known in a free state. Behaves like tin 
tetrachloride with water. A compound Ti2Cl6, analogous to 
C2C16, is known. 
Titanium Fluoride—TiFl4—not known in a free condition ; 
forms beautifully crystallized double salts e. g., TiFl4, 2KF1, 
corresponding to the silico-fluorides (K2SiFl6.) 
Titanic Acid—H4Ti04—separates as a white, amorphous 
powder, on adding ammonium hydrate to the hydrochloric 
acid solution of the titanates. Dried over sulphuric acid it 
loses 1 molecule H20 and becomes TiO, (OH)2. Titanic acid, 
like silicic and stannic, forms polyacids. The hydrates dissolve 
in alkalies and strong acids, to form salts. 
Titanium Dioxide—Ti02, which may be procured crystal- 
lized as rutile, brookite and anatase. Glowed in a stream of 
hydrogen it changes to the oxide Ti203. Titanium dioxide 
is almost insoluble in the acids; only dissolved by hydro- 
fluoric acid. It forms titanates upon fusion with the alkalies. 
With strong acids the hydrates form salts, e. g., TiO. S04. 
Water decomposes this. The alkaline titanates (K2Ti03) 
are very unstable. Other titanates occur in nature e. g., 
CaTiO.,, MgTiOa, and the so-called Titanic Iron, FeTiO.,. 
With nitrogen, titanium yields various compounds. When 
the dioxide is heated in ammonia gas, a dark violet powder of 
the composition TiN2 results. The compound Ti5CN4—the 
so-called cyan-titanium nitride, is sometimes found in copper 
red metallic cubes, in blast furnace slag, when iron ores con- 
taining titanium have been fused. 230 
INORGANIC CHEMISTRY. 
ZIRCONIUM. 
Zr = 90. 
Zirconium occurs only rarely in nature, generally in sili- 
cates, and especially as zircon ZrSi04. Zirconium is obtained 
free in the same way as titanium, and may be isolated as 
an amorphous black powder or in crystalline metallic 
leaflets of specific gravity, 4.15. Zirconium tetrachloride—■ 
ZrCl4, and fluoride, ZrFl4—are very similar to the corres- 
ponding titanium compounds. 
Zirconic Acid or Hydrate—Zr(OH)4—is precipitated by 
ammonium hydrate, from acid solutions, as a white voluminous 
precipitate, which becomes Zr02, zirconium dioxide, upon 
ignition. Zirconic acid is insoluble in potassium and sodium 
hydrates; only on fusion with the alkalies and their carbon- 
ates does it yield zirconates—Na2Zr03 and Na4Zr04, which 
water decomposes. 
The oxide and hydrate dissolve when warmed with 
sulphuric acid, forming Zr(S04)2, which may be crystallized 
from water. 
THORIUM. 
Occurs very rarely, mostly in silicates (Thorite). Free 
thorium, separated by sodium from the chloride, is a dark 
gray powder, of specific gravity 7.7, which in the air burns 
to the dioxide. 
Th = 234. 
Thorium Hydrate—Th(OH)4—and Thorium Dioxide, 
Th02, do not form salts with the alkalies. In sulphuric acid 
they dissolve to the sulphate Th(S04)2, which crystallizes from 
water with four molecules of water. 
BORON. 
B = 11. 
Generally classed with the metalloids, and stands isolated 
among them ; it forms the transition from these to the 
metals, as appears plainly in 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). Like the metals, it does not combine with 
hydrogen, and its oxide B2Os, although really of an acid 231 
nature, approaches such undetermined metallic oxides as 
aluminium oxide, A1203. Boron is triatomic and forms com- 
pounds only of the form BX3. 
BORON TRICHLORIDE 
It is found in nature as boracic acid and in the form of 
borates, like borax (sodium salt), boracite (magnesium salt). 
It may be obtained free in an amorphous and crystallized state. 
The first results upon igniting boron trioxide with sodium, 
away from air contact; free boron and sodium borate are 
formed. On treating the fusion with water, the borate dis- 
solves, leaving the metal as a greenish-brown powder, which, 
when heated in the air, burns with strong brilliancy to the 
trioxide. Nitric and sulphuric acids change it to boric acid. 
Fused with phosphoric acid, it liberates phosphorus. Upon 
boiling Avith aqueous alkalies it dissolves, like beryllium, 
silicon and aluminium, with formation of borates :— 
The crystalline variety may be obtained by igniting boron tri- 
oxide with aluminium. The boron, separated by the aluminium, 
dissolves in 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, 
of specific gravity 2.63, which are more or less colored.* Its 
lustre, refraction of light and hardness, resemble that of the 
diamond. Crystalline boron is more stable than the amor- 
phous ; it does not oxidize upon ignition, and is only slightly 
attacked by acids. Fused with potassium and sodium hydrate 
both modifications yield sodium borate. | */ /' 
2B + 2K0H + 211,0 = 2B0.0K + 3H2. 
Boron Trichloride, BC13, may be prepared by heating boron 
in chlorine or conducting a stream of the latter over agiowing 
mixture of the trioxide and carbon (see SiCl4 and A12C16):— 
It is a colorless liquid of specific gravity 1.35, and boiling 
at 18°. Its vapor density equals 58.7 (H = 1), corresponding 
to the molecular formula BC13 = 117.5. The liquid fumes 
strongly in the air and decomposes with water into boric and 
hydrochloric acids:— 
B303 + 3 C + 3 Cl2 = 2 BC13 + 3 CO. 
BC13 + 3 H20 = B(OH)3 + 3 HC1. 
* According to late investigations, the crystals are not pure boron, 
but contain aluminium and carbon (Hampe). 232 
INORGANIC CHEMISTRY. 
The trichloride also results from the action of the penta- 
chloride of phosphorus upon the trioxide :— 
Boron Fluoride, BF1„, similar to silicon fluoride, and pro- 
duced according to the same methods, by the action of hydro- 
fluoric acid upon the trioxide, or by warming a mixture of the 
trioxide and calcium fluoride with sulphuric acid:— 
B203 + 3 PC15 = 2 BC13 + 3 POCI3 
B203 + 3 CaFl2 + 3 H2S04 = 3 CaSo4 + 3 H20 + 2 BF13. 
Is a colorless gas, burning strongly in the air, of specific 
gravity 34 (H = 1) and may, under strong pressure, be con- 
densed to a liquid. It dissolves extremely readily in water 
(700 volumes in 1 vol.) producing Hydrogen-Boro-fluoride, 
BF14H (= BF13.F1H), which remains in solution :— 
The reaction is analogous to the formation of hydrofluo- 
silic acid from silicon fluoride (see p. 155). Hydrogen boro- 
fluoride is a monobasic acid, only known in solution and in its 
salts. 
4BF1S + 3H20 — 3HBF14 + H3B03. 
Boric Acid—H3B03 = B(0H)3—occurs in salts and free in 
nature. In some volcanic districts, especially in Tuscany, 
steam escapes from the earth (fumaroles, etc.) containing 
small quantities of it. These vapors condense in small 
natural water pools, or are conducted into walled basins. 
By evaporation and concentration of the aqueous solution 
boric acid separates; the same occurs naturally as sassolite. 
To prepare pure boric acid precipitate a hot solution of borax 
with hydrochloric acid. The acid separates in colorless, shin- 
ing scales; it dissolves in 25 parts, H20 of 14°, or in 3 parts 
at 100°. The solution indicates 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. By these reactions we are af- 
forded a ready means for its detection. 
Heated to 100° the acid loses 1 molecule H20 and passes into 
the anhydro or meta acid HB02, which at 140° is converted 
into Tetraboric acid—B407H2. When glowed boric anhy- 
dride or Boron trioxide B203 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. PERIODIC SYSTEM. 
Boric acid is a very weak acid ; it can be expelled from its 
salts by most other acids. When fused, in consequence of 
the difficult volatility of its anhydride, it removes the most 
acids from their salts. 
Salts of normal boric acid B(OH)3 are not known, while 
the ethers B.(O.CH3)3 are. The salts of metaboric acid BO. 
OK can be obtained crystallized, but they are very unstable. 
They are decomposed by carbon dioxide with production of 
salts of tetraboric acid :— 
The latter, from which the ordinary borates are derived 
(see borax) may be viewed as an anhydro-acid, formed by 
the union of 4 molecules of trihydric boric acid (compare 
p. 186):— 
4NaB02 + C02 = B407Na2 + C03Na2. 
On heating amorphous boron in a stream of N or ammonia, 
or by glowing a mixture of trioxide and carbon in nitrogen 
gas, there is formed boron nitride, BN. This is a white 
amorphous powder, which, heated in a flame, gives forth an 
extremely intense greenish-white light. On passing steam 
at 200° over this boric acid and ammonia are produced:— 
4B (0H)3 — 5H20 = B407H2. 
Bn + 3H20 = B (OII)3 + NH3. 
PERIODIC SYSTEM OF THE ELEMENTS. 
In the preceding pages we have studied four groups of ele- 
ments and their compounds with hydrogen, the halogens and 
oxygen. The remarkable relations of the elements of any one 
group, as, also, of the various groups to each other, to which 
we have repeatedly directed attention, appear more manifest 
if viewed in the same connection in which they present them- 
selves in the periodic system of elements. The position which 
these elements occupy in the system determines, up to a 
marked degree, their entire physical and chemical character. 
The system is based upon the grouping of the elements 
according to the magnitude of their atomic weights. For the 
longest time we have been cognizant of the remarkable re- 
lations existing between the atomic weights of analogous 
elements, but only recently has the law of periodicity under- 
lying them been affirmed by Medelejeff and Lothar Meyer, 234 
INORGANIC CHEMISTRY. 
and, according to this, the properties of the elements and their 
compounds present themselves as periodic functions of the 
atomic weights. 
Arranging the elements according to increasing atomic 
weight we observe that similar elements return after definite 
intervals. Thus they arrange themselves in several periods, 
consisting of the following horizontal series (for brevity the 
atomic weights are not attached to the symbols) :— 
1. Li Be B C N 0 FI 
2. Na Mg A1 Si P S Cl 
3) K Ca Sc Ti V Cr Mn ; Fe Co Ni ; Cu Zn Ga — As Se Br 
4 Rb Sr Y Zr Nb ;Mo — j Ru Rh Pd j Ag Cd In Sn Sb Te I 
J 5) Cs Ba (La Ce Di) — — ; — — — • — — — — 
16) Yb — Ta W — j Os Ir Pt j Au Ilg Tl Pb Bi 
The first series, lithium (Li) to fluorine (FI), and sodium 
vNa) to chlorine (Cl), present two periods of seven members 
each, in which the corresponding (above and below) mem- 
bers exhibit a great but not complete analogy. Sodium 
resembles lithium, magnesium, beryllium, chlorine, fluorine, 
etc. Then follow two periods, consisting of 17 elements each: 
Potassium (K) to bromine (Br), and rubidium (Rb) to iodine 
(I). The series 5 and 6 are incomplete, and together constitute, 
as it appears, a period. In the 7th series there are as yet but 
two elements: thorium = 234 and uranium = 240. Thus result 
3 great periods, whose corresponding members exhibit an almost 
complete analogy ; the elements K Rb Cs, Ca Sr Ba, Ga In 
Tl, As Sb )3i, etc., are so similar that they remind us of the 
homologous series of the carbon compounds (compare p. 144), 
and, therefore, can be designated as homologous elements. 
Only in the third great period (series 5 and 6) do the middle 
members exhibit any variations. 
On now 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 Li; Ca Sr Ba 
resemble Mg and Be. Then the similarity gradually lessens, 
disappears apparently in the middle members, and only 
again appears toward the end of the periods: I and Br 
resemble chlorine and fluorine, Te and Se sulphur and 
oxygen, Bi Sb As phosphorus and nitrogen, etc. The 
character or the function of the three great periods is therefore PERIODIC SYSTEM. 
235 
other than that of the two small periods. But in all five 
periods a gradual, regular alteration in the properties of the 
adjoining heterologous elements exhibits itself. Particularly 
distinct is this in the measurable physical properties, all of 
which, in the middle of the periods (both of the large and 
small), show a maximum or minimum, as is visible in the 
specific gravity of the elements in solid condition (compare 
further the atomic volumes p. 245). 
Specific gravity, 
Na 
0.97 
Mg 
1.7 
A1 
2.5 
Si 
2.5 
P 
2.0 
S 
1.9 
Cl 
1.3 
K 
0.86 
Ca 
1.6 
Sc 
Ti 
V 
5.5 
Cr 
6.8 
Mn 
7.2 
Fe 
7.9 
Co 
8.5 
Ni 
8.8 
Cu 
8.8 
Zn 
7.1 
Ga 
5.9 
As 
5.6 
Se 
4.6 
Br 
2.9 
These relations show themselves very clearly in a graphic 
representation, by making the atomic weights the abscissae, 
the value in numbers of the properties the ordinates; then 
the individual periods represent segments of curves, which 
blend to a curve with alternating maxima and minima. 
The same regularity exhibits itself with the 2 small periods, 
even in chemical properties, especially in the valence of the 
elements in their compounds with hydrogen or the hydro- 
carbon groups CH3, C2 H5, etc., (compare p. 166 and p. 239). 
The hydrogen valence rises and falls periodically with the 
condensation of the substance (corresponding to the specific 
gravity) 
NuR 
MgR2 
III 
A1R3 
IV 
SiH4 
III 
ph3 
11 
sh2 
C1H, 
On the other hand, the maximum valence of the elements 
in the salt-forming oxides increases successively:— 
Na20 
MgO 
III 
A1203 
IV 
Si04 
p205 
VI 
so3 
VII 
C1207 
The chemical valence expresses itself somewhat differently 
in the three great periods. In them we have a double period- 
icity; thus, e. g., with the salt-forming oxides :— 
K20, 
11 
CaO, 
III 
Sc203 
IV 
Ti02 
v2o5- 
VI 
-Cr03 
VII 
Mn207 
VI 
Fe03 
IT II 
Co02 NiO 
Co20 
11 
ZnO 
III 
Ga2 0 3 
IV 
v 
As203 
VI 
Se03 
VII 
Br207. 236 
INORGANIC CHEMISTRY. 
In consequence of this double periodicity, the first seven and 
the last seven members of the 3 great groups, as regards their 
valence (and consequently also their compounds), resemble 
the 7 members of the 2 small periods. To bring out this 
double periodicity and analogy, the 7 first and 7 last 
members of the great groups are divided into two series and 
arranged under the corresponding 7 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. 
234), as they have no analogues, come to stand apart. In 
this manner arises the following table, in wrhich the 7 (or 10) 
vertical columns include analogous elements:— 
Li 
Be 
B 
C 
N 
0 
FI 
Na 
Mg 
A1 * 
Si 
P 
S 
Cl 
K 
Ca 
Sc 
Ti 
V 
Cr • 
Mn 
Fe Co Ni 
Cu 
Zn 
Ga 
— 
As 
Se 
Br 
Rb 
Sr 
Y 
Zr 
Nb 
Mo 
— 
Ru Rh Pd 
Ag 
Cd 
In 
Su 
Sb 
Te 
I 
Cs 
Ba 
(LaCeDi) 
— 
— 
— 
— 
— 
Yb 
— 
Ta 
W 
— 
Os Ir Pt 
Au 
Hg 
T1 
Pb 
Bi 
— 
— 
Exactly the same grouping of the elements simultaneously 
with their atomic weights, in round numbers, is produced in the 
adjoining table (p. 237). In this arrangement we must always 
bring into consideration that the principal analogy (homology) 
of the three great periods finds expression in the three unin- 
terrupted horizontal series (p. 234), and that the decomposition 
of the latter into every two series only corresponds to the 
secondary, double analogy with the small periods. It may be 
further remarked that in the second small period the three 
last numbers, P, Sand Cl, show a complete homology with the 
corresponding numbers of the large periods—as is expressed 
in the table. PERIODIC SYSTEM 
237 
i 
Group. 
II 
Group. 
Ill 
Group. 
IV 
Group. 
V 
Group. 
VI 
Group. 
VII 
Group. 
VIII 
Group. 
H-COM POUNDS. 
Highest Salt- 
— 
— 
— 
mh4 
mh3 
mh2 
MH 
(M,H) 
forming Oxides. 
M20 
MO 
m203 
mo2 
m205 
mo3 
m2o7 
mo3 
mo2 
MO 
Periods. Series. 
H 1 
1st 1st 
Li 7 
Be 9 
B 11 
C 12 
N 14 
O 16 
FI 19 
2d 2d 
Na 23 
Mg 24 
A1 27 
Si 28 
P 31 
S 32 
Cl 35 
3d { 3d 
K 39 
Ca 40 
Sc 45 
Ti 48 
V 51 
Cr 52 
Mn 55 
Fe 56 
Co 58 
Ni 58 
Cu 63 
Zn 65 
Ga 70 
—73 
As 75 
Se 79 
Br 80 
4th { 6th 
Rb 85 
Sr 87 
Y 89 
Zr 90 
Nb 94 
Mo 96 
—100 
Ru 104 
Rh 104 Pd 106 
l 6th 
Ag 108 
Cd 11 2 
In 113 
Sn 118 
Sb 122 
Te 126 
I 127 
f 7th 
Cs 132 
Ba 137 
(Lal39 Ce 140 Dil44) 
— 
8th 
— 
_ 
6th l 
1 9th 
— 
— 
Yb 173 
— 
Ta 182 
W 184 
— 
Os (198) 
Ir 193 
Pt 196 
Uoth 
Au 197 
Hg 200 
T1 204 
Pb 206 
Bi 210 
— 
— 
— 
— 
— 
(Th 234) 
Ur 240 
ELEMENTS OF THE PERIODIC SYSTEM. 238 
INORGANIC CHEMISTRY. 
When the periodic grouping of the elements was first presented, 
some, at that time not sufficiently well established, atomic weights had 
to be more or less altered. Thus, the atomic weight of iridium, for- 
merly 75.8, was made 118, and that of uranium 240 (before 120). All 
such alterations, through recent investigations, have been proven to 
be established. Further, the atomic weight of tellurium (formerly 
determined to be 128) had to be less than that of iodine (126.6) ; 
this also, has been established by recent researches, which place it 
at 126.8. There is, therefore, no doubt that the atomic weight of 03- 
mium (found 198) will also prove to be somewhat less,—only the more, 
because at present a smaller atomic weight (192.7) than formerly 
accepted (p. 197) has been discovered for iridium. Hence, the periodic 
system offers a control for the numbers of the atomic weight, while 
formerly they appeared to be irregular and, at the same time, acci- 
dental. 
Further, upon the basis of the periodic system, the existence 
of new, not yet known, elements may be ascertained, which 
would correspond to yet unoccupied, free places, or vacancies 
in the table. In fact, two such formerly vacant places have 
been tilled up by the discovery of gallium (Ga = 69.8) and 
scandium (Sc == 45) ; their properties have shown themselves 
to be perfectly accordant with those deduced from the periodic 
system. At present, of the elements of the first four periods only 
two are wanting (see p. 234) : the first homologue of manga- 
nese (with atomic weight of about 100), and the lowest homo- 
logue of tin (atomic weight, about 73). The vacancies in the 
series 5 and 6 are partly explained from the somewhat vary- 
ing function of the third great period (p. 234) ; partly also by 
the most recently discovered or characterized elements, erbium, 
terbium, wasium, and norwegium, whose atomic weights have 
not yet been determined with certainty. These would prob- 
ably find a position here. 
The entire character of a given element is determined to a 
very high degree by the law of periodicity; hence, all physi- 
cal 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 elements, and 
confine ourselves here to a notice of some general relations; 
further, the connection of atomic weight with the valence and 
the thermo-chemical phenomena. 
Particularly distinct does the relation of metalloids to metals 
show itself in the periodic system. The first members of all 
periods (on the left side) consist of electro-positive metals, 
forming the strongest bases, the alkalies—Cs III) K Na Li PERIODIC SYSTEM. 
239 
and metals of the alkaline earths—Ba Sr Ca Mg and Be. 
The basic character diminishes successively, in the following 
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 com- 
prises 5 metalloids, (B C N O FI), the second only four 
(Si P S Cl), the fourth and fifth periods each only 3 (or 2) 
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 derivatives are also afforded by the metalloids 
with the monatomic hydrocarbon groups (as CH;), CaH5, C3 
H7, etc.), which, in many respects, resemble hydrogen. 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 :— 
Na(CH3), 
Mg(CH3)2, 
III 
Al(CH3)3. 
C1CH3. 
IV 
Si(CH3)4, 
III 
P(CH3)3, 
S(CHS)2, 
Their stability gradually diminishes with the increasing basic 
nature of the metals ; hence, in the three large periods, this 
power extends only to Zn, Cd and Hg (beginning with bro- 
mine, I and Bi). 
In consequence of the opposite (metalloidal and metallic) 
character of the two ends of the periods, there are in the 
table representing the double periodicity of the great periods 
(p. 236 and 237) two sub-groups each, with the seven vertical 
groups; on the left with the more positive basic, and on the 
right 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 VII form the gradual transition from the 
last to the first. 
The fundamental deduction resulting necessarily from the 
law of periodicity is, that the various elementary atoms must 
be aggregates or condensations of one and the same original 
substance, a necessary correlative postulate of the recognized 240 
INORGANIC CHEMISTRY. 
unity of all forces. Only then will it be understood that the 
properties of the elements are functions of the atomic weight. 
That this original matter is not hydrogen, as supposed (hypo- 
thesis of Prout), follows from the fact that the atomic weights 
of the elements are not multiples of that of hydrogen ; this is 
shown by the determinations, made with such accuracy by 
Stas, for some of the elements. These atomic weights referred 
to hydrogen = 1 are the following:— 
Oxygen 0 = 15.960 
Chlorine Cl = 35 360 
Bromine Br = 79.750 
Iodine I = 126.533 
Silver Ag — 107.660 
Potassium K = 39.040 
Sodium Na — 22.980 
The ordinary atomic weights as they appear in this book are 
given in round numbers, which is sufficient for general 
purposes. If the atomic weights of all the elements had 
been fixed with equal accuracy, it would probably be possible 
to approach more closely the consideration of the nature of 
the periodic function. 
Periodicity of Chemical Valence.—Group I of the table 
comprises the monatomic metals, group II the diatomic. In 
group III is the triatomic metalloid, boron (which does not 
furnish a hydrogen compound), and the triatomic metals A1 Sc 
Y and Ga In Tl. In the tetratomic group of carbon the quan- 
tivalence arrives at its maximum ; from here the valence 
gradually decreases with increasing atomic vreight; the group 
of nitrogen is triatomic, the group of oxygen is diatomic, that 
of the halogens monatomic. This valence is derived from the 
compounds with hydrogen and hydrocarbons (compare p. 
239), or where such do not exist, as in case of boron and 
many metals, from the halogen compounds :— 
IV 
ch4 
III 
NHS 
II 
0H2 
F1H 
LiCl 
NaCl 
Be Cl 2 
MgCl2 
III 
BC1S 
A1C13 
iy 
cci4 
SiCl4 
hi 
nci3 
pci3 
oci, 
sci2 
FI* 
cr2 
The elements of the 4 first groups are not capable of yield- 
ing 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 
p. 163). The higher valence of these elements shows itself yet 
plainer in the more stable oxygen compounds. On arrang- PERIODIC SYSTEM. 
241 
ing together the highest oxides of the seven groups capable of 
forming salts (salt-building oxides), we get this series:— 
1 
Li20 
II 
BeO 
III 
b203 
IV 
co2 
n205 
VI 
so3 
VII 
IaO,. 
The elements of the four first groups in their oxygen com- 
pounds of course exhibit the same valence as in the compounds 
with hydrogen (or hydrocarbon radicals) and the halogens ; 
in the three last series, however, there is noticed a constant 
increase of valence up to oxygen. 
Besides the highest oxides, remarkable for their greater sta- 
bility, the elements of the three last groups form lower oxides, 
as in this manner they return to hydrogen valence:— 
III 
P203 
IV 
so2 
I2<v 
II 
SC12 
Ill 
C1203. 
C120. 
The hydroxyl compounds of the elements of the 7 groups 
are analogously constituted to the oxides. With them, we 
have the following limiting series, as expression of the maxi- 
mum valence (compare p. 165) :— 
I 
Na(OH) 
Mg(OH)a 
III 
Al(OH), 
IV 
Si(OH)4 
P(OH)5 
S(OH)6 
VII 
Cl(OH)7 
The hydroxyl compounds of the elements of the 4 first 
groups exist in free condition, excepting that of carbon, 
C(OH)4, which is only represented in its derivatives. The 
strong basic character of the hydrates of group I (NaOH) 
diminishes, step by step, down to the weak acid hydrate, 
Si(OH)4. The hydrates of the three last groups are of acid 
nature, and mostly unstable or not known. 
By elimination of 1, 2 and 3 molecules H20 they yield the 
ordinary highest acids:— 
PO(OH)3 
Phosphoric Acid. 
S02(0H)2 
Sulphuric Acid. 
VII 
C1030H 
Perchloric Acid. 
In the same way behave the non-saturated hydrates 
III 
P(OH), 
S(0H)4 
Cl(OH)5 
S(OH)2 
III 
Cl(OH), 
Cl(OH) 242 
INORGANIC CHEMISTRY. 
From the hydrate, S(OH)4 is derived sulphurous acid, 
SO(OH)2 (p. 179); from the hydrate, Cl(OH)5, chloric acid, 
C102,OH ; from the hydrate, Cl(OH)3, chlorous acid, CIO,OH. 
The hydrates P(OH)3, S(OH)2 and ClOH are very unstable, 
and the two first appear to readily pass into H2PO,OH and 
HSO,OH (compare p. 180). 
It has already been shown, in the case of per-iodic, sulphuric 
and nitric acids, how the usually, so-called hydrates with 
water of crystallization (regarded as molecular compounds) 
find explanation by the acceptance of the existence of such 
hydroxyl compounds. The same may be carried out 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 quantivalence of 
the elements have their complete expression in the periodic 
system, are regulated by it, and hence we must conclude that, 
in fact, the valence is not only a property attaching to the 
elements per se, but is also influenced by the nature of the 
combining elements; the hydrogen valence is constant, the 
valence to oxygen and halogens, on the contrary, is variable, 
according to definite rules. Valence, therefore, is a relative 
function of the elements. 
Periodicity of Thermo-Chemical Phenomena. We ob- 
served in the ease of the elements of the chlorine and oxygen group, 
that, in their hydrogen compounds the heat liberation decreased success- 
ively with increasing atomic weight (p. 57), while, with the chlorine and 
oxygen derivatives (corresponding to the growing metallic character), 
generally an increase occurs. Similar relations exhibit themselves 
with the halogen, oxygen and sulphur compounds of the metals, as 
will later be more closely exemplified with the individual groups. 
Here it is sufficient to call attention to the relations in the heterolo- 
gous series. 
In the formation of the hydrogen compounds of the elements of the 
two first periods, according to present data, the following quantities of 
heat are set free : — 
(C,H4) 
19.5 
(N,H3) 
11.9 
(0,H2) 
67.2 
(F^H) 
(Si,H4) 
33.2 
(P,H3) 
36.5 
(S,Ha) 
4.5 
(01, H 
22.0*) 
* The numbers given here for the heat intensity, as well as those fol- 
lowing, indicate so-called great calories, which contain 1000 ordinary 
or small calories; therefore, to convert them into the latter, multiply 
them by 1000. PERIODIC SYSTEM 
243 
In the halogen compounds the heat is more regular:— 
Li, Cl 
93.8 
(Be,Cl2) 
(B,C13) 
104 
(C,C14) 
(N,cis) 
—38.1 
(0,C12) 
—18 
(FI,Cl) 
(Na,Cl) 
97.7 
(Mg,Cla) 
151.0 
(A1C1,) 
160.9 
(SiCl4) 
157.6 
(P,C1.) 
68.9 
(S,C12) 
(Cl,Cl) 
With the bromides, it is less throughout, and the least with the 
iodides. Consequently, a maximum appears to lie in the middle of 
the periods. On calculating, however, the heat intensity, which corres- 
ponds to 1 equivalent of the elements (united with 1 equivalent of 
chlorine) we obtain successively decreasing numbers, corresponding to 
the decrease of the basic-metallic character of the elements :— 
(Na,Cl) 
97.7 
75.5 
53.6 
(S^CU) 
39.2 
(¥0 
23 
(s40 
(Cl,Cl). 
Perfectly similar relations are furnished by the oxides:— 
(Na20) 
100.2 
(Mg,0) 
145.8 
(AlaO.) 
388.8 
(Si,Oa) 
219 
(P205) 
363 
(S,03) 
104 
(C1207). 
Calculated upon 1 equivalent, the heat intensity is: 
50.1 72.9 64.6 54.7 36.3 17.3 
That the heat of combination is smaller with sodium oxide than with 
magnesium oxide, depends partly upon the solubility of the first, as 
this property is also to be included as a thermal function. A like 
diminution of heat intensity shows itself also with the heterologous 
elements in their compounds of similar form :— 
(Mn, Cl2) 
111.9 
(Mn, 0) 
94.7 
(Fe, Cla) 
82 
(Fe, 0) 
68.2 
(Co, Cl2) 
76.4 
(Co, O) 
63.4 
Ni, Cla) 
74.5 
(Ni, 0) 
60.8 
(Cu, Cla) 
51.6 
(Cu, 0) 
40.8 
(Zn, Cla) 
97.2 
(Zn, 0) 
85.4 
Remarkable are also the following series 
Ag, Cl, 
29 3 
Cd, Cla, 
93.2 
In, Cl3, 
Sn, Cl4 
127.2 
Sb, Cl s 
87 
Te, Cla 
(Ag20) 
5.9 
(Cd, 0) 
63.6 
(In2, 0») 
(Sn, 02) 
133.5 
(Sba, 03) 
(Te,0) THE METALS. 
Although there is no sharp line of demarkation between metals 
and non-metals, yet these two classes of bodies form, in their 
entire deportment, a distinct contradiction, as appears plainly 
in the periodic system of elements. In physical respects the 
character of metals determines itself by the external appear- 
ance and by the conductivity of heat and electricity; chemi- 
cally, 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 diminishes gradually 
and passes over into an acid-forming one. 
At ordinary temperatures all the metals excepting mercury 
are solid, slightly volatile bodies. They are opaque, and only 
some, like gold, when beaten into thin leaflets, permit the pass- 
age of light to a limited extent. In compact mass they exhibit 
metallic lustre and possess mostly a whitish-gray color; only 
gold and copper are brilliantly colored. In powder form 
almost all the metals are black. The most crystallize in 
forms of the regular system; only a few, showing a metal- 
loidal character, are not regular. Thus, antimony and bis- 
muth crystallize hexagonally, and tin is quadratic. The 
specific gravity of the metals is very different, and varies from 
0.59 to 22.4 as seen from the following arrangement — 
PHYSICAL PROPERTIES OF THE METALS. 
Lithium, 0.59 
Potassium, 0.86 
Sodium, 0.97 
Rubidium, 1.52 
Calcium, 1.58 
Magnesium, 1.75 
Aluminium, 2.56 
Chromium, 5.9 
Arsenic, 5.9 
Antimony, 6.7 
Zinc, 7.1 
Tin, 7.3 
Iron, 7.8 
Cobalt, 8.5 
Cadmium, 8.6 
Copper, 8.9 
Bismuth, 9.8 
Silver, 10.5 . 
Lead, 11.3 
Palladium, 11.8 
Thallium, 11.9 
Mercury, 13.5 
Gold, 19.3 
Platinum, 21.4 
Iridium, 22.3 
Osmium, 22.4 
In general the specific gravity of the metals, as also of the 
metalloids, increases with the atomic weights; more especially 
do they stand in a sharp periodic dependence with reference 
to the latter. The first members of all periods possess low PHYSICAL PROPERTIES 
245 
specific gravities; the latter grow gradually until the middle of 
the period, where the maximum is attained, and then they 
again decrease (p. 235). These relations show themselves more 
fully if, instead of the specific gravity, we compare the specific 
volumes or atomic volumes; i. e., the quotients from the 
atomic weights (A) and specific gravities (d) :— 
A 
= specific volume. 
These quotients express the relative volumes of the atoms 
(in solid or liquid state). Thus the atomic volume of lithium 
_T- = 11.9, that of potassium = 45.4: i. e., the potassium 
atom occupies a 3.8 times larger space than that of the lithium 
atom. The periodic alterations of the atomic volumes are set 
opposite to those of the specific gravities, as the former are 
obtained by the division of the atomic weights by the specific 
gravity. Therefore the atomic volumes decrease gradually, 
commencing from the first members of the periods (Li Na 
K Rb), attain a minimum in the middle of the periods and 
then increase again until to the last members (Cl Br I). On 
the other hand, with the homologous elements (the vertical 
series) with increase of atomic weight there is almost always 
seen an increase of the atomic volumes. 
Since in the 3 large periods the alterations of the atomic 
volumes (as also of all other physical properties) indicate a 
simple periodicity (not double, like the valence), they are 
expressed in the following tables, by progressive series (page 
234):— 
ATOMIC VOLUMES OF THE ELEMENTS 
Li 
11.9 
Be 
5.7 
B 
4.1 
C*) 
3.6 
N 
Of 
17 
FI 
Na 
23.7 
Mg 
13.8 
Al 
10.7 
Si 
11.2 
P 
13.5 
S 
15.7 
Clt) 
25.6 
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga — As Se Br 
45.4 25.4 9.3 7.7 6.9 7.2 7.0 6.7 7.2 9.7 11.6 — 13.2 16.9 26.9 
Rb Sr Y Zr Nb Mo — Ru Rh Pd Ag Cd In Sn Sb Te 
56.1 34.9 — 21.7 15.0 11.1 — 9 2 8.6 9.2 10.2 12.9 15.3 16.1 18.2 20.5 
I Cs Ba(La CeDi) — — — — — 
25.6 — 26.5 22 20 22 
Yb — Ta W — Os Ir Pt Au Hg Tl Pb Bi — — 
— 16.9 9.6 — 9.4 8.7 9.3 10.2 18.7 17.1 18.1 21.1 
Th — Ur 
30 13 
* As diamond. 
f Liquid 246 
INORGANIC CHEMISTRY. 
It is exceedingly noteworthy, that the elements standing at the 
beginning and end of the periods (one side the alkalies Li Na K Rb, 
and the alkaline earth metals Be Mg Ca Sr Ba—on the other, the 
halogens and the elements of the oxygen group), possess the greatest 
chemical energy, and there is scarcely a doubt that between chemical 
energy and atomic volume there exists a closer causal connection. 
We can suppose that the specifically light elements, with large atomic 
volumes, execute larger chemical oscillations, hence act together more 
readily and energetically. With this might be included, that, in en- 
ergetic reactions, greater quantities of heat are eliminated. Further, 
might be deduced, that, between the elements of greater but oppo- 
site oscillations (the alkalies and halogens) the expressions of chemi- 
cal valence are the simplest; that towards each other they deport them- 
selves as monatomic. 
The metals whose specific gravity is less than 5, are termed 
light, the rest heavy. The former have, in general, a greater 
chemical energy, therefore oxidize more easily and form 
strong basic oxides; their compounds are mostly readily 
soluble. The heavy metals possess, on the whole, a varying 
deportment. They are less energetic, less basic, and yield 
insoluble oxygen and sulphur derivatives; their naturally 
occurring compounds usually possess metallic lustre and are 
termed ores. 
Most metals are very malleable and tenacious; hence oan 
be beaten into thin plates and leaves and drawn out into 
wires ; most malleable are gold and silver. Only a few, like 
antimony, bismuth and tin, which also possess a metalloidal 
character, are brittle and may be pulverized. Heat will fuse 
all metals, although some require the high temperature of 
the oxy-hydrogen flame. 
The fusing points of the most important of them are the 
following:— 
Mercury, —39° 
Rubidium, -f-38° 
Potassium, 62° 
Sodium, 95° 
Tin, 228° 
Bismuth, 270° 
Cadmium, 315° 
Lead. 334° 
Zinc, 423° 
Aluminium, 750° 
Silver, 954° 
Gold. 1035° 
Copper, 1054° 
Cast Iron, 1150° 
Wrought Iron, 1500° 
Palladium, 1500° 
Platinum, 1779° 
Iridium, 1950° 
A greater volatility also corresponds with the greater fusi- 
bility. Mercury boils at 360° ; potassium and sodium about 
440°; cadmium at 860° ; zinc towards 1000°. The diffi- SPECIFIC HEAT. 
247 
cultly fusible metals may also be volatilized by the* galvanic 
current. 
All these physical properties bear a periodic dependence 
to the atomic weights, as will be more closely indicated ir 
the individual groups. 
SPECIFIC HEAT—ATOMIC HEAT 
Of all physical properties of the elements, from a chemical 
standpoint, their heat capacity is the most important, and can 
serve for the determination of the atomic weights. To heat 
one and the same quantity, by weight, of the different metals or 
substances to one and the same temperature, a very varying 
amount of heat is required. This is evident from the follow- 
ing experiments:— 
If we add to 1 kilogram of H20 at 0° 1 kilo, of H20 at 100°, 
the temperature of the mixture of 2 kilograms is 50°. The 
quantity of heat necessary to raise 1 part by weight of H20 1° 
is, for all temperatures from 0-100° almost the same; this is 
designated the heat unit or calorie. On bringing to 1 kilo- 
gram H20 at 0° 1 kilogram Hg at 100°, the temperature of 
the water and of the mercury after their compensation will 
equal only 3.2°. Consequently the mercury has, as it cooled 
about 96.8° (from 100 to 3.2°) given off 3.2 calories. The 
quantities of heat contained in equal parts by weight of water 
and mercury, therefore, deport themselves as 96.8 to 3.2, i. e., 
the specific heat of mercury (that of water being made = 1) 
is A* = 0.0332. 
96.8 
On comparing the specific heats of solid elements found 
in this way with their atomic weights, we discover that these 
are inversely proportional to the latter, and hence the product 
of the specific heat and atomic weight for all solid elements 
(few excepted), is a constant quantity. This fact was first 
discovered by Dulong and Petit (1819) and formulated in the 
following law : The solid elements possess the same atomic heat. 
In the subjoined table are presented the specific heats of the 
elements in solid condition (as far as they have yet been 
determined). W represents the specific heat, A the atomic 
weight, and the product, W X A, the atomic heat:— Elements. 
w | 
A 
w X A 
Hydrogen* 
5,880 
1 
5,6 
Lithium 
Li 
0,941 
7 
6,6 
Beryllium 
Be 
0,408 
9,3 
3,8 
Boron (amorphous) 
B 
0,254 
11 
2,8 
Graphite 1 
n 
0,174 
}12 
2, 1 
Diamond / 
0,143 
1,7 
Sodium 
Na 
0,293 
23 
6,7 
Magnesium 
Mg 
0,245 
24 
5,9 
Aluminium 
Al 
0,202 
27,3 
5,5 
Silicon (cryst) 
Si 
0,165 
28 
4,6 
.Phosphorus (yellow) 
P 
0,189 
31 
5,9 
Sulphur (rhombic) 
S 
0,178 
32 
5,7 
Potassium 
K 
0,166 
39 
6,5 
Calcium 
Ca 
0,170 
40 
6,8 
Chromium 
Cr 
0,100 
52,4 
5,2 
Manganese 
Mn 
0,122 
54,8 
6,7 
Iron 
Fe 
0,112 
56 
6,3 
Cobalt 
Co 
0,107 
58,6 
6,3 
Nickel 
Ni 
0,108 
58,6 
6,4- 
Copper 
Cu 
0,093 
63,3 
5,9 
Zinc 
Zn 
0,093 
65 
6,1 
Gallium 
Ga 
0,079 
69,9 
5,5 
Arsenic (cryst) 
As 
0,082 
75 
6,2 
Selenium (cryst) 
Se 
0,080 
79 
6,4 
Bromine (solid) 
Br 
0,084 
80 
6,7 
Zirconium 
Zr 
0,066 
90 
6,0 
Molybdenum 
Mo 
0,072 
96 
6,9 
Ruthenium 
Ru 
0,061 
103,4 
6,3 
Rhodium 
Rh 
0,058 
104 
6,0 
Palladium 
Pd 
0,059 
106,3 
6,3 
Silver 
Ag 
0,056 
107,6 
6,0 
Cadmium 
Cd 
0,054 
111,6 
6,0 
Indium 
In 
0,057 
113,4 
6,5 
Tin 
Sn 
0,054 
117,8 
6,5 
Antimony 
Sb 
0,052 
122 
6,4 
Tellurium 
Te 
0,047 
126,8 
6,0 
Iodine 
I 
0,054 
126,5 
6,8 
Lanthanum 
La 
0,045 
139 
6,2 
Cerium 
Ce 
0,045 
140 
6,2 
Didymium 
Di 
0,045 
144 
6,5 
Tungsten 
W 
0,033 
184 
6,1 
Osmium 
Os 
0,031 
(198) 
6,2 
Iridium 
Ir 
0,032 
192 7 
6,3 
Platinum 
Pt 
0,032 
196,7 
6,4 
Gold 
Au 
0,032 
196,2 
6,4 
Mercury (solid) 
Hg 
0,032 
200 
6,4 
Thallium 
T1 
0,033 
204 
6,8 
Lead 
Pb 
0,031 
206,4 
6,5 
Bismuth 
Bi 
0,030 
210 
6,5 
* As Palladium hydride. SPECIFIC HEAT. 
249 
From the table, it is evident that the atomic heats of the 
most elements lie between 5.9 and 6.8, and equal, upon an 
average, 6.4. 
Only in the case of few elements is the atomic heat some- 
what less (S, P, Si, Al) or considerably less (C, B, Be), than 
the mean. Such are those with low atomic weight, "having 
a metalloidal character and occupying the middle of the two 
small periods. These variations bear distinct periodic depend- 
ence to the atomic weights:— 
Na 
6.7 
Be 
3.8 
B 
2.5 
C 
1.9 
N 
0 
FI 
K 
6.5 
Mg 
6.0 
A1 
6.5 
Si 
4.6 
P 
5.6 
s 
5.7 
Cl 
The variations in the middle are in part explained by the 
fact that most of the elements in their different modifications 
(crystalline, amorphous, malleable) possess a somewhat dif- 
ferent heat capacity, as observed with carbon. More important, 
however, is the influence of temperature. The figures in the 
table indicate mostly the heat capacities at medium tempera- 
tures. It was known before that these show a slight increase 
with the temperature, but only recently, through H. E. Weber, 
has it been proven that, for the elements C, B and Ti, which, 
at medium temperatures, possess a remarkably low atomic 
heat, the increase is very considerable; that, beyond a definite 
temperature, the atomic heat becomes tolerably constant, and 
then almost agrees with the law of Dulong and Petit:— 
w. 
A. 
wx A. 
Diamond, graphite, above 600°, 
0.45 
12 
5.4 
Boron, above 600°, 
0.5 
11 
5.5 
Silicon, above 200°, 
0.2 
28 
5.6 
It is probable that, for all elements, a definite temperature 
exists, at which their heat capacities become accurately com- 
parable. 
From this close agreement of the found atomic heat of the 
metals with the medium, follows, without doubt, the occurrence 
of a regularity, and we must conclude that the slight varia- 
tion?, apart from the inaccuracy of the observation?, are influ- 
enced by secondary causes. Hence the specific heat may 250 
INORGANIC CHEMISTRY. 
serve for the derivation of the atomic weight of metals; the 
atomic weight is equal to the constant quantity, 6.4 divided by 
the found specific heat:— 
A =6.4 
W 
The atomic weights derived from the specific heat—the so- 
called thermal atomic weights, agree almost in 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. By analysis the 
equivalent weight 37.8 (InCl) of indium is fixed with great 
accuracy ; it is, however, unknown whether the atomic weight 
is double or triple that quantity. The specific heat of indium 
is 0.0569, from which the atomic weight would be 64 = 
112.5—a number closely approaching three times the equiva- 
lent weight of indium 113.4 (=37.8x3). From this it follows 
that the true atomic weight of indium is 113.4 and that it is 
triatomic, InCl3. 
In their solid compounds the elements retain the specific heat attach- 
ing to them in their free, solid state ; hence the molecular heat is 
equal to the sum of the atomic heats of the elements constituting the 
molecules—law of Neumann and H. Kopp. Hence the atomic heat 
of elements not known in solid condition may be derived from the 
molecular heat of their compounds. Thus are found the following 
atomic heats: for nitrogen, 6.4; for chlorine, 6.4; for oxygen, 4; for 
fluorine, 5 ; for hydrogen, 2.3. 
In the free gaseous state the elements have mostly a slighter atomic 
heat, as seen from the following table:— 
A 
Wf 
A X W 
Oxygen 
16 
0.156 
2.5 
Hydrogen.. 
6 
2.405 
2.4 
Nitrogen... 
14 
0.172 
2.4 
Chlorine.... 
85.4 
0.093 
3.3 
* It is only for iron that the atomic weight derived from its few vola- 
tile compounds is twice as large (112) as that resulting from the heat 
capacity—of free iron, hence we must conclude that in these com- 
pounds two iron atoms are present. 
f By constant volume. ISOMORPHISM. 
251 
The law, that the atoms in solid condition possess the 
same heat capacity (A. W = A.'W.'), finds an interesting 
analogy and exemplification in the result coming from the 
mechanical gas theory, and in the proposition of Avogadro, 
that the molecules in gas condition, at like temperatures, have 
similar degree of motion (M. v. = M\ v') and experience 
like increases. As in the gas condition the molecules, so for 
the solid the atoms, are the smallest mass particles, which 
have the same heat energy. The velocity of their heat motion, 
both for the molecules and for the atoms, is so much the 
greater, the smaller their masses. 
ISOMORPHISM 
As indicated in the preceding pages, the atomic weights of 
the elements may be directly derived from the heat capacity 
of solids, while from the gas density of the volatile compounds 
the molecular weights, and from the latter, then, indirectly, are 
ascertained the atomic weights (compare p ). A third, 
if, although less general and certain, means of determining 
the atomic and molecular weights is afforded by isomorphism. 
By this is understood the phenomenon observed by Mitscher- 
lich (1819), that chemically similar bodies possess the same or 
almost the same crystal form; as an essential mark of iso- 
morphous bodies is added their ability of crystallizing to- 
gether—to form so-called isomorphous mixtures. Conversely, 
therefore, from the isomorphism of two compounds may be 
concluded an analogous chemical composition, a similar num- 
ber of atoms in the molecule. From this, such quantities of 
the elements which can replace each other in isomorphous 
compounds, are accepted as the relative atomic weights. 
Thus, for example, for the metals calcium, strontium and 
barium, not yielding any volatile compounds, the atomic 
weights were not derived from the heat capacity; the isomor- 
phism, however, of many of their derivatives with those of 
magnesium determined the same; the quantities of these 
elements, replacing 24 weight parts of magnesium (1 atom), 
were accepted as the true atomic weights. 
In the present state of chemistry, but merely secondary im- 
portance attaches to isomorphism as a method of determining 
atomic weights. The phenomena of pleomorphism, according 
to which one and the same substance frequently possesses sev- 
eral crystalline forms, teach us that the latter are not only 252 
INORGANIC CHEMISTRY. 
dependent upon the chemical molecules, but that the latter 
(according to as yet unknown laws) may unite to more com- 
plicated crystal molecules. Hence, isomorphism affords a 
means for determining the molecular size of solid substances. 
Upon the other hand, many cases are known where chemi- 
cally dissimilar compounds possess similar isomorphous crys- 
talline form. Thus, dimorphous calcium carbonate (CaC03), 
as calcite is isomorphous with soda Saltpetre (N03Na), while 
as aragonite it is isomorphous with potassium saltpetre 
(KN03). Consequently, isomorphism is only to be applied 
with care, in chemical conclusions. Yet, generally, it is seen 
that chemically similar bodies have like crystalline form, es- 
pecially if the similarity of the elements according to groups 
be taken into consideration, as expressed in the periodic sys- 
tem. Thus, the isomorphism of the sodium compounds with 
the silver and cuprous derivatives, of the permanganates with 
the perchlorates (C104K), of the chromates with the sulphates 
(S04Na2), confirms the relations presented in the periodic 
system. Details upon this will be noticed in the consideration 
of the individual groups. 
CHEMICAL PROPERTIES OF THE METALS. 
The metals combine, generally without difficulty, with the 
metalloids, and form Avith them well characterized compounds 
the properties of which are essentially different from the ele- 
ments composing them. The greater the chemical difference of 
two bodies (metals and non-metals, bases and acid) the more 
energetic, in general, is their effort to unite, and the more 
different and more stable the resulting products. As we have 
seen, the analogous metalloids (the group of chlorine, of 
sulphur) form but slightly characteristic derivatives with each 
other. Similarly also the metals, by fusion with each other, 
form indefinite metal-like compounds, known as alloys. 
Alloys, for the solid condition, are essentially the same as 
solutions for the liquid. Solutions and alloys constitute the 
change from the mechanical mixtures to the real chemical 
compounds. In both instances the constituents possess only a 
slight affinity for each other, and, therefore, unite in almost all 
proportions to the so-called undetermined compounds (see p ). 
We, however, know that in solutions definite compounds fre- 
quently preexist; thus in an aqueous solution of sulphuric AMALGAM 
acid the hydrate H2S042H20; in aqueous nitric acid the 
hydrate HN03, H20. Also in the solutions of the salts crystalliz- 
ing with water of crystallization at definite temperatures, definite 
compounds with water (e. g., Naa S04, 10II2O, Co CJ2, 611,0) 
preexist. Similarly constituted combinations also appear to 
be present in the alloys, which represent compounds according 
to atomic relations. Antimony and tin form a crystalline 
compound of the composition Sb2Zn3. The crystalline form is 
always influenced by definite chemical compounds. This 
double character of the alloys manifests itself in their prop- 
erties. In many respects they show the average deportment 
of the metals from which they arise. By combining the 
various metals we can procure alloys of the desired properties; 
on this is founded the technical application of the same. Thus, 
to gold and silver, which are very soft in pure condition, can 
be imparted a greater hardness if they are alloyed with cop- 
per ; and the latter, again, may be rendered harder by fusion 
with zinc. In other properties of alloys, however, is exhibited 
the character of a chemical compound. Their temperature 
of fusion generally is not the average of the metals constitut- 
ing them, but always lies lower. An alloy of 8 parts lead, 15 
Bi, 4 Sn and three of Ca melts at 65°, although each of the 
single metals fuses above 200°. 
Mercury is able to dissolve almost all metals forming al- 
loys known as amalgams, which can crystallize. With hydro- 
gen, which, in a chemical respect, possesses complete metallic 
character, the most metals do not combine, probably owing to 
its volatility. Only palladium, potassium and sodium furnish 
the compounds Pd2, K2H and Na2H, which deport themselves 
as alloys, while copper forms a pulverulent compound 
(CuH). That antimony yields a gaseous product (SbH3), de- 
pends upon its pronounced metalloidal character. The ability 
of individual metals of the platinum and iron group to permit 
the passage of hydrogen at a red heat depends, probably, upon 
a chemical attraction; hydrogen first dissolves and is then 
evaporated again. 
Halogen Compounds. The metals unite directly with the 
halogens to salt-like compounds, which are not decomposed by 
water at ordinary temperatures, and, in general, are very 
stable; on the opposite hand, the halogen compounds of the 254 
INORGANIC CHEMISTRY. 
metalloids (excepting those of carbon) are easily broken up by 
water. These compounds are also produced by the action of 
the haloid acids upon the free metals, their oxides, hydrates 
and carbonates, whereby they plainly characterize themselves 
as salts of the haloid acids. A third procedure for the forma- 
tion of chlorides and bromides, essentially analogous to the 
first, rests upon the simultaneous action of carbon and chlo- 
rine, or bromine upon the oxides (see Chloride of aluminium 
and silicon). 
Corresponding to the various valences of the metals exist 
the following compound forms of halogen derivatives:— 
i 
KC1 
ii 
ZnCl2 
hi 
InCl, 
IV 
SnCl4 
v 
TaCl5 
VI 
WnCl6. 
The higher valence of the elements pronounces itself more 
emphatically in their more stable oxygen compounds. 
OXIDES AND HYDROXIDES—HYDRATES. 
The 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 dis- 
solve readily in water and form strong basic hydroxides or 
hydrates (KOH, Ca(OH)2), which, mostly, are not decomposed 
by glowing. Other metals (the so-called heavy metals) oxi- 
dize and decompose water only at higher temperatures ; their 
oxides are insoluble in water, generally afford no hydrates, as 
the latter, upon heating, readily decompose into oxides 
(anhydrides) and water:— 
Zn(OH)a =ZnO + H20. 
They are of a less basic nature, and their soluble salts 
usually exhibit acid reaction. Some metals, finally, 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:— 
CuS04 + 2K0H = K2S04 + Cu(OH)a. OXIJDES AND HYDRATES 
The different valence of the metals exhibits itself most distinctly 
in their salt-forming oxygen derivatives. Corresponding to the eight 
groups of the periodic system of the elements we have* the follow- 
ing eight forms or types of the highest salt-producing oxides (see p. 
241) :— ' 
I 
K20 . 
II 
MgO 
III 
A1203 
IV 
Sn02 
V 
Bi205 
VI 
CrO, 
VII 
Mna07 
VIII 
0s04 
From these are derived the hydroxides or hydrates: — 
KOH, 
fig(OH), 
hi 
Al(OH), 
IV 
Sn(OH)4 
Bi(OH)5 
Cr(OH)6 
VII 
Mn(OH), 
The oxides and hydrates of the two first forms possess a strong basic 
character and furnish salts with acids. In the oxides and hydrates of 
the 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 pass into metahydrates, which retain the acid 
character. Thus, from Al(OH)3 is derived AlO.OH, which yields salt- 
like compounds, e. g-, AlO.OK; from Sn(0H)4 is derived the stannic 
acidSnO(OH)2 and its salts, as Sn03K2. The oxides of the three 
last groups, finally, are only of an acid nature, and afford salts with 
bases. Their corresponding highest hydrates are very unstable or do 
not exist; inasmuch as, giving up 1, 2 and 3 molecules of water, they 
yield the ordinary acids (p. 241):— 
BiO,H 
Bismuthic 
acid. 
Cr04H2 
Chromic 
acid. 
VII 
Mn04H2 
Permanganic 
acid. 
HNO, 
Nitric acid. 
so4h2 
Sulphuric acid. 
C104H, 
Perchloric acid. 
Like the metalloids, the metals of the four last series form lower 
oxides and hydrates (p. 241) in which they exhibit a lower atom- 
icity :— 
Sn (OH)2, 
III 
Bi(OH)„ 
Mo (0H)4, 
fin (0H)4 
These lower oxides have a basic character, and it is the more pro- 
nounced the further removed they are from the limiting form. 
Throughout they resemble, in deportment, the corresponding combi- 
nation forms of the metals of the three first groups. 
The metals of the two first groups have higher oxygen 
compounds, called peroxides, e. g., Na202, Ba02. These do 
not form corresponding salts, and readily lose an atom of 
oxygen. By the action of dilute acids hydrogen peroxide is 
produced:— 
Ba02 +2HCl = BaCl2 +H202. 256 
INORGANIC CHEMISTRY. 
Based on this reaction, it is very probable that in the per- 
oxides, as.in that of hydrogen, the oxygen atoms are ar- 
ranged in a chain-like manner :— 
Na— 
Na—0/ 
Ba/°\ 
\o/ 
When concentrated acid acts upon them, oxygen is evolved, 
and salts of the lowrer oxides result; heated with hydrochloric 
acid, chlorine is generated :— 
Ordinarily, all higher oxides which are not able to form salts and 
which evolve chlorine with hydrochloric acid are termed peroxides, 
e. g., Pb02,lead peroxide, and Mn02, manganese peroxide. However, 
the compounds do not, indeed, possess the structure of true peroxides. 
IV 
Lead dioxide, Pb02, is throughout analogous to tin dioxide, Sn02, 
and is capable of combining with bases ; therefore, we must grant 
in it a direct union of the two oxygen atoms with tetra-tomic lead. 
So, too, is manganese probably tetratomic in manganese peroxide. 
The difference of these oxygen compounds from the true peroxides 
shows itself in that they are not able to form hydrogen peroxide. 
Finally, some monatomic metals are capable of forming oxides 
containing 4 atoms of metal, e. g., K40, Ag40 ; these compounds are 
termed quadrant oxides or suboxides. 
Ba02 + 4HC1 = BaCl2 + 2 H20 + Cl2. 
Sails. By the action of bases upon acids, salts and water 
are formed:— 
These are also produced by the direct union of basic with 
acid oxides: Na20 S03 = Na2S04, and by the action of 
metals upon the acids. Hence, usually, the salts are viewed 
as acids in which hydrogen is replaced by metals. Upon 
inquiring, however, into the composition of salts we discover 
them so constituted that a diatomic oxygen atom connects the 
metal with the acid radical (p. ). 
NaOH + N03H = N03Na + H20 
The salts, therefore, according as it is more practicable, can 
be regarded as acid derivatives, and also as derived from the 
basic hydrates by replacement of hydrogen. 
As we have seen, the polybasic acids form by replacement 
of one or several hydrogen atoms the primary, secondary, 
K—0—H 
Potassium hydrate. 
K—0—N02 
Potassium nitrate. 
H—0—N02 
Nitric acid. 257 
tertiary, etc., salts. Similarly are derived from polyatomic 
metals (or the polyacid, polyhydric bases) primary, secondary, 
etc., salts:— 
ni ( OH 
Bi -J OH 
(no3 
in/OH 
Bi-N03 
\N03 
ni r no. 
Bi\ NO, 
(NO, 
Primary bismuth 
nitrate. 
Secondary 
bismuth nitrate. 
Tertiary 
bismuth nitrate. 
Such salts in which not all the hydroxyl groups of the 
polyacid hydroxide are replaced by acid residues are called 
basic:— 
Pb/on 
Fb\N03 
7 r oh 
Zn { ci 
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 poly or 
anhydro acids; similarly, also, can polyhydric bases form poly- 
hydrates :— 
.OH 
Cu( 
>° 
Cu( 
xOH 
/OH 
Pb( 
>° 
Pb( 
>° 
Pb( 
xOH 
from which, by replacement of hydroxides by acid residues 
basic salts are obtained (see copper and lead). 
By replacement of the hydrogen atoms in the polyhydric 
acids or bases by various radicals arise the so-called mixed or 
double salts:— 
/K 
so4/ 
Cu 
so4x 
XK 
/K 
SO/ 
\ 
Al 
so./ 
)K 
P04 4 
i H 
Pot. Am. Phosphate. 
/Cl 
PbC 
XC03 
PbX 
lb\Cl 
Pot. Copper Sulphate. 
f NO, 
CrJNO, 
l Cl4 
Pot. Aluminium Sulphate. 
The halogen double salts are usually viewed as molecular 
compounds :— 
MgCl2.KCl, AuClj.KCl, PtCl4.2KCl. 258 
INORGANIC CHEMISTRY. 
If, however, the fluorides of boron and silicon BF13, IvFl, 
SiFl4, 2KF1 be derived from peculiar, atomic constituted acids, 
HBFh, H2SiFlB, then a peculiar union of atoms may be 
regarded as existing in the frequently very similar and isomor- 
phous metallic double chlorides. 
ACTION OF METALS UPON SALTS AND ACIDS. 
We have seen that the metals by solution in acids are able 
to form salts. In this case the hydrogen is directly replaced 
by the metal and separated in a free condition (providing in 
the moment of its formation it does not act upon the acid) :— 
Zn -j- S04H2 — ZnS04 -j- H2. 
In the same manner do the metals deport themselves with 
the salts. Zinc introduced into a solution of copper sulphate 
is dissolved to sulphate and metallic copper deposits :— 
Zn + CuS04 = ZnS04 + Cu. 
Herein is shown the perfect analogy between acids and salts. 
In chemical nature hydrogen is a metal. Hence the acids 
may be viewed as hydrogen salts ; hydrogen sulphate for sul- 
phuric acid, hydrogen nitrate for nitric acid, etc. The simi- 
larity of salts and acids shows itself, too, in their acidity. All 
soluble salts of the metals, whose hydrates are weak bases, 
exhibit acid reaction, and color blue litmus paper red. Only 
the salts of the strong basic metals, like potassium and cal- 
cium, show a neutral or basic reaction—providing the base is 
stronger than the acid. 
The displacement of metals from their salts by others, was 
formerly regarded as exclusively influenced by their elec- 
trical deportment. Indeed, the more electro-positive basic 
metals replace the electro-negative, less basic. In the follow- 
ing series each metal throws out from solution those preced- 
ing it: Au, Pt, Ag, Hg, Cu, Pb, Sn (Fe, Zn). Iron and zinc 
precipitate almost all the heavy metals from solutions of their 
salts. The most strongly positive potassium is able to displace 
all other metals. This is most evident by the action of molten 
potassium upon the haloid salts—a reaction which frequently 
serves for the separation of the metals in free condition :— 
AlClg + 3K = A1 + 3KC1. ACTION OF METALS. 
259 
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 vola- 
tility of the hydrogen; in fact, we know that hydrogen, 
under pressure, is capable of separating gold, silver, and some 
other metals from their salt solutions. 
Formerly great importance was attributed to the electrical behavior 
of the elements, and all were arranged in an electro-chemical series, in 
which oxygen figured as the negative and potassium the positive pole, 
0 -f- K. The opinion prevailed that the chemical affinity ot the 
elements depended upon their electrical differences, and that chemical 
union occurred because the opposite electricities united — electro- 
chemical theory of Berzelius. Now, however, we know that in the 
expression of chemical affinity only secondary importance is attached 
to the electrical deportment of bodies. Although in general the 
affinity corresponds to the electrical difference, yet this does not 
always occur. Thus the strongly negative chlorine expels bromine or 
iodine from their hydrogen, and nearly all their metallic compounds ; 
conversely, chlorine and bromine are displaced by iodine from their 
oxygen compounds as (C10SH) and C104H). Similarly, lead separates 
tin from its chloride, SnCl4, while on the other hand, tin throws out 
lead from the solution of its oxides in alkalies. 
At present it is established that the mutual deportment of the metals 
is dependent upon and regulated by their thermo-chemical relations. 
A metal displaces another from its oxygen salts, as also from its oxides, 
sulphides or halogen compounds, if the heat of formation of the 
resulting bodies is greater than the ones acting; this agrees with the 
principle of greatest heat development. Thus copper displaces silver 
from its sulphate, because the heat of formation of the copper sulphate 
(in aqueous solution) is about 33570 calories greater than that of silver 
sulphate. Sulphuric acid dissolves most metals with heat liberation, 
because their heat of formation— 
(S,04,H2) = 192900 
(S,04,Ha,Aq) = 210760 
is less than that of the most sulphates. The heat of formation of lead 
sulphate (Pb,S,04) equals 213500; therefore lead is not dissolved by 
concentrated but by dilute sulphuric acid, because in the latter the 
produced lead sulphate is not soluble. For the same reason potassium 
displaces almost all the other metals ; on the opposite hand, potassium 
is separated by sodium amalgam, with formation of potassium 
amalgam, as the heat of formation of the latter is much greater than 
that of sodium amalgam, and therefore, in the equation— 
(K, X) + (Na,Hg) = (Na, X) + K,Hg) 
the heat upon the right side overbalances. (Berthelot). 
Although the affinity relations dependent upon the quantity of heat 
frequently also correspond with the electrical differences of the free 
elements, this is so influenced that the electro-motive energy is induced 
by the heat, and proportional to the same (see p. 262). The latter 260 
INORGANIC CHEMISTRY. 
constitu'es the primary cause; it varies in the different compounds and 
explains the opposing deportment of the elements. Chlorine displaces 
iodine in iodides, not because it is more strongly electro-negative, 
but because the heat of formation of the chlorides is greater. Con- 
versely, from chloric acid chlorine is eliminated by iodine. 
ELECTROLYSIS OF SALTS. 
On subjecting a salt in a fused or dissolved condition to the 
action of a galvanic current, it is decomposed, so that the 
metal separates upon the negative pole and the acid group or 
halogen upon the positive:— 
4* — 
NaCl = Na + Cl. 
The oxygen salts behave in the same way; the metal upon 
the negative pole, the acid residue upon the positive:— 
_l_ — 
CuS04 = Cu + so4. 
As the liberated acid residue cannot exist in a free condition, 
a secondary reaction occurs, by which it generally, especially 
in the electrolysis of aqueous solutions, breaks up into oxygen 
and an acid oxide, which with the water of the solution forms 
the acid anew :— 
Thus, in the electrolysis of salts, the metal and oxygen 
separate out—the first at the negative, the latter at the posi- 
tive pole. That indeed the decomposition occurs in the manner 
indicated is confirmed by the fact that free acid arises at the 
positive pole. 
All neutral salts are decomposed in like manner. If, 
however, the metal contained in the salt acts upon water 
when free, manifestly a secondary reaction occurs at the 
negative pole. The real electrolytic decomposition of potas- 
sium sulphate would then take place according to the follow- 
ing equation :— 
so4 + h2o = so4h2 + 0. 
_1_ — — 
so4k2 = k2 + so3 + 0. 
The separated potassium decomposes water with form- 
ation of potassium hydrate and the disengagement of 
hydrogen:— 
K + HOH = KOH + H. ELECTROLYSIS OF SALTS. 
261 
Therefore, hydrogen and potassium hydrate occur as defi- 
nite decomposition products, at the negative pole; at the 
positive, however, oxygen and sulphuric acid. On coloring 
the liquid exposed to the electrolysis with a little violet syrup 
that part at the -}- pole will be reduced by the acid formed, 
while that at the —pole will have a green color from the base. 
That, in fact, the electrolytic decomposition of potassium 
sulphate and similar salts proceeds in the manner given, may 
be experimentally proven by using mercury as negative 
electrode; then the separated potassium unites to an amalgam 
with the mercury, which acts only gradually upon the water. 
Formerly it was supposed that the alkali salts, upon electrolysis, 
were directly decomposed into metallic and acid oxides, which with 
water, yielded the hydrates (KOH and S04H2); the appearance of 
H and 0 was attributed to the simultaneous electrolytic decomposition 
of water (a view which was set aside by the behavior of the other salts). 
Based on this erroneous idea all salts were held to be binary com- 
pounds of the metallic oxides (bases) with acid oxides (acids), 
e.f/., K20, SO3 = K2S04, K20, N205 = 2KN03—dualistic theory 
of Berzelius. The acids and bases were also thought to be binary com- 
pounds of a metallic oxide or acid anhydride with water: — 
K20,H20 = 2K0II, S03H20 = H2S04. 
The acid oxides or anhydrides were termed acids and the true acids 
hydrates. 
Other compounds are decomposed in the same way as the 
salts. Thus molten caustic potash KOH, breaks up into K 
and OH ; the first separates in metallic form upon the nega- 
tive pole (and gradually acts upon KOH with hydrogen disen- 
gagement), while at the positive pole water and oxygen appear 
—produced by decomposition of the at first formed hydrogen 
peroxide :— 
It is, therefore, probable that also water is decomposed in 
analogous manner:— 
<0H)2 = H20 + 0. 
the peroxide produced at first breaks up, however, mostly into 
water and oxygen. 
Considering the quantity relations which are deposited from 
various compounds, by the same electric current, we Avill dis- 
cover that invariably a like number of valences is dissolved in 
like time, i. e., equivalent quantities in the idea of the valence 
theory are separated (p. 162). Thus in the simultaneous de- 
2H0H = H2 + 02H2; 262 
INORGANIC CHEMISTRY. 
composition of hydrochloric acid, water and ammonia (pp. 64, 
88, 119) equal volumes of hydrogen (= 1 part) are liberated, 
while at the positive pole 1 volume chlorine (= 35.5 parts), 
2 volume oxygen (= 8 parts), and & volume nitrogen (= 4.66 
parts) appear. The electrolytically decomposed quantities 
stand, therefore, in the ratio:— 
HC1, 
H20, 
2 
h3n 
3 
In the same way, from all metallic chlorides (and other salts) 
equal quantities of chlorine are set free (as the chlorine atoms 
in all are alike), whilst the quantity of the precipitated metals 
agrees with their chemical activity. The electrolytically de- 
composed quantities of different salts stand in the following 
relation :— 
CuCl2, 
2 
Cu2Cl2, 
2 
SbCl3, 
3 
Fe2CI6) 
6 
SnCl4, 
4 
HgCl2, 
2 
Hg2(N03)2. 
2 
AgN03, 
As the quantity of heat liberated in the formation of equivalent 
quantities of the compounds, consequently, too, that necessary for the 
decomposition, is very different; equivalent quantities, however, being 
separated by the galvanic current, the energy of the latter must dis- 
tribute itself unequally upon the various electrolytes and, indeed, in 
proportion to the heat of decomposition. Joule has experimentally 
proven that the electro-motive energy disengaged by a galvanic ele- 
ment is proportional to (if no secondary actions occur) or equal to the 
amount of heat of the reaction producing it. Thus the electro-motive 
energy of a Daniell’s element (combination of zinc in dilute sulphuric 
with copper in copper sulphate) depends on the chemical replacement 
of copper in its sulphate by zinc, a reaction setting free 50100 calories. 
As, however, 69000 calories are requisite for the decomposition of 
one molecule of water, it is obvious that the electrolysis of water 
cannot be effected by one Daniell’s element, but that the combination 
of at least two of these is necessary. 
A chemical action will frequently occur when two salts in 
solution or fusion come together. The resulting phenomena 
Berthollet endeavored (close of preceding century) to explain 
in the following manner, in that he referred them to pure 
physical causes and excluded every special chemical affinity. 
In the opinion of Berthollet, in the solution of two salts 
four always arise. For example, on mixing solutions of 
copper sulphate and sodium chloride, there exist in solution 
copper sulphate, sodium sulphate, copper chloride and sodium 
chloride:— 
2CuS04 + 4NaCl yield 
CuS04 + NaT2S04 + CuCl2 + 2NaCl. TRANSPOSITION OF SALTS. 
263 
That copper chloride is really present in the solution 
together with the sulphate, follows, from the fact that the 
blue color of the latter, by the addition of sodium chloride, 
acquires a greenish color, peculiar to the copper chloride. 
Suppose one of the four salts formed in the solution is insol- 
uble or volatile, the reaction will occur somewhat differ- 
ently. Upon adding barium chloride to the copper sulphate 
solution, at the beginning, as in the first case, four salts will 
be formed. The barium sulphate produced separates, how- 
ever, in consequence of its insolubility, the equilibrium of 
the four salts will be disturbed, and new quantities of CuS04 
and BaCl2 act upon each other to complete transposition :— 
CnS04 + BaCl2 = BaS04 + CuCl2. 
And the same may therefore be explained by the insolubility 
of the barium sulphate. On adding HC1, or soluble chlorides 
to the solution of a silver salt all the silver is precipitated as 
chloride, as the latter is insoluble. 
Take another example. On adding sulphuric acid to a 
solution of potassium nitrate there is apparently no percept- 
ible alteration. We may suppose that the four compounds, 
KN03, K2S04, H2S04 and HN03, are present in the solution. 
Upon warming the latter volatile nitric acid will evaporate, 
and, in proportion to its separation, new quantities of potassi- 
um nitrate and hydrogen sulphate will act upon each other 
until the transposition is complete :— 
2KN03 + H2S04 = K2S04 + 2HN03. 
The decomposition of potassium nitrate by sulphuric acid, 
therefore, is the consequent of the volatility of the nitric acid. 
Sulphuric acid decomposes sodium chloride in the cold, 
because hydrogen chloride is volatile. Carbonates are 
decomposed even by very weak acids, because the carbonic 
acid, H2C03, at once separates gaseous carbon dioxide, C02. 
By such physical causes, in many instances, the chemical 
transpositions may be explained, and there is no doubt that 
to them attaches an important role. It is, however, not justifi- 
able to ignore any special chemical affinity between the 
various substances, as happened with Berthollet. Inde- 
pendent of all physical causes, chemical affinity is a self- 
active cause. This is obvious in the solutions of salts. Mix, 
e. g., ferric chloride with potassium acetate, and there is 
obtained a dark red solution, in consequence of the formation of 264 
INORGANIC CHEMISTRY. 
iron acetate. Although an insoluble salt is not produced here, 
yet the rearrangement of the two salts, evident from the 
optical properties of the solution, is a perfect one ; in solution 
are only iron acetate and potassium chloride :— 
FeaCl6 + 6C12H302K = (C2H302)6Fe2 + 6KC1. 
Pot. acetate 
The transposition is determined by the strong affinity of 
potassium for chlorine and by the weak basic nature of ferric 
oxide. If the difference between the affinities of the bases 
and salts is not so great, then four salts will exist in solution ; 
their quantity, however, will be proportional to the different 
affinities and determined by the equilibrium of all the forces 
of attraction. Thus in the previously mentioned solution of 
copper sulphate and sodium chloride are contained four salts, 
the quantities of copper chloride and sodium sulphate are, 
however, much greater than those of copper sulphate and 
sodium chloride (proven by the optical properties of the 
solution), because the affinity of sulphuric acid for sodium is 
greater than the same for copper. 
To make this plainer, let us examine the following example : 
barium sulphate is almost perfectly insoluble in water and 
acids, hence always results in the action of sulphates upon 
barium salts:— 
BaCl2 + K2S04 = BaS04 + 2KC1. 
If barium sulphate, however, be boiled with potassium car- 
bonate, the reverse reaction occurs; barium carbonate and 
potassium sulphate are produced :— 
Although the carbonate is somewhat less insoluble than the 
sulphate, the transposition is yet a complete one, because 
the affinity of S04 for potassium is greater than for barium. 
The relative affinity of various compounds for each other 
is, as yet, but little investigated. A criterion for the magni- 
tude of the same is afforded by the quantity of heat disengaged 
in chemical unions. Investigations in this direction have re- 
cently been taken up and executed with much zeal. 
S04Ba + K2C03 =BaC03 +K2S04. 
The thermo-chemical investigations of recent date have shown 
that mutual transpositions of the salts, or the action of the acids upon 
bases and salts—just as all other chemical affinity relations, are de- 
pendent upon and regulated by the law of the greatest heat develop- 
ment. The relations are here, however, sometimes more complicated, ALKALI METALS. 
265 
as the production of acid salts or double salts, further, the decomposi- 
tion (dissociation) of many salts by water, into their component bases 
and acids, must be taken into consideration. By regarding these 
secondary reactions all chemical transpositions, according to Berthel- 
lot, find their explanation in the heat regulating them, without making 
it necessary to accept a particular affinity function of the acids and 
bases, which is termed their avidity. Even the physical properties of 
the resulting compounds, their volatility and insolubility, are, contrary 
to the supposition of Berthellot, only of secondary and slight influence. 
GROUP OF THE ALKALI METALS. 
Potassium, 39 
Rubidium, 85.2 
Caesium, 132.5 
Lithium, 7 
Sodium, 23 
(Ammonium), 
The metals of this group are decidedly the most pronounced 
in metallo-basic character, and this constitutes a visible con- 
trast with the elements of the chlorine group, the most ener- 
getic among the metalloids. This contradictory character of 
both groups is seen, too, in their monovalence; in their com- 
binations with each other, their affinities saturate by single 
atoms. The more distinct the chemical character of two 
elements and the more unlike they are, so much the simpler 
and the more definite in general are the expressions of 
equivalence between them. 
The alkali metals in physical and chemical properties 
exhibit great similarity. They oxidize readily in the air, 
decompose water violently, even in the cold, with the forma- 
tion of strong basic hydrates, readily soluble in water, 
cilled alkalies (caustic potash, caustic soda), hence the name 
alkali metal. They are not decomposed by ignition. Their 
chemical energy increases with increasing atomic weight 
(more correctly atomic volume, p. 245) ; sodium is more 
energetic than lithium, potassium more than sodium, and 
rubidium more than potassium. Caesium is not known in 
free condition, but, to conclude, from its compounds, it possesses 
a more basic character than rubidium. We see that in other 
analogous groups (of chlorine, oxygen, phosphorus, carbon), 
with the increasing atomic weight the metalloidal, negative 
character diminishes and the basic increases. 
At the same time with the atomic weights the specific gravi- 
ties increase; as, however, the increase of the first is greater 
than that of the latter, so are the atomic volumes (the quo- 266 
INORGANIC CHEMISTRY. 
tients ■>ipAgr ) always greater. The increasing fusibility and 
volatility correspond to the increase of the atomic volumes; 
rubidium distills already at a red heat, while lithium only 
volatilizes with difficulty:— 
Li 
Na 
K 
Bb 
Cs 
Atomic Weight, 
Specific Gravity, 
7 
23 
39 
85 
132 
0.59 
0.97 
0.86 
1.62 
(2,4) 
Atomic Volume, 
11.9 
23.7 
45.4 
56.1 
— 
Fusion Temperature, 
180° 
95.6° 
62.5° 
38.5° 
— 
If, consequent!}7, the alkali metals exhibit in their chemical 
deportment a great similarity, we will discover yet more marked 
relations between potassium, rubidium and calcium upon the 
one hand, and lithium and sodium on the other, as appears in 
the periodic system of the elements. Especially is this no- 
ticed in the salts. The first three metals form difficultly soluble 
tartrates and chlorplatinates (see platinum). Their carbon- 
ates deliquesce in the air, while those of sodium and lithium 
are stable under similar circumstances ; the latter is, indeed, 
tolerably insoluble ill water. The phosphates deport them- 
selves similarly; lithium phosphate is very difficultly soluble. 
It must be remarked that the normal carbonates and phos- 
phates of all other metals are insoluble. In lithium, then, 
which possesses the lowest atomic weight, it would seem the 
alkaline character has not yet reached expression, and ap- 
proaches, in many resjiects, the elements of the second group, 
especially magnesium, just as beryllium approaches aluminium. 
This is indicated in the table, p. 237, by the position of the 
elements. The elements of the two small periods are, indeed, 
similar, but not completely analogous, while, in K Rb Cs the 
homology of the three great periods finds expression. 
The affinity relations of the alkalies are expressed and explained by 
their thermo-chemical relations. Generally, with increase of atomic 
weights, the heat liberation is greater; thus, e. g., in the formation of 
the chlorides and hydrates (the numbers I’epresent large calories, 
p. 243). 
(Li,Cl) = 93.8 
(Na,Cl) = 97.7 
(K,C1) = 105.6 
(Na 20) = 100.2 
(Li,Cl,Aq) = 102.2 
(Na,Cl.Aq) = 91.5 
(JK.Cl.Aq) =101.1 
(Na,0,ll) = 102.0 
(Li,O.H.Aq) = 117.4 
(Na,0,H.Aq) = 111.8 
(K.O,H, Aq) =116.4 
(K,0,H) = 103.9 POTASSIUM. 
267 
The varying depoi'tment of the lithium compounds, which frequently 
show a greater heat disengagement than the compounds of sodium, fir ds 
expression in the position of lithium in the periodic system. On the 
cor.trary, it is very probable that, with the true homologues of potas- 
sium—rubidium ana caesium, a constant increase of heat occurs. 
From the above numbers is explained, upon the basis of the prii ci- 
ple of the greatest heat liberation, how that sodium and lithium are 
displaced from their chlorides, etc., by potassium. In like manner, 
also, are the most other metals separated by potassium, because the 
heat of formation of the potassium compounds, generally, is much 
greater (see p. 259). From the comparison with the heat of formation 
of water (H20 = 690C0 calories) is explained further its ready de- 
composition by the alkali metals. All metals which, in the production 
of their oxides, MezO, or hydrates, MeOH, disengage more than 
69000 calories, decompose water, and with so much more energy, the 
greater the difference of heat. The insolubility of the oxides consti- 
tutes an obstacle to the action; this, however, by addition of 
neutral solvent acids, may be removed (see Aluminium). Conversely, 
all oxides with a less heat of formation are reduced by hydrogen. 
POTASSIUM. 
In nature potassium is found principally in silicates, viz.: 
feldspar and mica. By disintegration of these silicates occur- 
ring in many rocks potassium passes into the soil and is 
absorbed by plants; the ashes of the latter consist chiefly of 
different potassium salts. The chloride and sulphate are also 
found in sea water, and in large deposits at Stassfurt, at 
Magdeburg, and in Galicia, where they were left by the 
evaporation of the water of enclosed seas. Metallic potassium 
was first obtained by Davy, in the year 1807, by the decom- 
position of the hydrate, by means of a strong galvanic current. 
At present it is prepared by igniting an intimate mixture of 
carbon and potassium carbonate :— 
K = 39.1. 
KaC03 + 2C = 2K + 3C0. 
Such a mixture may be made by the carbonization of 
organic potassium salts, like crude tartar. It is then ignited 
to white heat, in an iron retort, and the escaping potassium 
vapors collected in receivers of peculiar construction, filled 
with rock-oil. The latter, an hydrocarbon, serves as the best 
means of preserving potassium, which would otherwise oxidize 
in the air and decompose other liquids. 
In a fresh section, potassium shows a silver white color and 
brilliant metallic lustre. At ordinary temperatures it is soft, 268 
INORGANIC CHEMISTRY. 
like wax, and may be easily cut It melts at 62.5 and is 
convei’ted, at a red, heat into a greenish vapor. In the air it 
becomes dead in color at once, by oxidation ; heated, it burns 
with a violet flame. Decomposes water energetically, with 
formation of potassium hydrate and the liberation of water. 
If a piece be thrown upon water, it will swim on the 
surface with a rotatory motion; by the reaction so much 
warmth is disengaged that the generated hydrogen and the 
potassium inflame. Finally, there usually results a slight 
explosion, by which 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. Potassium 
combines directly and very energetically with the halogens. 
On conducting hydrogen over metallic potassium heated to 300- 
400°, there results potassium hydride, K2H, a metallic, shining, 
brittle compound, which upon stronger heating, readily in vacuo, is 
again decomposed. Exposed to air it ignites spontaneously. The 
similarly obtained sodium hydride, Na2H does not possess this latter 
property. 
Potassium forms three oxygen compounds, of which only 
the following oxide yields corresponding salts. 
Potassium Oxide—K20—results from the oxidation of 
thin pieces of metallic potassium in dry air, and by heating 
potassium hydrate with metallic potassium in a current of 
hydrogen:— 
It is a white powder, fusing at a high temperature, and 
evaporating somewhat. With water it gives potassium 
hydrate, with evolution of much heat. 
Potassium peroxide, K02 or K204, and potassium suboxide, lv40, 
are very unstable, and pass readily into potassium oxide. The first is 
formed, together with potassium oxide, by the combustion of potassium 
in dry air or oxygen, and is a yellow mass. The suboxide has a violet 
color, due to the oxidation of potassium vapors. 
Potassium Hydrate, or Caustic Potash—IvOH—is obtained 
by the action of potassium or its oxide upon water. For its 
preparation potassium carbonate is decomposed by calcium 
hydrate (slacked lime) :— 
2KOH + K2 = 2K20 + 1I2. 
The solution of 1 .part potassium carbonate in 10-12 parts 
water is boiled with 1 part slaked lime in an iron pot, until a 
Iv2C03 + Ca(OH)2 = CaC03 + 2K0H. POTASSIUM BROMIDE. 
269 
filtered portion causes no effervescence, when hydrochloric 
acid is added ; i. e., there is no longer any carbonic acid 
present. The turbid liquid becomes clear upon standing, as 
the insoluble calcium carbonate subsides. The clear solution 
of potassium hydrate is poured off, evaporated, the residue 
melted in a silver dish (wThich it does not attack) and poured 
into moulds. The thus prepared caustic potash is not 
entirely pure, but contains potasssium chloride and other 
salts. To get chemically pure hydrate potassium nitrate is 
fused with copper filings and the fusion lixiviated with water. 
Potassium hydrate forms a white, crystalline, tolerably 
easily fusible mass, which volatilizes undecomposed at a very 
high temperature. Exposed to the air it deliquesces, as it 
absorbs water and carbon dioxide and changes into carbon- 
ate. In alcohol, and especially water, it is very soluble. The 
solution reacts strongly alkaline, saponifies the fats, and has a 
corrosive action upon the skin and organic tissues; hence it 
cannot be filtered through paper. From concentrated solu- 
tions at low temperature the hydrate KOH -f 2H20 crys- 
tallizes out. 
The haloid salts of potassium are formed by the direct union 
of the halogens with potassium, and by saturation of the 
hydrate or carbonate with haloid acids. They are readily 
soluble in water, have a salty taste, and crystallize in cubes. 
When heated they melt and are somewhat volatile. 
Potassium Chloride—KC1—occurs in Stassfurt in large 
deposits, as sylvite, and combined with magnesium chloride as 
carnallite (MgCl2, KC1 + 6H20). The latter salt serves as 
the chief source for the preparation of potassium chloride, 
which meets with varied application in the arts, and also for 
the preparation of potassium carbonate. The chloride crys- 
tallizes in vitreous cubes, of specific gravity 1.84. 100 parts 
water dissolve at 0° C., 30 parts, at 100°, 59 parts of the salt. 
Potassium Bromide—KBr—is generally obtained by warm- 
ing a solution of potassium hydrate with bromine, when the 
bromate also is produced :— 
6K0H + 8Br2 = 5KBr + KBr03 + 3H20. 
The solution is evaporated to dryness, mixed with charcoal 
and ignited, which reduces the bromate to bromide:— 
KBr03 + 8C = 3C0 +T£Br. 270 
INORGANIC CHEMISTRY. 
It is readily soluble in water and alcohol, and is employed in 
photography and medicine. 
Potassium Iodide—KI—may be prepared like the preced- 
ing. Usually it is obtained according to the following 
method: Iodine and iron filings are rubbed together under 
water, and potassium carbonate added to the solution of iron 
oxide ; this will precipitate ferrous-ferric oxide; carbon diox- 
ide escapes and potassium iodide will be found in solution. 
It forms large white crystals, is easily fusible and tolerably 
volatile. Its specific gravity equals 2.9. At medium tem- 
perature it dissolves in 0.7 parts water. The aqueous solu- 
tions dissolves iodine in large quantity. Many metallic, insol- 
uble iodides dissolve without difficulty in it, forming double 
iodides, e. g., Hgl2, 2KI. The iodide is employed in medicine 
and photography. 
Potassium Fluoride— KF1. Obtained by dissolving the car- 
bonate in aqueous hydrofluoric acid, forms easily soluble cubes. 
The aqueous solution attacks glass. It is greatly inclined to combine 
with other fluorides; KF1. HF1, BF13. KF1. On adding hydrofluo- 
silicic acid to the solution of potassium salts, a gelatinous precipitate 
of potassium silicoflaoride is thrown down; this is very difficultly solu- 
ble in water. 
Potassium Cyanide—KCN. By saturating potassium hydrate 
with hydrocyanic acid, and by heating yellow potassium prussiate (see 
Iron) we can produce this salt. It forms a white, easily fusible mass, 
which deliquesces in the air. The solution maybe easily decomposed. 
It crystallizes in cubes, has an alkaline reaction, smells like prussic 
acid, as this is set free by the carbon dioxide of the air. By fusion 
potassium cyanide reduces many oxides, and hence is employed 
in reduction processes. It is just as poisonous as prussic acid. It is 
applied in many ways, especially in photography and for galvanic 
silvering and gilding. 
Potassium Chlorate—KC103. Upon conducting chlorine 
gas through a hot concentrated potassium hydrate solution, 
the following reaction occurs :— 
6K0H + 3C12 = 5HG1 + KC103 + 3H20. 
When the solution cools, the difficultly soluble potassium 
chlorate separates out. Technically it is generally made 
by the action of chlorine upon a mixture of calcium hydrate 
and potassium chloride. The reaction occurs in two phases: 
first calcium chlorate is formed;— 
6Ca (0H)2 + 6C12 = 5CaCl2 -f Ca (C103) + GH20. 271 
This then reacts with the potassium cliloride 
POTASSIUM HYPOCHLORITE. 
Ca (C103)2 + 2KC1 = 2KC103 + CaCl2 
From the hot solution potassium chlorate crystallizes in 
shining tables of the monoclinic system, which are difficultly 
soluble in water (100 parts at ordinary temperature dissolve 
6 parts of the salt). Its taste is cooling and astringent. 
When heated it melts (at 330°), giving up a portion of its 
oxygen and changing to the Perchlorate—KC104—which on 
further heating decomposes into oxygen and potassium chlo- 
ride (see p. 172). As it gives up oxygen readily, it serves 
as a strong oxidizing agent. With hydrochloric acid it lib- 
erates chlorine:— 
KC103 + 6HC1 = KC1 + 3H,0 + 3Cla. 
Mixed with sulphur or some sulphides it explodes on 
heating and when struck a sharp blow. The igniting material 
upon the so-called Swedish (parlor) matches consists of anti- 
mony sulphide and potassium chlorate; rubbed upon the 
friction surface coated with red phosphorus they ignite. 
Potassium Hypochlorite—KCIO—is formed when chlorine 
is permitted to act upon a cold solution of potassium hydrate: 
It only exists in aqueous solution; when the latter is evap- 
orated the salt is decomposed into chloride and chlorate :— 
2K0H + CI2 = KC1 + KC10 + H20. 
3C10K = 2KC1 + C103K. 
The solution has a chlorine odor and bleaches strongly, 
especially upon the addition of acids. The bleaching solutions 
occurring in trade (Eau de Javelle) are prepared by the action 
of chlorine upon solutions of soda and potassium carbonate ; 
they also contain free hypochlorous acid. 
The oxy-salts of bromine and iodine are perfectly analogous to those 
of chlorine. Potassium Bromate—KBr03—and Potassium Iodate— 
KlOj—are prepared by the action of bromine or iodine upon potas- 
sium chlorate ; in these there is a direct substitution of chlorine (p. 
176). If the latter be passed through a hot solution of potassium 
iodate in potassium hydrate—the periodate of potassium, K104, arises; 
it is difficultly soluble, and upon heating decomposes into 0 and KI. 03, 
which then breaks up into potassium iodide and oxygen. 
Together with the normal periodates, K103. Nal03, exist other 
salts which are derived from the highest hydroxyl compound I(OH)7 
audits anhydro-derivatives (p. 174). These salts are very numerous 
and partly monoperiodates, fO(OH)5 and I02(OH)3, partly poly- 
periodates, produced by the condensation of several molecules of the 
highest hydrates, like I203(0H)8 and I205(0H)4. 272 
INORGANIC CHEMISTRY. 
Potassium Sulphate—K2S04—is formed in the action of sul- 
phuric acid upon potassium chloride, and as a bye-product in 
many technical operations. Crystallizes without water, in 
small rhombic prisms of a bitter, salty taste, and dissolves in 
10 parts HaO of ordinary temperature. It is employed prin- 
cipally for the preparation of potassium carbonate, according 
to the method of Le Blanc (see Soda). 
The acid or primary salt—HKS04—crystallizes in large 
rhombic tables, very readily soluble in water. It fuses about 
200°, loses water and is converted into Potassium pyrosulphate 
—K2S2Ot (p. 187)—which at 600° yields K2SO, and S03. 
The salts of sulphurous acid—the primary S03KH and the second- 
ary S03K2—form when S03 comes in contact with a potassium car- 
bonate 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 efferves- 
cence ceases and then cooled, pyrosulphite—K2S205—corresponding 
to the pyrosulphate, will crystallize out. 
Potassium Nitrate, Saltpetre, KN03, does not occur any- 
where in large quantity, but is widely distributed in the upper 
strata of the earth and in some regions of the hot zone (in 
Egypt and East India) it is exposed by disintegration. It is 
produced whenever nitrogenous organic substances decay in 
the presence of potassium carbonate—conditions, which are 
present in almost every soil. Upon the intentional introduc- 
tion of the same depends the artificial nitre production in the 
so-called saltpetre plantations. Manures and various animal 
offals are mixed with wood ashes (potassium carbonate) and 
lime, arranged in porous layers, and submitted to 2-3 years’ 
action of the air, whereby, from the slow oxidation of the ni- 
trogen, nitrates are produced. The heaps are then treated 
with water and potassium carbonate added to the solution, 
which contains potassium, calcium and magnesium nitrates, 
to convert the two last salts into potassium nitrate:— 
The precipitate of calcium and magnesium carbonate is 
filtered off and the solution evaporated. This procedure was 
formerly universally employed in the manufacture of potas- 
sium nitrate. At present, however, almost all of it is obtained 
by the decomposition of the sodium salt, occurring in large 
deposits in Chili, by means of potassium carbonate or chlo- 
ride:— 
Ca(N03)2 + K2C08 = CaCOg + 2KN03. 
NaN03 + KC1 = NaCl + KN03. POTASSIUM CARBONATE. 
273 
Warm saturated solutions of sodium nitrate and potassium 
chloride are mixed and boiled. Then four salts are formed, 
of which sodium chloride, the least soluble in hot water, sepa- 
rates. On cooling the solution, potassium nitrate, the least sol- 
uble in cold water, crystallizes ; sodium chloride is nearly 
equally soluble in the warm and cold water, for which reason 
the portion not separated by boiling remains in solution. 
Potassium saltpetre crystallizes without water in large six- 
sided prisms. It is far more soluble in hot than in cold water ; 
100 parts of water dissolve 244 parts at 100°, but at 0° only 
13 parts. It possesses a cooling taste, fuses about 340°, and 
decomposes by further heating into oxygen and potassium ni- 
trite, KN02. Heated with carbon it yields potassium carbon- 
ate :— 
4K N03 + 6C = 2K2C03 + 3C02 + 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 cannon). Upon 
an average, the powder consists of 7-5 per cent. KN03. 12 per cent, 
sulphur and 13 per cent, carbon, which closely corresponds to the 
atomic composition 2KN03 -f- S + 3C. Ihe decomposition of the 
powder by burning may be approximately expressed by the following 
equation:— 
The effectiveness of the powder, therefore, depends upon the disen- 
gagement of carbon dioxide and nitrogen gas, the volume of which is 
almost 100 times as great as that of the decomposed powder. 
2KN03 + S + 3C = K2S + 3C02 + Na. 
Potassium Nitrite, KN02, is obtained by the fusing of salt- 
petre with lead, which withdraws one atom of oxygen from 
the former. A white, fusible mass results; this deliquesces 
in the air. 
The potassium salts of phosphoric add: K3P04, K2HP04 
and KH2P04, meet with no practical application, as they are 
easily soluble in water and crystallize poorly; therefore, the 
sodium salts are generally used. The borates, too, B02K and 
B407K2 + 5H20 (See Borax), crystallize writh difficulty. 
Potassium Carbonate—K2C03—ordinarily known as pot- 
ashes, is a principal ingredient of plant ashes. The field 
plants absorb potassium salts from the earth ; these are then 
transformed, in them, into salts of organic acids. By the 
combustion of the plants the organic acids are destroyed and 
potassium carbonate produced. The ashes are lixiviated with 274 
water, the filtrate evaporated and the residue ignited. The 
crude potashes thus obtained contain, besides the carbonate, 
also chloride, sulphate and other salts. To purify them, treat 
with a little water, which will dissolve the easily soluble car- 
bonate, leaving the other ingredients, for the most part, be- 
hind. Thus purified potashes are obtained. This method for 
the production of the potash from plant ashes was formerly ex- 
tensively pursued in America, Hungary and Russia; it is not 
much used at present, because, upon the one hand, potassium 
carbonate is, in practice, replaced by the cheaper sodium 
carbonate ; on the other hand, the immense deposits in Stass- 
furt and Galicia afford an inexhaustible source of supply for 
potassium salts. Considerable quantities of potassium car- 
bonate, which, at present, are almost entirely limited to the 
production of Bohemian or crystal glass, have been recently 
obtained from Stassfurt, according to the methods for 
preparing sodium carbonate from the chloride (see Soda). 
Chemically pure potassium carbonate is obtained most conve- 
niently by the ignition of cream of tartar. 
The commercial carbonate is a white, deliquescent powder. 
From concentrated aqueous solution it crystallizes, with 1| 
molecules of water, in monoclinic prisms; at 100° it loses 
i molecule water. The solution has a caustic taste and 
shows an alkaline reaction. When C02 is conducted through 
the liquid it is absorbed and primary jootassium carbonate 
produced:— 
INORGANIC CHEMISTRY 
This salt, ordinarily called bi-carbonate, crystallizes in 
monoclinic prisms, free from water. It dissolves in 3-4 parts 
water and exhibits neutral reaction. Heated to 80°, it decom- 
poses into K2C03C02 and water. 
C03K2 + H20 + C02 = 2 KHC03. 
Potassium Silicate, water-glass, does not possess a constant 
composition and cannot be obtained crystallized. It forms by 
the solution of silicic acid or amorphous silicon dioxide in po- 
tassium hydrate, or by the fusion of silica with potassium hy- 
drate or carbonate. It constitutes a transparent, glassy mass, 
soluble in water. The concentrated solution dries when ex- 
posed, to a glassy, afterwards opaque, mass. Potassium (and 
also sodium) water-glass has an extended application, especi- 
ally in cotton printing, for the fixing of colors (stereochromy), 
in rendering combustible material fireproof, in soap boiling, 
etc. 275 
SULPHUR COMPOUNDS. 
Potassium Sulphydrate, KSH, is obtained when potassium 
hydrate is saturated with hydrogen sulphide:— 
SULPHUR COMPOUNDS. 
Evaporated in vacuo it crystallizes in colorless rhombohedra, 
of the formula, 2KSH -f- ILO, which deliquesce in the air. 
At 200°, it loses its water of crystallization, and at a higher 
temperature fuses to a yellowish liquid, which solidifies to a 
reddish mass. Like the hydrate, it has an alkaline reaction. 
On adding an equivalent quantity of the hydrate to the 
sulphydrate solution, we get potassium sulphide:— 
KOH + H3S = KSH + H20. 
KSH + KOH = K2S + H20. 
Potassium Sulphide, K2S, also obtained by fusing potassium 
sulphate with carbon :— 
K2S04 + 2C = K2S + 2C02. 
Fused, it solidifies to a red crystalline mass. From con- 
centrated aqueous solution it crystallizes with 5 molecules of 
H20, in colorless prisms, which deliquesce in the air. The 
solution absorbs oxygen from the latter, and is decomposed 
into potassium hyposulphite and caustic potash:— 
2K2S + H20 + 20 2 = K2S203 + 2K0H 
Potassium sulphydrate and sulphide precipitate insoluble 
sulphides from the solutions of many metallic salts. By 
acids they are decomposed with liberation of hydrogen 
sulphide. 
When the aqueous solution of the sulphide is boiled with 
sulphur the polysulphides K2S3, K2S4 and K2S5 are formed, 
which after fusion solidify to yellowish-brown masses. The 
aqueous solutions of the polysulphides arc decomposed by 
acids, with disengagement of H2S and separation of sulphur 
(milk of sulphur). The so-called liver of sulphur ( Hepar sul- 
furis), a liver-brown mass, Avhich is used in medicine, is 
obtained by the fusion of potassium carbonate with sulphur, 
and consists of a mixture of potassium polysulphides with 
potassium sulphate. 
The aqueous solution of the potassium, as of the sodium 
sulphide, dissolves some metallic sulphides and forms sulpho- 
salts with them (p. 214). 276 
INORGANIC CHEMISTRY. 
When dry ammonia is conducted over heated potassium, 
potassamide (NH.2K), a dark blue liquid which solidifies to 
a yellowish-brown mass, results. Water decomposes it into 
potassium hydrate and ammonia. When potassamide is 
ignited away from the air ammonia escapes and leaves behind 
potassium nitride, a blackish compound which is spontaneously 
inflammable. 
Recognition of the Potassium Compounds.—Almost all 
the potassium compounds are easily soluble in water, with the 
exception of a few, which, therefore, serve for the character- 
ization and separation of potassium. Tartaric acid added to 
the solution of a potassium salt gives a crystalline precipitate 
of acid potassium tartrate. Platinic chloride (PtCl4) produces 
in solutions a yellow, crystalline precipitate of PtCl4, 2KC1. 
Potassium compounds introduced into the flame of an alcohol 
or gas lamp impart to the same a violet coloration. The 
spectrum of the flame is characterized by bright lines, a red 
and violet (see Spectrum Analysis). 
RUBIDIUM and CiESIUM. 
Kb = 85.4. 
Cs = 132. 
Rubidium and Caesium are the perfect analogues of potassium (p. 
265). They were discovered by means of the spectroscope, by Bun- 
sen and Kirchhoff, in 1860. Although only occurring in small quan- 
tities they are yet very widely distributed, and frequently accompany 
potassium in mineral springs, salt, and plant ashes. The mineral 
lepidolite contains 0.5 percent, of rubidium ; in the very rare pollucite, 
a silicate of aluminium and caesium, upwards of 30 per cent, of caesium 
oxide is present. The spectrum of rubidium is marked by two red 
and two violet lines ; caesium, by two distinct blue lines ; hence, the 
names of these elements. 
With platinum chloride rubidium and caesium form double chlorides 
(PtCl4, 2RbCl), which are more insoluble than the double platinum 
salt of potassium, and, hence, may answer for the separation of these 
elements from potassium. In a free state they may be separated by 
decomposition of the molten chlorides, by means of the galvanic cur- 
rent. Rubidium is also obtained by the ignition of its carbonate with 
charcoal. Metallic rubidium is of a silver white color, with a some- 
what yellowish tinge; its vapor is greenish blue. Csesium has only 
been obtained alloyed with mercury. SODIUM HYDRATE. 
277 
SODIUM. 
Na — 23. 
Sodium is widely distributed in nature, especially as chlo- 
ride in sea water and as rock salt; also in silicates. The 
metal was obtained in 1807, by Davy, by the action of the 
galvanic current upon fused sodium hydrate. At present, 
like potassium, it is obtained upon a large scale by glowing a 
mixture of sodium carbonate and carbon in an iron retort:— 
Na2COs + 2C = 2Na + 3C0. 
The liberated sodium vapors are condensed on flat iron 
receivers of peculiar construction, and the liquefied sodium 
collected under rock-oil. 
Sodium, in external properties, is very similar to potassium. 
It melts at 95.6°, distills at a red heat, and is converted into 
a colorless vapor, which burns with a bright yellow flame in 
the air. It oxidizes readily on exposure, and decomposes 
water even in the cold, although less energetically than potas- 
sium. A piece of sodium thrown upon water swims about 
upon the surface with a rotatory 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. 
Sodium Oxide—Na20, and suboxide, Na40, are very similar 
to the corresponding potassium compounds ; the peroxide is some- 
what different. It is obtained by burning sodium in an oxygen cur- 
rent Its formula is Na202. When heated it absorbs iodine vapors, 
forming the compound Na2OI2 (Na202 + I2 = Na2OI2 + 0), solu- 
ble in water, but decomposed by acids into free iodine and sodium 
salt. This compound, as also some others, seems to indicate that 
sodium has several equivalences. 
It is very remarkable that upon heating sodium oxide it is decom- 
posed by hydrogen, with separation of metallic sodium and formation 
of sodium hydrate :— 
Na20 + H = NaOH + Na. 
This is explained by the fact that the heat of formation of NaOH 
(102 0 C.) is greater than that of Na20 (100.2 C.) and hence the 
reaction occurs, according to the above equation, accompanied by 
heat disengagement. Conversely, therefore, NaOH cannot be decom- 
posed by sodium (Beketoff). 
Sodium Hydrate — NaOH— like potassium hydrate, is 
formed by boiling a solution of sodium carbonate with cal- 
cium hydrate:— 
Na2C03 + Ca (0H)a = CaC03 + 2NaOH. 278 
INORGANIC CHEMISTRY. 
At present it is directly produced in the soda manufacture 
by adding a little more carbon to the fusion (see Soda). 
The sodium hydi’ate which solidifies after the fusion is a 
white, radiating, crystalline mass, and resembles caustic pot- 
ash 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 deliquesces 
perfectly, as the resulting carbonate is also deliquescent). The 
aqueous solution of sodium hydrate resembles that of potas- 
sium. From the concentrated solution separate at 0° crystals 
of NaOH+ 31H20. 
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 %. Rock salt forms large 
deposits in many districts—at Stassfurt and Wielizca in 
Galicia. 
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, 
Eure sodium chloride first separates, and this is collected by itself. 
ater, there crjrstallizes a mixture of the sodium chloride and magne- 
sium sulphate; finally potassium chloride, magnesium chloride and 
some other salts appear (among them potassium iodide and bromide), 
the separation of which constitutes a particular 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 also of other solutions, at first almost pure 
ice separates; the enriched sodium chloride solution is then concen- 
trated in the usual way. 
The rock salt is either mined in shafts or where the strata are not so 
large and are admixed with other varieties of rock, a lixiviation process 
is employed. Borings are made in the earth and water run into them, 
or into any already formed openings. AVhen the water has saturated 
itself with sodium chloride, it is pumped to the surface and the brine 
then further worked up. 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 
these it will be distributed and evaporated by the fall; the concen- 
trated brine collects in the basin below, and is then evaporated over a 
free fire. 
Sodium chloride crystallizes from water in transparent 
cubes, which by slow cooling arrange themselves in hollow, Glauber’s salt. 
279 
four-sided pyramids. It is only slightly more soluble in hot 
than in cold water; 100 parts at 0° dissolve 36 parts salt; at 
100° 39 parts. The saturated solution, therefore, contains 26 
% sodium chloride. The specific gravity of the crystals equals 
2.13. If the saturated solution be cooled below 10°, large 
monoclinic tables (NaCl -f 2H20) separate; these lose water 
at 0° and become cubes. 
The ordinary sodium chloride contains, usually, 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, as 
mechanically enclosed wTater escapes. At red heat the chlo- 
ride fuses and vaporizes at higher temperatures. 
Sodium bromide and iodide at ordinary temperatures crystallize 
with 2 molecules of H20, which they lose again at 30°. 
Sodium chlorate (NaC103) and perchlorate (NaC104) are consider- 
ably more soluble in water than the corresponding potassium salts. 
Sodium Iodate—NaI03—obtained same as the potassium salt, and 
at ordinary temperatures crystallizes with 3 molecules of H20 in silky 
needles. If chlorine gas be conducted through the warmed solution of 
sodium iodate in sodium hydrate, on cooling the periodate 10 | |oH) *2 
(see p. 175) crystallizes out. This, when dissolved in nitric acid, be- 
comes the normal salt (Na I04 + 3H20. 
Sodium Sulphate—Na2S04—crystallizes at ordinary tem- 
peratures with 10 molecules of water of crystallization, and is 
then known as Glauber’s salt (Sal mirabile Glauberi). It 
occurs in many mineral waters, and in large deposits, with or 
without water, in Spain. It is a bye-product in the manu- 
facture of sodium chloride from sea water and brine. It is 
produced in large quantities by heating salt with sulphuric 
acid— 
2NaCl + H2S04 = Na2S04 -f 2HC1, 
and is used in making soda (sodium carbonate). In modern 
times the sulphate has been obtained by a transposition of 
sodium chloride with magnesium sulphate at winter tempera- 
ture—a procedure which is prosecuted‘chiefly in Stassfurt, 
where immense quantities of magnesium sulphate exist:— 
2NaCl + S04Mg = MgCl2 + S04 Na2. 
Sodium sulphate crystallizes at ordinary temperatures with 10 
molecules of H20, in large, colorless, monoclinic prisms, which 
in the air weather and fall into a white powrder. When the 280 
INORGANIC CHEMISTRY. 
salt is heated to 33°, it fuses iu its own water of crystallization; 
by further increase of temperature it gradually loses this, be- 
comes solid, and again fuses at a red heat. The solubility of 
Glauber’s salt (Na2S04 -j- 10 H20) shows thefollowing interest- 
ing deportment: 100 parts of water dissolve, at 0°, 12 parts ; at 
18°, 48 parts ; at 25° 100 parts; at 30°, 200 parts; at 33°, 
237 parts of the hydrous salt. At the last temperature the 
solubility is greatest; by further increase of heat, it gradually 
diminishes; at 50°, 100 parts water dissolve only 263 parts; 
at 1003, 238 parts of the salt. While, ordinarily, the solu- 
bility increases with temperature, Glauber’s salt exhibits 
a varying deportment. This is explained in that the hydrate, 
+10 H20, in aqueous solution, above the tempera- 
ture of 33°, decomposes into w7ater and the salt, Na2S04 + H20, 
which is less soluble in water. The decomposition does not 
occur at once, but only gradually, with increasing tempera- 
ture, for which reason the quantity of the salt dissolved 
gradually grows less. Here we have an example of dissocia- 
tion taking place in aqueous solution. The solution, saturated 
at 33°, becomes turbid upon cooling, and a portion of the dis- 
solved salt separates in anhydrous, small, rhombic octahedra. 
The following interesting deportment in the solution of Glauber’s 
salt may also be noticed. When the solution, saturated at 33°, is 
cooled down to the ordinary temperature, and even lower, not the 
slightest separation of crystals occurs, although the salt is vastly more 
insoluble at lower temperatures than at 33°. Many other salts form 
similar supersaturated solutions, although they are less striking than 
that of Glauber’s salt. The supersaturated solution of the latter may 
be agitated and twirled about without consequent crystallization. If, 
however, a glass rod, or some other solid body, be introduced into 
the solution, it will solidify, suddenly, to a crystalline mass. The 
particles of dust floating about in the air will have a like effect; 
therefore, to preserve the supersaturated solution, the vessel con- 
taining it should be kept well corked. By accurately made in- 
vestigations, it has been determined that the crystallization of the 
supersaturated Glauber salt solution is only induced by contact with 
already formed crystals. These must then be present ever}7where in 
the atmosphere, because only solids exposed to the air, and not purified 
bring about the crystallization. Hence the formation of a crystal of 
Glauber’s salt is always dependent upon the previous existence of a 
similar crystal—just as the production of cells is only caused by cells. 
In the crystallization of a supersaturated Glauber salt solution con- 
siderable heat is disengaged, and the mass increases in temperature. 
This is because the latent heat of all substances in the liquid condition 
is greater than in the solid. At 10°, occasionally, and of their own 
accord, transparent crystals, Na2S04 + 7 H20, separate from the SODIUM HYPOSULPHITE. 
281 
supersaturated solution. Exposed to the air and in contact with solid 
bodies, these crystals are changed to anhydrous sodium sulphate and 
Glauber’s salt. 
This salt is employed in medicine as a purgative. Finds 
extended application technically for the fabrication of glass 
and the preparation of soda. 
The primary or acid sodium sulphate—NaHS04—is 
obtained by the action of sulphuric acid upon the neutral 
salt or upon sodium chloride : — 
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 hydrate or carbonate. The 
secondary sulphite, Na2S03, crystallizes at ordinary temperatures 
with 7 molecules of H20; in the presence of sodium hydrate or by 
warming the solution, it separates in the anhydrous state. The 
primary sulphite—NaHS03—gives up sulphur dioxide in the air, and 
is oxidized to sodium sulphate. 
NaCl + H2S04 = NaHS04 + HC1 
Sodium Hyposulphite—Na2S203—is prepared by boiling 
the aqueous solution of neutral sulphite with flowers of sul- 
phur :— 
Na2S03 -f S = Na2S203. 
It crystallizes with 5 molecules of H20, in large monoclinic 
prisms, dissolves very readily in ivater, and in the air is some- 
what deliquescent. At 56° it melts in its water of crystalliza- 
tion ; loses all water at 100° and decomposes by further heating 
into Na2S04 and Na2S5. When the dry salt is heated in the 
air the polysulphide burns with a blue flame. Acids decom- 
pose the aqueous solution with separation of sulphur and 
evolution of sulphur dioxide:— 
. S203Na2 + 2HC1 =2NaCl + S02 + S +H20. 
Like the sulphate it readily affords supersaturated solu- 
tions. The hyposulphite is used as a reducing agent; chlo- 
rine, bromine and iodine are converted by it into the corres- 
ponding halogen salts: — 
2S203Na2 + I2 = S406Nii2 -f- 2NaI. 
Sodium tctrathionate. 
An iodine solution is instantaneously decolorized by sodium 
hyposulphite ; sulphuric acid and sodium chloride are pro- 
duced. Upon this reaction rests the application of sodium 282 
INORGANIC CHEMISTRY. 
hyposulphite as an antichlor in chlorine bleaching, to remove 
the excess of the chlorine, which has a destructive effect 
upon the tissue. In consequence of its property of dissolv- 
ing the halogen silver derivatives, it is employed in photo- 
graphy. 
Sodium Carbonate (Soda)—Na^COs.—This technically very 
important salt occurs frequently in nature. In some districts, 
like in Hungary and in Africa, it disintegrates from the soil, 
and occurs also in the so-called sodium seas (in Egypt, upon 
the coast of the Caspian Sea). It is contained in the ashes of 
many sea plants, chiefly the algse, 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. The ash of the sea plants (in Normandy called varec, 
in England kelp) formerly served as the principal source of 
the soda manufacture. At present it is, however, almost ex- 
clusively made in large quantities from sodium chloride, 
according to a method devised in 1808 by Leblanc. 
According to this method the sodium chloride is converted, 
by warming with sulphuric acid, into sodium sulphate (p. 279). 
When the latter is dry, it is mixed with charcoal and chalk 
and glowed in a reverberatory furnace. Two principal phases 
may be distinguished in this reaction. First, the carbon re- 
duces the sodium sulphate:— 
Na2S04 + 2 C = Na2S + 2 C02. 
The sodium sulphide then acts upon the calcium carbonate 
to form the calcium sulphide and sodium carbonate:— 
Na2S + CaC03 = CaS + Na2C03. 
At the same time, by the high temperature, a portion of the 
calcium carbonate is changed to oxide and carbon dioxide; 
the appearance of the monoxide, which burns with a bluish 
flame, indicates the end of the action. The chief products in 
the soda fusion are, then, sodium carbonate, calcium sulphide 
and oxide; in addition, different other sulphur salts are 
formed in smaller quantity. The fusion is lixiviated with hot 
water ; the sodium carbonate dissolves, and there remains be- 
hind an insoluble compound of calcium sulphide with oxide, 
CaO, 2CaS, the soda residue. By evaporation of the solution 
and ignition of the residue, we get the commercial or crude 
calcined soda, containing different admixtures, among them SODIUM CARBONATE. 
283 
sodium hydrate. The latter is formed by the action of excess 
of carbon upon sodium carbonate: — 
Na2C03 + C = Na20 -f 2 CO. 
By purposely adding more carbon to the fusion, sodium hy- 
drate is obtained, together with the carbonate. To purify the 
crude soda it is recrystallized from water; large, transparent 
crystals, Na2C03 -f- 10 H20, crystallized soda, separate out; 
in solution remains sodium hydrate. 
Considerable quantities of soda are, at present, obtained from 
cryolite, a compound of aluminium fluoride with sodium fluoride 
(A1F13, 3 NaFl), which occurs in great deposits in Iceland. The pul- 
verized mineral is ignited with burned lime; insoluble calcium fluoride 
and a very soluble compound of aluminium oxide with sodium oxide, 
called sodium aluminate (see Aluminium) are produced: — 
The mass is treated with water and carbon dioxide conducted into the 
liquid, which causes the precipitation of aluminium oxide, and sodium 
carbonate dissolves:— 
2 (A1F1„ 3 NaFl) + 6CaO = 6 CaFl2 + A1203, 3 Na20. 
A1203, 3 Na20 + 3 H20 + 3 C02 = A12(0H)6 + 3 Na2C.03. 
Latterly, a third procedure has appeared. It depends upon the double 
decomposition of a solution of sodium chloride with primary ammonium 
carbonate, by heating, under high pressure :— 
NaCl -f C03(NH4)H = NaHC03 + NH4C1. 
The difficultly soluble primary sodium carbonate separates from 
solution, leaving ammonium chloride dissolved, which can afterwards 
be converted again into carbonate by aid of calcium carbonate. In this 
way, one and the same quantity of ammonium carbonate will suffice for 
the conversion of an indefinite quantity of sodium chloride into soda. 
The technical difficulties which at first opposed the extension of this, 
in chemical respects, so simple a process, are now mostly removed, 
and we can expect that the so-called ammonia process for the soda 
manufacture will replace, at least partially, that of Leblanc. 
At ordinary temperatures sodium carbonate crystallizes 
Avith 10 molecules of H20(Na2C03 + 10H2O) in large mono- 
clinic crystals, which disintegrate upon exposure and become 
a white powder. It melts at 50° in its water of crystalliza- 
tion, and upon additional application of heat a pulverulent 
hydrate—Na2C03 -f- 2H20—separates, which in dry air has 1 
molecule of H20, and at 100° loses all of this. At 30°-50° 
rhombic prisms of the composition C03Na2 -}- 7H20, crystallize 
from the aqueous solution. The anhydrous salt absorbs water 
from the air but does not deliquesce. It melts at a red heat 284 
INORGANIC CHEMISTRY. 
and volatilizes somewhat at a very high temperature. 100 
parts H20 dissolve 10 parts at 0°, and at 38° 138 parts of the 
dry salt. At more elevated temperatures the solubility is less, 
owing, as in the case of the sulphate, to the formation of less 
soluble lower hydrates. Sodium carbonate has a strong 
alkaline reaction; acids liberate carbon dioxide from it. 
Primary Sodium Carbonate—Natrium bicarbonicum— 
NaHC03—is produced by the action of carbon dioxide upon 
hydrous secondary carbonate:— 
Na2C08 + C02 + H20 = 2NaHC03. 
It crystallizes Avithout water, in small monoclinic tables ; dis- 
solves, however, at ordinary temperatures in 10-11 parts water, 
and possesses feeble alkaline reaction. By heating and boil- 
ing the solution it passes into the secondary carbonate with 
disengagement of carbon dioxide. By rapid evaporation small 
monoclinic prisms of the so-called sodium sesqui carbon ate— 
C:)08Na4 -f- 3H20, separate; this also deposits in the sodium 
seas of Hungary and Egypt. 
Sodium Nitrate—NaN03—Chili saltpetre, is found in im- 
mense deposits in Peru. It crystallizes in rhombohedra very 
similar to cubes, hence designated cubic saltpetre. In water 
it is somewhat more easily soluble than potassium saltpetre. 
In the air it attracts moisture, hence is not adapted to the 
manufacture of gunpowder. In other respects it is perfectly 
similar to potassium nitrate. It is largely used in the manu- 
facture of nitric acid, as it is. much cheaper than the 
potassium salt. 
Sodium Phosphates. The sodium salts of phosphoric acid 
are less soluble and crystallize better than those of potassium. 
The tri-sodium phosphate—Na3P04—is made by saturating 1 
molecule of phosphoric acid with 3 molecules NaOH, and 
crystallizes in six-sided prisms with 12 molecules of H20. It 
reacts strongly alkaline, absorbs carbon dioxide from the air, 
and is converted into the secondary salt. 
Di-sodium Phosphate—Na.2HP04—is the most stable of 
the sodium phosphates, and hence, generally employed in 
laboratories (Natrium phosphoricum). It may be obtained 
by saturating phosphoric acid with sodium hydrate to feeble 
alkaline reaction. It crystallizes at ordinary temperatures with 
12H20 in large monoclinic prisms which disintegrate rapidly SODIUM BORATE. 
285 
upon exposure. It is soluble in 4 parts water, and shows a 
feeble alkaline reaction. The solution absorbs carbon dioxide 
abundantly, without suffering any alteration. When heated 
the salt loses wTater, melts about 300°, becoming Sodium Pyro- 
phosphate—Na4P207—which crystallizes with 10 molecules 
of II20, and upon boiling with nitric acid passes into primary 
sodium phosphate. 
The primary or monosodium phosphate—Na H2P04—crys- 
tallizes with 1 molecule of H20 and exhibits an acid reaction. 
At 100° it loses its water of crystallization, and at 200° 
becomes Na2H2P207, disodium pyrophosphate, which at 240° 
forms sodium metaphosphate—NaPOs— 
We get various modifications of the metaphosphate accord- 
ing to the conditions of fusion and cooling; they are probably 
polymerides corresponding to the formulas Na2P206, Na3P309, 
etc. Upon heating sodium metaphosphate with metallic 
oxides the latter dissolve, and salts of orthophosphoric acid 
are formed, e, g.:— 
H2Na2P207 =2NaP03 + H20. 
NaP03 + CaO = NaCaPO*. 
In this manner, with various metals characteristic colored 
glasses (phosphorus beads) are obtained, which in blow-pipe 
analysis serve for the detection of the respective metals. 
The salts of arsenic acid are perfectly analogous to those of phos- 
phoric acid. Of the antimouiates may be mentioned the disodium- 
pyroantimoniate Na2H2Sb207 + 6H20, which is insoluble in cold 
water. 
Sodium Borate. The normal salts of boric acid B(OH)3 
and metaboric acid BO.OH, (see p. 233) are not very stable. 
The ordinary alkaline borates are derived from tetraboric 
acid (H2B407), which results from the condensation of 4 
molecules of the normal boric acid :— 
The most important of the salts is borax, which at ordinary 
temperatures crystallizes w7ith 10 molecules of H20, in large 
monoclinic prisms, Na2 B407 + 10PI2O. Borax occurs naturally 
in some lakes of Thibet, whence it was formerly imported under 
the name of tinkal. At present it is artificially prepared by 
boiling or fusing boric acid with sodium carbonate. At 
ordinary temperatures the crystals dissolve in 14 parts wTater; 
at 100° in one-half part; the solution is feebly alkaline. 
4B(OH)3 — 5H20 = H2B407. 286 
INORGANIC CHEMISTRY. 
When warmed to 70° octaliedra crystallize from the solution 
and have the composition Na2B407 + 5H20, octahedral borax. 
Upon heating, both salts puff' up, lose water and yield a white, 
porous mass (burned borax), which at a red heat fuses to a 
transparent, vitreous mass (Na.2B407). In fusion this dissolves 
many metallic oxides, forming transparent glasses (borax 
beads), which frequently possess characteristic colors ; thus 
copper salts give a blue, chromic oxide, a green glass. 
Therefore borax may be employed in blow-pipe tests for the 
detection of certain metals. Upon this property of dissolving 
metallic oxides depends the application of borax for fusion 
and soldering of metals. 
Sodium Silicate—sodium water glass—is analogous to the 
potassium salt, and is most readily made by fusing quartz with 
sodium sulphate and charcoal. 
The sulphur compounds of sodium are also analogous to 
those of potassium. 
Almost all the sodium salts are easily soluble in water, 
sodium pyroantimoniate—H2Na2Sb207—excepted ; therefore, 
this may be used in precipitating sodium from its salts. 
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. 
RECOGNITION OF SODIUM COMPOUNDS. 
LITHIUM. 
Lithium only occurs in nature in small quantities, hut 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 compound silicate, it occurs in lepi- 
dolite or lithia mica; as phosphate (with iron and manganese) 
in triphylite. 
The metal is separated from the chloride by means of the 
galvanic 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 white light. 
Li = 7. 287 
The lithium salts are very similar to those of sodium, 
closely approach, however, those of magnesium (p. 266). 
AMMONIUM COMPOUNDS. 
Lithium Chloride—LiCl—crystallizes, at ordinary tempera- 
tures, in anhydrous, regular octahedra ; below 10°, however, 
with two molecules.of 1I20, it deliquesces in the air. 
Lithium Phosphate—Li:,P04 + i11,0—and Lithium Car- 
bonate—Li2C03—are with difficulty soluble in water; there- 
fore, precipitated from solutions of lithium salts by sodium 
phosphate or carbonate. By strong ignition the carbonate 
loses carbon dioxide. As regards these two salts, lithium 
approaches the metals of the calcium group (p. 266). Its 
compounds color the flame a beautiful red; the spectrum 
shows an intense red line. 
AMMONIUM COMPOUNDS. 
Upon page 119 we observed that ammonia combines 
directly with acids to form salt-like compounds, which are 
analogous to metallic salts, especially those of potassium. 
The monatomic group, NH4, playing the role of metal in these 
derivatives, is called ammonium, and its compounds ammo- 
nium compounds. The metallic character of the group NH4 
is confirmed by the existence of ammonium amalgam, which, 
as regards its external appearance, is very similar to the sodium 
and potassium amalgams. Ammonium amalgam may be pre- 
pared by letting the galvanic current act upon ammonium 
chloride, NH4C1, viz., by immersing the negative platinum 
electrode into a depression in the ammonium chloride, which 
is filled with mercury. There then separates, just as in the 
decomposition of potassium or sodium chloride —a metal at 
the negative pole—ammonium which forms the amalgam with 
mercury. The amalgam may also be obtained if sodium 
amalgam be covered with a concentrated solution of ammon- 
ium chloride :— 
(Hg + Na) and NH4C1 yield Hg + NH4 and NaCl. 
Sodium amalgam. Ammonium amalgam. 
Ammonium amalgam forms a very voluminous mass with a 
metallic appearance. It is very unstable, and decomposes 
rapidly into mercury, ammonia and hydrogen. 
On dissolving in water, ammonia yields a strong alkaline 
solution ; however, no proofs are present to lead us to 288 
accept the existence of ammonium hydroxide (NH4OH) in 
solution. On the opposite hand, there exist organic derivatives 
of ammonium hydrate, in which the hydrogen of the ammo- 
nium is replaced by hydrocarbon residues; e. g., tetramethyl 
ammonium hydrate—N(CH3)4OH. These are thick liquids, 
of strong basic reaction, which, throughout, are very similar 
to potassium and sodium hydrate. 
INORGANIC CHEMISTRY. 
Ammonium Chloride — NH4C1 — is sometimes found in 
volcanic districts, and formerly vTas obtained by the dry 
distillation of camel’s dung (Sal ammoniacum). At present 
it is prepared, almost exclusively, by saturating the ammo- 
nia water from gas works with hydrochloric acid. The solu- 
tion is evaporated to dryness and the residue heated in iron ves- 
sels, when the ammonium chloride sublimes as a compact, 
fibrous mass. It dissolves in 2.7 parts of cold and one part of 
boiling water, and crystallizes from the solution in small, 
mostly feather-like, grouped octahedra or cubes, of sharp, salty 
taste. When heated, ammonium chloride sublimes without 
melting ; at the same time a dissociation into NHS and HC1 is 
sustained, but these products recombine again to ammonium 
chloride, on cooling. The dissociation is complete at 350°, 
and the vapor density then equals 13 (H =1) correspond- 
ing to that of a mixture of similar molecules, of NH;) 
(8.5) and HC1 (18.2). A like decomposition is sustained by 
the ammonium chloride when its solution is boiled ; ammonia 
escapes and the solution contains some free hydrochloric acid. 
Ammonium Sulphate—(NH4)2S04—is obtained by saturat- 
ing the ammonia water from gas works with sulphuric acid. 
It crystallizes without water in rhombic prisms, soluble in two 
parts of cold and one part of hot water. It fuses at 140°, and 
by further heating decomposes into ammonia, nitrogen, water 
and ammonium sulphite. 
Ammonium Nitrate—NH4N03—is isomorphous with potas- 
sium nitrate and deliquesces in the air. When heated it 
melts, and then decomposes into hyponitrous oxide and water 
(p. 203). 
Ammonium Nitrite—NH4N02—is present in minute quan- 
tities in the air, and results from the action of the electric 
spark upon the latter when moist, and also in the oxidation of 
phosphorus. It may be obtained by the saturation of aqueous 
ammonia with nitrous acid—in a perfectly pure condition, by AMMONIUM niOSPIIATES, 
289 
the decomposition of silver or lead nitrite by ammonium 
chloride. Heat decomposes it into nitrogen and water (p. 107). 
Ammonium Carbonate. The neutral or secondary salt, 
(NII4)2CO;i, separates as a crystalline powder, when ammonia 
gas is conducted through a concentrated solution of the so- 
called sesquicarbonate. In the air it yields up ammonia 
and becomes the primary or acid salt, NH4HC03, which 
when heated to 58°, decomposes into carbon dioxide, am- 
monia and water. 
The ordinarily occurring, commercial, so-called sesquicar- 
bonate of ammonium (C02)2, (NH3)3, ILO, which can be re- 
garded as a compound of primary ammonium carbonate with 
ammonium carbamate, C03 (NH4)H -J- NH2, C02, NH4 (see 
organic chemistry), arises in the decay of many nitrogenous 
hydrocarbons, e. g., the urine, and was formerly prepared by 
the dry distillation of bones, horn, and other animal substances. 
At present it is obtained by heating a mixture of ammonium 
chloride, or sulphate, with calcium carbonate. Then it sub- 
limes as a white, transparent, hard mass, which gives off am- 
monia and carbon dioxide in the air, falling into a white 
powder of primary ammonium carbonate. The latter, ob- 
tained by the weathering of the two first salts, or by saturating 
ammonium hydrate with carbon dioxide, is a white, odorless 
powder, more insoluble in water. In aqueous solution it 
gradually loses carbon dioxide and is changed to secondary 
carbonate. 
Ammonium Phosphates. The most important of these is 
the secondary ammonium-sodium phosphate, P04(NHl) Nall 
-f- 4H20, ordinarily termed salt of phosphorus. It is found 
in guano and decaying urine. It can be obtained by crystal- 
lization of di-sodium phosphate and ammonium chloride :— 
Na,HP04 + NH4C1 = NH4NaHP04 + NaCl. 
It consists of large, transparent, monoclinic crystals. When 
heated it fuses, giving up water and ammonia and forming 
a transparent glass of sodium metaphosphate NaP03 (p. 285). 
It will serve in blowpipe tests for the detection of various 
metals. 
The tertiary ammonium phosphate—(NH4)3P04—separates 
upon mixing concentrated solutions of phosphoric acid and 
ammonia. Upon drying, it loses ammonia and passes into 290 
the secondary salt (NII4)2HP04, which, upon boiling its solu- 
tion changes to the primary salt, P04(NH4)H2. 
INORGANIC CHEMISTRY. 
Ammonium Sulphide—(NH4),S—results upon mixing 1 
vol. H2S with 2 vols. NH3 at —18°. It is a white crystalline 
mass, decomposing, at ordinary temperatures, into NH4HS 
and NH:1. In aqueous solution it is obtained by saturation 
of ammonium hydrosulphide solution with ammonia. 
Ammonium Hydrosulphide.—NH4SH—is produced upon 
conducting hydrogen sulphide into an alcoholic ammonia solu- 
tion. In aqueous solution it is obtained by saturating aqua 
ammonia with hydrogen sulphide. At first, the solution is 
colorless, but assumes a yellow color on standing in contact 
with the air, owing to the formation of ammonium polysul- 
phides—(NH4)2Sn. The so-called yellow ammonium sulphide 
is more simply obtained by the solution of sulphur in the 
colorless hydrosulphide. Both solutions are often employed 
in laboratories for analytical purposes. 
Recognition of Ammonium Compounds. All ammonium 
salts are volatile and decompose upon heating. The alkalies 
and other bases liberate ammonia from them, which is recog- 
nized by its odor and the blue color imparted to red litmus 
paper. Platinum chloride produces in solutions of ammo- 
nium chloride a yellow crystalline precipitate of ammonio- 
platinum chloride PtCl4.2NH4Cl. Tartaric acid precipitates 
primary ammonium tartrate. 
METALS OF THE SECOND GROUP. 
The second group of the periodic system (see table' p. 237) 
comprises chiefly the diatomic metals, forming compounds 
only according to the diatomic type, MeX2, and in their 
entire deportment exhibiting many analogies. Their 
special relations and analogies are more closely regulated by 
the law of periodicity. 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 many properties approaches 
aluminium; just like lithium, attaches itself to magnesium, 
(p. 266). The metals, calcium, strontium and barium, consti- 
tute the second members of the three great periods, are among 
themselves perfectly homologous (p. 234), and according to ALKALINE EARTH METALS. 
291 
their strong basic character attach themselves to the alkali 
metals K Kb and Cs. The members of the second sub-group 
corresponding to them, zinc, cadmium and mercury, belong 
really to the right negative sides of the three great periods. 
They fall in with the heavy metals, are much less basic, 
and resemble the alkaline earth metals only in their com- 
bination forms. In consequence of the double periodicity 
of the three great periods both sub-groups (Ca Sr Ba and 
Zn Cd Hg) exhibit many analogies with magnesium and 
beryllium. 
Calcium. 
GROUP OF THE ALKALINE EARTHS. 
Strontium. 
Barium. 
Ca = 40. 
Sr = 87.2. 
Ba = 130.8. 
The metals of this group are termed alkaline earth metals, 
because their oxides in their properties attach themselves on 
the one side to the oxides of the alkalies, upon the other to the 
real earths (alumina, etc.) In properties they show the 
same gradation as the elements of the potassium group, and 
as regards their atomic weight bear the same relation to each 
other. With increase in atomic weight their chemical energy 
and basicity become greater. Barium decomposes water ener- 
getically, and oxidizes more readily than strontium and calci- 
um. In accord with this we find barium hydrate a stronger 
base; it dissolves tolerably easily in water, does not decompose 
upon ignition and rapidly absorbs carbon dioxide from the air. 
Barium carbonate is also very stable, fuses at a white heat, 
and only disengages a little carbon dioxide. Calcium 
hydrate is much more difficultly soluble in water, and upon 
ignition breaks up into water and calcium oxide; the car- 
bonate also yields up carbon dioxide upon similar treatment. 
In its entire character strontium stands between barium and 
calcium. All these affinity relations find full expression in 
the heat of formation of the corresponding compounds. 
While thus the alkaline earth metals, in free condition and in 
their hydrates, are similar to the alkalies, they essentially distin- 
guish themselves from them by the insolubility of their 
carbonates and phosphates, and still more by their sulphates. 
In water and acids barium sulphate is almost insoluble, while 
that of calcium dissolves in 400 parts water; strontium sul- 
phate occupies a medium position. 292 
INORGANIC CHEMISTRY. 
The metals of this group do not form any volatile compounds and 
their specific heats have not yet been determined. As the determina- 
tion of the vapor densities of the elements or their volatile compounds, 
further the ascertainment of the specific heat of the metals, afford the 
only two direct means for the derivation of the true atomic weights, it 
was allowable to place the atomic weights of the calcium group equal to 
i 
their equivalent weights (Ca = 20, CaCl). But the great analogy of 
their compounds with those of the metals of the magnesium group, for 
instance, their isomorphism, argues with great probability that the 
metals of the group are diatomic, and that the present accepted double 
atomic weights are the true ones (compare p. 251). This conclusion 
at present for calcium is confirmed by the experimental determination 
of its heat capacity. 
CALCIUM. 
Calcium belongs to the class of elements most widely distributed 
upon the earth’s surface. As calcium carbonate (limestone, 
marble, chalk) and the sulphate (gypsum, alabaster), it repre- 
sents immense deposits in all stratified formations. As phos- 
phate it constitutes phosphorite, as fluoride, fluorite, both of 
which are abundant. As silicate it is found in most of the 
oldest crystalline rocks. 
The metal is obtained by the electrolysis of the fused 
chloride; further, by heating calcium iodide with sodium, or 
calcium chloride with sodium and zinc. Although the 
affinity of calcium for oxygen is less than that of the alkalies, 
yet the oxide (also BaO and SrO) cannot be reduced to 
metal by ignition with carbon, iron or sodium—due, probably, 
to the non-fusibility of the oxide. 
Calcium is a yellow, shining metal, of specific gravity 1.55- 
1.6. In dry air it is tolerably stable, in moist it covers itself 
with a layer of hydrate. It decomposes water with consider- 
able energy. It fuses at a red heat, and in the air burns with 
a brilliant yellow light. 
Calcium Oxide—Lime—CaO—may be obtained pure by 
igniting the nitrate or carbonate. On a large scale it is pre- 
pared by burning the ordinary limestone or marble (CaC03) 
in lime-kilns. It is a grayish-white mass, which does not 
fuse even at the highest temperatures. The oxy-hydrogen 
flame thrown upon a piece of lime causes it to emit an 
extremely intense white light (Drummond’s Lime Light). In 
Ca — 40. CALCIUM PEROXIDE. 
293 
the air lime attracts moisture and C02, becoming calcium 
carbonate. Burned lime unites with water with evolution of 
much heat, breaking up into a white voluminous powder of 
calcium hydrate Ca(OH)2—slaked lime. 
When limestone contains large quantities of aluminium, magnesium, 
carbon, or other constituents, the lime from it slakes with difficulty, 
and is known as poor lime, to distinguish it from pure, fat or rich lime, 
which readily becomes a powder with water. 
Calcium Hydrate—Ca(OH)2—slaked lime—is a white, 
porous powder, forming a thick paste, milk of lime, with 
water. It dissolves with difficulty in cold water (1 part in 
160 parts), but still more difficultly in warm water; the solu- 
tion saturated in the cold (lime water) becomes cloudy upon 
warming. It reacts strongly alkaline. In the air it attracts 
carbon dioxide, and forms calcium carbonate. At a red heat 
it decomposes into oxide and water. 
Slaked lime is employed in the preparation of ordinary 
mortar, a mixture of calcium hydrate, water and quartz 
sand. The hardening of the mortar in the air depends 
principally upon the fact that the calcium hydrate com- 
bines with the C02 of the air to form the carbonate; at 
the same time, by the action of the hydrate upon the silicic 
acid of the sand calcium silicate is produced, whereby the 
durability of the mortar increases with time. 
Hydraulic mortar, or cement, is produced by a gentle burn- 
ing of a mixture of limestone or chalk with aluminium silicate 
(clay) and quartz powder. On stirring the powdered burnt 
mass with water it soon hardens, and is not dissolved by 
water. Some naturally occurring limestones, containing up- 
wards of 20 per cent, clay, yield hydraulic cements, without 
any admixtures after burning. The composition of these latter 
is variable, also the process of their hardening ; it depends 
principally, however, upon the formation of calcium and alu- 
minium silicates. 
Calcium Peroxide—Ca02—is precipitated as a hydrate in 
crystalline leaflets, if lime water be added to a solution of 
barium peroxide in dilute hydrochloric acid; it is very un- 
stable. 
The halogen derivatives of calcium, like those of other 
metals, are prepared by the solution of the oxide or carbonate 
in the haloid acids. Also formed by the direct union of cal- 294 
INORGANIC CHEMISTRY. 
cium with the halogens; calcium burns in the vapors of 
chlorine, bromine and iodine. Technically, calcium chloride 
is often obtained as a bye-product, as in the preparation of am- 
monia. 
Calcium Chloride—CaCl2—crystallizes from aqueous solu- 
tion with 6 molecules ofH20,in large, six-sided prisms, which 
deliquesce in the air. In vacuo it loses 4 molecules H20. 
When heated it melts in its water of crystallization, loses 
water, but only after it is exposed above 200° does it become 
anhydrous, when it becomes a white, porous mass. The dry 
salt fuses at a red heat and solidifies to a crystalline mass, 
which attracts water energetically, and may be employed in 
the drying of gases and liquids. The dry calcium chloride 
also absorbs ammonia, forming the compound CaCl2.8NH;j. 
The crystallized hydrous salt dissolves in water with reduc- 
tion of temperature ; by mixing with snow or ice the temper- 
ature is lowered to —48°. Upon fusing the dry chloride in 
the air, it will partially decompose into the oxide and hydro- 
gen chloride. 
Calcium bromide and iodide are very similar to the chlo- 
ride. 
Calcium Fluoride—CaFl2—occurs in nature as fluorite, in 
large cubes or octahedra, also massive. Often discolored by 
impurities. It is found, in sparing quantities, in ashes of 
plants, bones, and the enamel of the teeth. A soluble fluoride 
added to the solution of calcium chloride throws down in- 
soluble calcium fluoride as a white voluminous precipitate. 
The fluoride is perfectly insoluble in water and is only decom- 
posed by strong acids. It easily fuses at a red heat, serving, 
therefore, as a flux in the smelting of ores. When heated it 
shows phosphorescence. 
Calcium Hypochlorite—Ca(C10)2—is not known in a pure 
condition. The so-called bleaching lime, or chloride of lime, 
obtained by conducting chlorine, at ordinary temperatures, 
over slaked lime, contains calcium hypochlorite as active 
principle. 
According to the analogy with the action of chlorine upon 
potassium, or sodium hydrate, the reaction, in the case of cal- 
cium hydrate may be expressed by the following equation:— 
2Ca(OH)2 + 2C12 = Ca(OCl)2 + Ca€l2 + 2H20. CHLORIDE OF LIME. 
295 
From this, chloride of lime would have to be looked upon 
as a mixture of calcium hypochlorite and water. In accord- 
ance with the equation of the reaction, the completely chlori- 
nated chloride of lime must contain 48.9 % chlorine, which is 
never the case, as invariably a portion of the calcium hydrate 
appears to remain unaltered. When chloride of lime is treated 
with water, calcium hypochlorite and calcium chloride dis- 
solve, the hydrate, for the most part, remaining. It was 
thought the atomic constitution, must be ascribed 
to the mixture of calcium hypochlorite and calcium chloride, 
according to similar molecules:— 
Ca(OCl)2 + CaCl2 = 2CaOCl2. 
According to the more recent investigations of Stahlschmidt, the 
active constituent of chloride of lime consists of a basic calcium hypo- 
chlorite, and the action of chlorine upon calcium hydrate 
takes place according to the following equation :— 
From this the completely saturated chloride of lime does not contain 
more than 39 per cent. Cl, which agrees with actual observation. The 
formation of calcium hydrate by the action of water is explained by the 
decomposition of the basic calcium hypochlorite:— 
3Ca(OH)2 + 2C12 = 2 Ca02HCl + CaCl2 + 2H20 
2Ca02 HC1 = Ca(OCl)2 + Ca(OH)2. 
Chloride of lime is a white, porous powder, of a chlorine- 
like odor. The aqueous solution reacts strongly alka- 
line and bleaches. In the air it decomposes, as the carbon 
dioxide of the former liberates hypochlorous acid. Even 
in closed vessels it gradually breaks up, with elimination 
of oxygen ; the decomposition is hastened by 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 expel chlorine from 
chloride of lime, and in just twice the quantity that the hypo- 
chlorite in the chloride contains :— 
When sulphuric acid acts, the calcium chloride present 
participates in the reaction :— 
Ca (C10)a + 4HC1 = CaCl2 + 2H20 + 2Cla. 
Ca(CiO)2 + CaCl2 + 2H2S04 = 2CaS04 + 2C12 + 2H;0. 296 
INORGANIC CHEMISTRY. 
Upon this is founded the application of chloride of lime for 
the production of chlorine in chlorine bleaching and disinfec- 
tion. 
The quantity of chlorine set free by acids from the chloride 
of lime represents its quantity of so-called active chlorine ; 
good chloride of lime should contain at least 25%. 
Upon boiling the aqueous solution of chloride of lime calcium 
chlorate and chloride are produced:— 
On this is based the application of chloride of lime for the produc- 
tion of potassium chlorate (KC103) by a transposition of calcium 
chlorate with potassium chloride. 
When to the solution of bleaching lime a small quantity of cobaltic 
oxide is added, upon warming, a regular stream of oxygen is disen- 
gaged ; 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 contact action of the oxides. 
The reaction is explained, doubtless, in the same way as the action of 
hydrogen peroxide upon certain oxides (see p. 92). The feebly com- 
bined oxygen atom in cobaltic oxide unites with the oxygen of the 
calcium hypochlorite to form free oxygen :— 
3Ca(C10)2 = (C103)2Ca + 2CaCl2. 
Ca (C10)a + 2Co203 = CaCl2 + 202 + 4CoO. 
Cobaltic . 
oxide. 
Cobaltous 
oxide. 
The resulting cobaltous oxide is then again converted by the chlo- 
ride of lime into cobaltic oxide, which acts upon a fresh quantity of 
bleaching lime. 
Calcium Sulphate—CaSO*—is very abundant in nature. 
In anhydrous condition it forms the mineral anhydrite, 
crystallizing in forms of the rhombic system. With two 
molecules of water it occurs as gypsum, in large mono- 
clinic crystals or in granular, crystalline masses (Alabaster, 
etc). Also upon precipitating the soluble calcium salts with 
sulphuric acid, CaS04 + 2H20 separates as a fine crystalline 
powder. Calcium sulphate is only difficultly soluble in water; 
1 part at average temperatures dissolves in 400 parts H20. 
When heated to 200° gypsum loses all its water and becomes 
burnt gypsum, which pulverized and mixed with water forms 
a paste, that in a short time hardens to a solid mass. The 
hardening is dependent upon the reunion of anhydrous cal- 
cium sulphate with 2 molecules of H20. On this depends 
the use of burned gypsum for the production of moulds, fig- 
ures, etc. In case gypsum has been too intensely heated, 
(dead-burnt gypsum) it will no longer harden with water; the 
naturally occurring anhydrite behaves in the same manner. CALCIUM PHOSPHATE. 
297 
Calcium Nitrate—Ca (N03)2—is produced by the decay of 
nitrogenous organic substances in the presence of lime, there- 
fore it frequently is found disintegrated upon walls (in 
cattle stables). From water it crystallizes in monoclinic 
prisms having four molecules of water ; the anhydrous salt 
deliquesces in the air. By the action of potassium carbonate 
or chloride calcium nitrate may be transposed into potassium 
nitre (p. 272). 
Calcium Phosphate. The tertiary phosphate—Ca3 (P04).j 
—is found in slight quantities in the most of the mountain 
rocks. In combination with calcium fluoride it crystallizes as 
apatite. In compact masses, more or less intimately mixed 
with other constituents, it constitutes, as phosphorite, immense 
deposits in Spain, France, Germany and Russia. When 
these minerals disintegrate the calcium phosphate passes into 
the soil and is absorbed by the plants. In the latter it accu- 
mulates chiefly in the seeds and grains. In the animal king- 
dom it is principally found in the bones, the ashes of which 
contain upwards of 85% calcium phosphate. The tertiary 
calcium phosphate is entirely insoluble in water. If disodiutn 
phosphate be added to the aqueous solution of a calcium salt 
and then ammonium hydrate, it will separate as a gelatinous 
precipitate, which, after drying, forms a white amorphous 
powder. In acids, even acetic, it is very readily soluble. 
The secondary calcium phosphate—P04CaH -f- 2H20—is 
sometimes present in guano, in the form of small, shining 
prisms, and separates as an amorphous precipitate if disodium 
phosphate be added to a solution of calcium chloride mixed 
with some acetic acid. 
The primary phosphate—Ca(H2P04)2—is produced by the 
action of sulphuric or hydrochloric acid upon the two first 
phosphates. It is readily soluble in water and deliquesces in 
the air. Heated to 200° it decomposes into pyrophosphate, 
metaphosphoric acid and water:— 
2Ca(H2P04)2 = Ca2P207 + 2HP03 + 3H20. 
On igniting the mixture with charcoal the metaphosphoric 
acid is reduced to phosphorus. In this manner the latter is 
extracted from calcium phosphate. 
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 298 
INORGANIC CHEMISTRY. 
bone meal and pulverized phosphorite were formerly employed. As, 
however, the phosphoric acid is contained in these substances as tri- 
calcium phosphate, which is not easily absorbed by the plants, at 
present the primary phosphate is extensively employed as a fertilizer, 
or, better, the mixture resulting from the action of sulphuric acid upon 
the tertiary salt. Superphosphate is the name applied to the result- 
ing mass.* 
Calcium Carbonate—CaCO:(—is very widely distributed in 
nature. It crystallizes in two crystallographic systems, hence 
dimorphous. In rhombic crystals with the specific gravity 
3.0 it forms aragonite. In hexagonal rhombohedra with 
specific gravity 2.7 it occurs as calcite. Iceland spar, em- 
ployed for optical purposes, is perfectly pure, transparent cal- 
cite. The common calcite, which constitutes immense mountain 
chains, is an amorphous, or indistinct crystalline stratum, and 
usually is mixed with other constituents, like clay. When 
the limestone is granular and crystalline it is termed marble. 
Dolomite, which also constitutes large layers, is a compound 
of calcium and magnesium carbonate, containing generally 
excess of the former. Chalk is very pure amorphous calcium 
carbonate, consisting of the microscopic shells of sea animals. 
Further, calcium carbonate is a regular constituent of all 
plants and animals ; the shells of eggs, of mussels, also 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, amor- 
phous powder, which soon becomes crystalline. In the cold it 
assumes the form of calcite ; upon boiling the liquid it changes 
generally into aragonite crystals. 
In pure water the carbonate is almost insoluble; dissolves 
somewhat in water containing carbon dioxide, as it very prob- 
ably is changed to primary carbonate—Ca(HC03)2. For this 
reason we find calcium carbonate dissolved in all natural 
Avaters. When the solution stands exposed, more rapidly, on 
warming, carbon dioxide escapes and 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 decomposed by 
acids with evolution of carbon dioxide. At a red heat it 
decomposes into CaO and C02. GLASS 
299 
Calcium Silicate—CaSi03—occurs as white, crystalline 
wollastonite. It is also a constituent of most natural silicates 
and of the artificial silicate fusions—of glass. 
G-lass.—The silicates of potassium 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, an 
amorphous, transparent mass, of average fusibility, results ; it is only 
slightly attacked by water and acids—it is glass. To prepare 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 so- 
dium a mixture of sulphates with charcoal can be employed; the car- 
bon reduces the sulphates to sulphides, which form silicates when 
fused with silicon dioxide. 
The following are varieties of glass:— 
Soda glass—a mixture of sodium and calcium silicates—is easily 
fusible, and is employed for window panes and ordinary glass vessels. 
Potash, or Bohemian Glass, also Crown Glass, consists of calcium 
and potassium silicates, is not so easily fused, is harder, and withstands 
the action of water and acids better than soda glass; therefore em- 
ployed in the manufacture of chemical glass ware. 
Glass Crystal, or Flint Glass, is composed of potassium and lead 
silicate. It is not as hard, tolerably readily fused, refracts light 
strongly, and when polished, acquires a clear lustre. On this account 
it is employed for optical purposes (for lenses, prisms) and used in 
ornamental glassware. Strass—a lead glass containing boron tri- 
oxide is used to imitate precious stones. The opaque varieties of 
enamel consist of lead glass and contain insoluble admixtures, as tin 
dioxide and calcium phosphate in the fused glass. 
Ordinary window glass is obtained by the fusion of rather impure 
materials ; in consequence of the presence of ferrous oxide it is 
ordinarily colored green. To remove this coloration, manganese 
peroxide is added to the fusion. It oxidizes a portion of the ferrous 
to ferric oxide, the silicate of which is colored slightly yellow, while 
manganese forms a violet silicate. Both colors, violet and green, 
almost neutralize each other as complementaries. The colored glasses 
contain silicates of colored metallic oxides; chromium and copper 
color green ; cobalt, blue ; cuprous oxide, a ruby red, etc., etc. 
The sulphur compounds of calcium are very much like 
those of the alkalies. Calcium Sulphide—CaS—is most 
readily obtained by heating the sulphate with carbon, and is a 
whitish-yellow mass. By solution in water is obtained 
Calcium Hydrosulphide—Ca(SH)2—which decomposes on 
boiling the aqueous solution. When calcium oxide is ignited 
with sulphur in a closed crucible a yellowish-gray mass, con- 
sisting of calcium polysulphides and sulphate, is obtained. 
Milk of lime boiled with sulphur yields a deep yellow solution 300 
INORGANIC CHEMISTRY. 
of calcium polysulphides. By acids, very finely divided 
sulphur—milk of sulphur—together with evolution of H2S, is 
precipitated from solutions of the polysulphides. If the 
reverse, the addition of a solution of polysulphides to excess 
of dilute acids, be made, hydrogen persulphide will separate. 
STRONTIUM. 
This element is rather rare in nature, and is principally 
found in strontianite (strontium carbonate) and celestite 
(strontium sulphate). Its compounds are very similar to 
those of calcium. 
Sr =- 872. 
The metal is obtained by the electrolysis of fused strontium 
chloride. It is a brass-yellow metal, of specific gravity, 
2.5. In the air it oxidizes and burns, when heated, with a 
bright light. Water decomposes it at ordinary temperatures. 
Of the compounds of strontium we may mention the 
following:— 
Strontium Oxide—SrO—is most readily obtained by glow- 
ing the nitrate. With water it unites, with strong evolution 
of heat, to Strontium Hydrate—Sr(OH)2--which is more 
readily soluble in water than calcium hydrate. From aqueous 
solution it crystallizes with 8 molecules of H20. When ignited 
it decomposes into SrO and H20, but with more difficulty than 
calcium hydrate. 
Strontium Chloride—SrCl2 6H20—crystallizes from 
water in hexagonal tables, which deliquesce in the air; it is 
somewhat soluble in alcohol. 
Strontium Sulphate—SrSO*—is much more difficultly 
soluble in water than calcium sulphate, but not so much so as 
barium sulphate. 
Strontium Nitrate—Sr(N03)2—is obtained by dissolving 
the carbonate in nitric acid, and is readily soluble in water. 
From warm solutions it crystallizes in anhydrous octahedra, 
but from cold with 4 molecules H20, in monoclinic prisms. 
Mixed with combustible substances it colors the flame a 
beautiful carmine red, and for this reason is employed in 
pyrotechny. BARIUM. 
301 
Strontium Carbonate—SrC03—is precipitated from aque- 
ous solutions of strontium salts, as an amorphous, insoluble 
powder, by soluble carbonates. By ignition it breaks up into 
SrO and C02, however, more difficultly than calcium car- 
bonate. 
BARIUM. 
Barium occurs in nature in large masses, as heavy spar or 
barium sulphate, and as witherite (barium carbonate). All 
its compounds are distinguished by their high specific gravity, 
hence the name barium, from ftapbs, heavy. In accordance 
with its general character barium is a stronger basic metal 
than strontium and calcium (p. 291). 
The barium salts are either prepared from the natural 
witherite, by dissolving it in acids, or from heavy spar. The 
latter is almost insoluble in all acids; to obtain the other 
compounds from it, it must first be converted into sulphide. 
For this purpose a mixture of barium sulphate with carbon is 
heated to redness, whereby the sulphate is reduced to sulphide, 
which is soluble in water and readily transposed by acids. 
Metallic barium was first obtained by the electrolysis of the 
fused chloride. The following method is more convenient: 
Sodium amalgam is added to a hot saturated barium chloride 
solution; the sodium displaces the barium, which forms an 
alloy with the mercury. The resulting liquid barium amal- 
gam is kneaded with water, to remove all the sodium, and 
then heated in an hydrogen stream, to volatilize the mercury. 
Barium is a bright yellow metal, of specific gravity 3.6. 
It fuses at a red heat, but does not vaporize. It is rapidly 
oxidized in the air ; it decomposes water very energetically, 
even at ordinary temperatures, like sodium. 
Ba - 136.8. 
Barium Oxide—BaO—is obtained by the ignition of barium 
nitrate. It is a gray, amorphous mass, of specific gravity 
4.0, and fusible in the oxy-hydrogen flame. With water it 
yields the hydrate, with evolution of much heat. 
Barium Hydrate—Ba(OH)2—is precipitated from concen- 
trated solutions of barium salts by potassium or sodium 
hydrate, not, however, by ammonium hydrate. At ordinary 
temperatures it dissolves in 20 parts, upon boiling, in 3 parts 
water. From aqueous solution it crystallizes with 8 mole- 302 
INORGANIC CHEMISTRY. 
cules of H20 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 car- 
bonate. At a red heat it fuses without decomposition and 
solidifies to a crystalline mass. 
Barium Peroxide—Ba02—is produced when barium oxide 
is heated in a stream of air or oxygen, and always contains 
oxide. To purify it, the commercial peroxide is rubbed 
together with water and added to very dilute hydrochloric 
acid, until the latter is almost saturated. To the solution 
containing barium chloride and hydrogen peroxide, is added 
excess of baryta water. Hydrated barium peroxide—Ba02 
4-H20—separates in shining scales, which, upon warming, 
lose water readily and break up into a white powder of 
barium peroxide. The latter, obtained directly from the 
oxide, is a compact, gray mass. 
The peroxide dissolves in dilute acids, with production of 
hydrogen peroxide. Concentrated sulphuric acid sets free 
ozonized oxygen from it. When strongly ignited (above 
400°) it decomposes into barium oxide and oxygen. 
Barium Chloride—BaCl2—crystallizes from aqueous solu- 
tion, with two molecules of H20, 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 octahedra, of the regular system, soluble in 12 parts 
of cold and 3 parts of hot water. It is employed for green 
flames in pyrotechny. 
Barium Sulphate—BaS04—found in nature as heavy spar, 
in rhombic prisms, with a specific gravity of 4.6. Artificially, it 
is obtained by the precipitation of barium salts with sulphu- 
ric acid as a white, amorphous powder, almost insoluble in 
water and acids. Under the name of permanent ivhite, it is 
used as a paint, as a substitute for poisonous white lead, from 
which it is also distinguished by its unalterability. 
Barium Carbonate—BaCO:,—as witherite, occurs in shin- 
ing, rhombic crystals, and is precipitated from barium solu- 
tions by soluble carbonates, as a white, amorphous powder. 
It fuses at a white heat, and loses some carbon dioxide. MAGNESIUM GROUP. 
303 
Barium Sulphide—BaS—is obtained by igniting the sul- 
phate with carbon. It dissolves in water, with decomposition 
into hydrate and hydrosulphide. 
Recognition of the Compounds of the Alkaline Earths. 
The carbonates and phosphates of this group are insoluble in 
water; hence precipitated from aqueous solution of the salts 
upon addition of soluble carbonates and phosphates (of the 
alkalies). The sulphates are also insoluble in acids (only cal- 
cium sulphate is somewhat soluble) ; for this reason they are 
thrown down from acid solutions by soluble sulphates or free 
sulphuric acid; the precipitation is complete, even with cal- 
cium, if alcohol be added to the solution. The hydrates of the 
alkaline earths, which are more or less soluble in water, are 
only precipitated by sodium or potassium hydrate from concen- 
trated solutions. Hydrofluosilicic acid produces, in solutions 
of barium salts, a crystalline precipitate of barium silico-fluo- 
ride, BaSiFl6. 
Very characteristic are the flame colorations produced by 
the volatile compounds; calcium salts impart a reddish 
yellow color; strontium, an intense crimson; barium, a 
yellowish green. The spectra correspond to these flame 
colors. 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, two less distinct, but very characteristic 
lines, an orange and a blue. Finally, the barium spectrum 
consists of several orange, yellow and green lines, among 
which a bright green is particularly prominent. 
In this group are usually included beryllium, magnesium, 
zinc and cadmium. However, these metals do not exhibit 
complete analogy, as clearly seen in the periodic system 
(p. 290). Generally, beryllium, which approaches alumini- 
um, differs, while magnesium is not only similar to zinc and 
cadmium, but also to the alkaline earths, calcium, strontium 
and barium. The similarity with the latter shows itself in 
the basic nature of magnesium, while with zinc and cadmium 
it chiefly consists in isomorphism of compounds. 
METALS OF TIIE MAGNESIUM GROUP, 304 
INORGANIC CHEMISTRY. 
Beryllium and magnesium bear the same relation to Ca Sr 
Ba as lithium and sodium to the metals of the potassium 
group. 
The alkaline character of the alkaline earths, which grad- 
ually diminishes from barium to calcium, becomes almost 
nothing in magnesium and beryllium, which possess the 
lowest atomic weights (see p. 291). Magnesium and berylli- 
um are scarcely capable, even at boiling temperature, of de- 
composing w7ater. Their oxides and hydrates are almost 
insoluble in it; the hydrates decompose, on gentle ignition, 
into oxides and water. Their carbonates are very unstable; 
their chlorides, too, suffer, even on drying, a partial decompo- 
sition into oxide and hydrogen chloride. More decidedly do 
magnesium and beryllium distinguish themselves from the 
alkaline earths by the solubility of their sulphates. The spe- 
cific properties of beryllium and magnesium are maintained 
in zinc and cadmium, which, with the former, constitute a 
natural group. Zinc and cadmium do not decompose water 
at boiling heat; their hydrates are insoluble in it, and little 
stable; their carbonates and chlorides easily undergo decom- 
position ; their sulphates are readily soluble in water. The 
similarity expresses itself further in the isomorphism of most 
of their compounds. Thus, magnesium and zinc sulphates 
crystallize with 7 molecules of H20, in perfectly similar forms. 
If the solution of a mixture of both salts be allowed to crys- 
tallize, we get crystals with variable quantities of zinc and 
magnesium ; the formation of such isomorpkous mixtures in ad 
libitum proportions, is a characteristic indication of the iso- 
morphism of chemically similar compounds. 
The difference between beryllium and magnesium upon the 
one side, and zinc and cadmium on the other, is shown dis- 
tinctly in their specific gravity. While the two first possess 
a low specific gravity (Be — 2.1 Mg — 1-75), zinc and cad- 
mium (with specific gravities 7.2 and 8.6) belong to the so- 
called heavy metals (see p. 246). 
The difference in specific gravity determines, also, many 
differences in chemical character. The light metals (especi- 
ally 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 insoluble in 
water, and, usually, in acids ; in these respects, magnesium 
and beryllium behave like the alkalies, while zinc and cad- MAGNESIUM. 
305 
mium are precipitated by hydrogen sulphide or alkaline sul- 
phides, from solutions of their salts, as sulphides. Further, 
the oxides of the light metals are very stable, and are only 
reduced by carbon, if they are readily fusible (like potassium 
and sodium oxides); the heavy metals, on the other hand, 
are easily separated from their oxides by carbon. Zinc and 
cadmium oxides are reduced by carbon, while those of mag- 
nesium and beryllium are not altered., All these affinity 
relations find closer expression and explanation in the thermo- 
chemical deportments. 
MAGNESIUM. 
Magnesium is 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 isomorphous mixture of 
calcium and magnesium carbonates. Further, it is present 
in most of the natural silicates ; its soluble salts are contained 
in almost all natural waters. 
Mg = 24. 
Metallic magnesium may be obtained by the electrolysis of 
the chloride or by heating the same with sodium. On a large 
scale it is prepared by heating the double chloride of magne- 
sium and sodium with metallic sodium :— 
MgCl 2. NaCl + 2Na = 3NaCl + Mg. 
The fusion is treated with water and the residual magnesium 
purified by distillation. 
Magnesium is a brightly shining, almost silver-white 
metal, of specific gravity 1.75 ; its atomic volume equals 13.7. 
It is tenacious and ductile, and when heated may be converted 
into wire and rolled out into thin ribbons. It fuses at a dark 
red and distills at bright glowing heat. At ordinary temper- 
ature it scarcely oxidizes in the air ; when heated it burns 
with an extremely intense white light, owing to the glowing 
non-volatile magnesium oxide. Magnesium is rich in chemi- 
cally active rays, and for this reason is employed for photo- 
graphing jn dark chambers. Instead of pure magnesium, 
ordinarily its alloy with zinc is employed ; this burns with an 
almost equally bright light. Boiling water is very slowly 
decomposed by magnesium. In dilute acids it dissolves easily 
to form salts; the alkalies do not attack it. 306 
INORGANIC CHEMISTRY. 
Magnesium Oxide—MgO—or magnesia, formed by the 
combustion of magnesium, is ordinarily obtained by ignition 
of the hydrate or the carbonate (magnesia usta). It is a 
white, very voluminous, amorphous powder, which finds appli- 
cation in medicine. The feebly ignited magnesia combines 
with water, with slight generation of heat, to produce magne- 
sium hydrate. 
Magnesium Hydrate—Mg(HO)2—is precipitated from solu- 
tions of magnesium salts, as a gelatinous mass, by potassium 
or sodium hydrate. Dried at 100° it is a white, amor- 
phous powder. In water and alkalies it is almost insoluble; 
moist litmus paper is, however, colored blue. Ammonium 
salts dissolve it tolerably easily, forming soluble double salts. 
Magnesium hydrate attracts carbon dioxide from the air and 
forms magnesium carbonate. By gentle ignition it yields the 
oxide and water. 
Magnesium Chloride—MgCl2—is present in traces in many 
mineral springs. It may be obtained by solution of the car- 
bonate or oxide in hydrochloric acid. By evaporation of its 
solution the salt crystallizes out with six molecules of H20 in 
deliquescent crystals isomorphous with calcium chloride. 
When heated these give up water; at the same time a partial 
decomposition of the chloride into oxide and hydrogen chlo- 
ride takes place:— 
The chlorides of beryllium and zinc, as also those of several 
other metals, behave similarly. To get anhydrous magnesium 
chloride ammonium chloride is added to the solution. The 
double salt MgCl2, NH4C1 + 6H20 is formed. When heated 
this at first loses water, and at 460° throws off ammonium 
chloride; anhydrous magnesium chloride remains. This is a 
leafy, crystalline mass, which fuses easily, and at a red heat 
distills undecomposed; in the air it is very deliquescent. 
Double salts similar to the above are also formed with 
potassium and calcium chloride. The potassium double 
salt—MgCl2, KC1 -|- 6II20—occurs as carnallite, in consid- 
erable deposits, at Stassfurt. 
MgCl2 + H20 = MgO + 2HC1. 
Magnesium Sulphate—MgS04—is found in sea water and 
in many mineral springs. With more or less water it is 
kieserite, which abounds extensively at Stassfurt. At ordi- 
nary temperatures it crystallizes with 7 molecules H20 — 307 
MAGNESIUM PHOSPHATES, 
MgS04 -f- 7H20—in four-sided rhombic prisms readily soluble 
in water (at 0° in 2 parts water). It has a bitter, salt-like taste, 
and serves as an aperient. From solutions heated to 70° it 
crystallizes with 6 molecules of H20; at 0°, however, with 
12 molecules. Heated to 150° these hydrates lose all their 
water of crystallization, excepting one molecule, which escapes 
above 200.° One molecule of water, therefore, in magnesium 
sulphate, is 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 6 molecules of 
H20 in monoclinic prisms, e. g.:— 
The sulphates of zinc and several other metals, as iron, cobalt, and 
nickel, in their diatomic forms, are very similar to magnesium sulphate. 
Their sulphates crystallize with 7 molecules of H20, are isomorphous, 
and contain 1 molecule of intimately combined water. With potas- 
sium and ammonium sulphates they form double salts, crystallizing 
with 6H20, and also isomorphous; e. g.:— 
MgS04, K2S04 + 6H20. 
ZnS04 + 7H20 ZnS04. K2S04 + 6H20 
FeS04 + 7H20 FeS04. K2S04 + 6H20 
The constitution of these double salts may be viewed in the same way 
as that of potassium—sodium sulphate, or of mixed salts of polybasic 
acids. We may suppose that in the given instance the diatomic 
metal unites two molecules of sulphuric acid :— 
/K 
S04 
/Mg + 6H20. 
S04 
XK 
Magnesium Phosphates. The tertiary phosphate (P04)2 
Mg3, in small quantities accompanies the tertiary calcium 
phosphate in bones and in plant ashes. The secondary phos- 
phate, MgHPCh -f- 7H20, is precipitated from the soluble 
magnesium salts, by disodium phosphate (Na2IIP04) as a salt 
difficultly soluble in water. In presence of ammonium salts 
in the same case magnesium-ammonium phosphate, MgNH4- 
P04 -f- 6H20, insoluble in water, is precipitated as a double 
salt. The latter is found in guano, forms in the decay 308 
INORGANIC CHEMISTRY. 
of urine, and sometimes is the cause of the formation of cal- 
culi. The primary salt, H4Mg (P04)2, has not been obtained. 
The magnesium salts of arsenic acid, H3As04, are very 
similar to those of phosphoric acid. Magnesium-ammonium 
arseniate (MgNH4As04 -f- 6II20) is also almost insoluble in 
water. 
Magnesium Carbonate, MgC03, occurs crystallized in 
rhombohedra in nature (isomorphous with calcite), as magne- 
sium spar; in compact masses, as magnesite. Combined with 
calcium carbonate, it forms dolomite, to which, when pure, the 
formula, CaC03, MgC03, is ascribed; ordinarily, however, it 
contains excess of calcium carbonate. On adding sodium or 
potassium carbonate to the aqueous solution of a magnesium 
salt, some carbon dioxide escapes, and a white precipitate 
forms, which consists of a mixture of magnesium carbonate 
and hydrate. If the precipitate be dried at low* temperature, 
a white, voluminous powder, the composition of which generally 
corresponds * to the formula, Mg(OH)2, 3C03Mg -f- 4H20 is 
obtained. 
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 carbonate, MgC03 -f- 3H20, 
separate. These, boiled with water, give up carbon dioxide 
and are again converted into basic carbonate. The naturally 
occurring magnesite sustains no change when boiled. 
With potassium and ammonium carbonate, magnesium 
carbonate yields isomorphous double salts, e. g., MgC03, 
K2C03 + 4H20. 
Of the silicates of magnesium, we may mention olivine 
(Mg2Si04), serpentine (Mg3Si207 -j- 2H20), talc (Si5014Mg4), 
(Si308Mg2 -f- 2H20), sepiolite. The mixed silicates of 
magnesium and calcium are very numerous; to these belongs 
asbestos. 
*This compound is viewed as a basic carbonate, with the following 
formula:— 
co3/'MgOH 
.Mg 
co3<; 
)Mg 
C°3\MgOH 
It is the salt employed in medicine under the name magnesia alba. BERYLLIUM. 
The fixed alkaline hydrates precipitate magnesium hydrate 
from magnesium salts ; the carbonates throw down basic mag- 
nesium carbonate. The precipitates are insoluble in pure 
water and the alkalies, but readily dissolve in solutions of 
ammonium salts. In presence of the latter, neither the alka- 
line hydrates nor carbonates cause precipitation. Disodium 
phosphate precipitates, in presence of ammonium salts, mag- 
nesium-ammonium phosphate, MgNH4P04 + 6H20, insoluble 
in water. 
Recognition of Magnesium Compounds. 
BERYLLIUM. 
Be = 9.3. 
Among the metals of the second group beryllium occupies a position 
similar to that of lithium in the first group; 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 in many respects attaches itself to magnesium, so does 
beryllium approach aluminium Like the latter, it is scarcely at all 
attacked by nitric acid, but dissolves easily in sodium or potassi: 
um hydrate, 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. Beryllium sul- 
phate, like that of aluminium, forms a difficultly soluble double salt 
with potassium sulphate. However, beryllium, in most of its com- 
pounds, stands nearer to magnesium than to aluminium. 
Beryllium occurs but rarely in nature,-principally in beryl, a double 
silicate of aluminium and beryllium—A.l2Be3(Si03)6. Emerald has 
the same composition, and is only colored green by a slight amount 
of chromium oxide. 
Metallic beryllium is obtained by the ignition of the chloride 
with sodium, and is a white ductile metal, of specific gravity 2.1. Its 
specific heat equals 0.4084, the atomic heat is, therefore, 3.8 (p. 248). 
It does not decompose water, even upon boiling. At ordinary tempera- 
ture it does not oxidize in the air. Heated, it burns in the air with a 
very bright light, but only if it be finely divided, not, however (as 
magnesium), in a compact mass. It is readily dissolved by dilute 
hydrochloric and sulphuric acids; also by potassium and sodium 
hydrates. 
Beryllium Chloride—BeCl2—is obtained, like aluminium chlo- 
ride, by ignition of a mixture of beryllium oxide and carbon in a 
stream of chlorine. It sublimes in shining needles, which deliquesce 
in the air. From aqueous solution it crystallizes with four molecules 
of II20 ; upon drying it suffers a decomposition, like magnesium chlo- 
ride. 
The salts of beryllium have a sweet taste, therefore it has, also, 
been called gludnum. From the soluble salts, ammonium hydrate 310 
INORGANIC CHEMISTRY. 
precipitates a white, gelatinous beryllium hydrate, Be(OH)2. This 
dissolves readily in sodium and potassium hydrate, but on boiling, 
is again separated from solution. When heated, the hydrate breaks 
up into water and beryllium oxide, BeO, which is a white, amorphous 
powder, of specific gravity 3.08. Its specific heat equals 0.2471. 
Beryllium Sulphate—BeS04—crystallizes from water at various 
temperatures, with four or seven molecules of H20, of which one is 
more closely combined. With magnesium sulphate it crystallizes in an 
isomorphous mixture. The double salt, S04Be, S04K2, + 3H20, is, 
like the alums, with difficulty soluble in water. 
ZINC. 
Zn = 65. 
The natural compounds of the heavy metals have generally 
a high specific gravity, frequently possess metallic lustre, 
occur usually in the older crystalline rocks in veins, and are 
termed ores. The most important zinc ores are the carbonate 
—ZnC03—the silicate and sphalerite or blende ZnS. The 
principal sources of these ores are in Silesia, England, Belgium, 
Poland and the United States. To get the metal the car- 
bonate or sulphide is converted into oxide by roasting in the 
air, this, then mixed with carbon, is ignited in earthen ware 
cylindrical tubes. In this manner the oxide is reduced :— 
ZnO + C = Zn + CO 
and the liberated zinc distilled off. In the receivers, together 
Avith the fused compact zinc is a gray, pulverulent mass, 
called zinc dust, which consists of a mixture of zinc oxide with 
finely divided metal. This material is used in laboratories as 
a strong reducing agent. 
Metallic zinc has a bluish-white color, and exhibits rough, 
crystalline fracture ; its specific gravity equals 7-7.2. At ordi- 
nary 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 fuses at 412° and distills about 
1000°. 
In moist air it coats itself with a thin layer of basic carbon- 
ate. Heated in the air it burns, with a very intense, bluish- 
white light, to zinc oxide. Compact zinc decomposes water 
only at a red heat; zinc dust, however, acts at ordinary 
temperatures. In dilute acids zinc is readily soluble; in 311 
ZINC SULPHATE. 
potassium or sodium hydrate, as also in ammonia, upon boil- 
ing, it dissolves with liberation of hydrogen. 
Owing to its slight alteration in the air zinc meets with 
extensive application as sheet-zinc for coating statues and in 
architectural adornment, and galvanizing sheet iron. It also 
forms an important constituent of many important alloys, like 
brass and argentan (see these). 
Zinc Hydrate—Zn(OH)2—is precipitated as a white amor- 
phous powder, from aqueous solution, by alkalies, and is 
soluble in excess of the reagent. When heated it decomposes 
into water and zinc oxide. 
Zinc Oxide—ZnO—is usually prepared by igniting the 
precipitated basic carbonate, and as zinc white, is employed as 
a stable white paint. The oxide obtained by burning the 
metal is a white, voluminous, flocculent mass, called flores 
Zinci or Lana philosophica. When zinc oxide is heated it 
acquires a yellow color, which disappears on cooling. 
In nature zinc oxide occurs as zincite, colored by other 
admixtures. 
Zinc Chloride—ZnCl2—anhydrous, is obtained by heating 
zinc in a stream of chlorine, by the evaporation of the solution 
of zinc in hydrochloric acid, and by distillation of zinc sul- 
phate with calcium chloride. It forms a white, deliquescent 
mass, fusing when heated and subliming without decomposi- 
tion. In pure aqueous solution the zinc chloride upon 
evaporation partially decomposes (like magnesium chloride) 
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 basic oxy-chlorides, e. g,, ZnClOH. Zinc chloride 
forms deliquescent double salts with the alkaline chlorides 
e. g., ZnCl2.2KCl. With ammonia it yields various com- 
pounds, of which ZnCl2.lsH3, is characterized by great 
stability. 
Zinc Sulphate—ZnS04—is formed by dissolving zinc in 
sulphuric acid. It is prepared upon a large scale by a gentle 
roasting of zinc blende (ZnS) ; the zinc sulphate is extracted 
by water. At ordinary temperatures it crystallizes from aque- 
ous solution with 7 molecules of H20 (zinc or white vitriol) 312 
INORGANIC CHEMISTRY 
in rhombic crystals, resembling those of magnesium sulphate 
very much. It forms double salts with the alkaline sulphates; 
these contain six molecules of water (p. 307). 
Zinc Carbonate—ZnC03—occurs native as smithsonite in 
hexagonal crystals. Sodium carbonate precipitates basic car- 
bonates of varying composition, from solutions of zinc salts. 
Zinc Sulphide—ZnS—is zinc blende, usually colored brown 
by ferric oxide or other admixtures. Ammonium sulphide 
precipitates it as a white compound, from zinc solutions. Insol- 
uble in water, but readily dissolved by dilute acids, excepting 
acetic; therefore, precipitated by hydrogen sulphide from 
zinc acetate solutions. This reaction serves to separate zinc 
from other metals. 
Zinc Silicate—Zn2Si04+H20—occurs native as calamine. 
CADMIUM 
Cadmium very often accompanies zinc in its ores. In the 
Silesian zinc ores as much as 5 per cent, are present, and in 
these it was discovered in 1819. Being more volatile than zinc, 
in obtaining the latter it first distills off, and may be easily 
separated from the first portions of the distillate. It is a 
white, tenacious and tolerably soft metal, of specific gravity 
8,6. It fuses at 315°, and boils at 860°. In the air it does 
not alter much. Heated, it burns with the separation of a 
brown smoke of cadmium oxide. It is difficultly soluble in 
dilute hydrochloric and sulphuric acids, but dissolves readily 
in nitric. Zinc throws out the metal from the soluble cad- 
mium salt solutions. 
Ca = 111.6 
St. Claire Deville found the specific gravity of cadmium vapors (at 
1040°) to be 3.9 (air — 1) or 56 (H = 1). Therefore, the molecular 
weight of cadmium is 112. Now, as the atomic weight of cadmium 
(determined from its specific heat and the vapor density of the volatile 
compounds) is also 112, it follows that the gas molecule of cadmium 
consists of but one atom. We saw that the molecules of other ele- 
ments in the gaseous state were composed of two or more atoms 
(02,N2,04,S6). Cadmium forms an exception to this rule. Mercury, 
and probably also other diatomic metals, like zinc, deport themselves 
similarly. These relations remind us of the behavior of the hydrocar- 
bon residues (radicals) ; while the diatomic or tetratomic groups, e. <7., 
ethylene C2H4 and acetylene C2H2, exist in free condition, the mona- 
tomic groups (as CH3,CN) cannot appear free, but double themselves, 
if separated from their compounds. CADMIUM. 
313 
Cadmium Hydrate—Cd(OH)2—which is precipitated as a 
white powder, from the soluble cadmium salts, by the alkalies ; 
it is insoluble in sodium and potassium hydrates, but easily 
soluble in ammonium hydrate. 
Of the cadmium compounds may be mentioned :— 
Cadmium Oxide—CdO—is prepared by ignitingthe nitrate. 
A brownish-black powder, consisting of microscopic octahedra. 
Cadmium Chloride—CdCl 2—crystallizes from aqueous so- 
lution, with two molecules of H,0, and may be dried without 
decomposition. The dry salt is fusible and volatile. 
Cadmium Iodide—Cdl2—is obtained by the direct action 
of iodine upon metallic cadmium in presence of water. Crys- 
tallizes from the latter in hexagonal tables. Used in photo- 
graphy. 
Cadmium Sulphate—CdS04—crystallizes from water, not 
like the sulphates of zinc and magnesium, with 7 molecules 
of H.O, but with 3ELO; the crystals weather in the air. 
Yet, with the sulphates of the alkali metals it forms double 
salts, e. {/., CdS04 K2S04 -f 6H20, which are perfectly analo- 
gous to those of zinc and magnesium, and isomorphous with 
them. 
Cadmium Sulphide—CdS—occurs native as greenockite, 
in yellow hexagonal prisms. From cadmium salt solutions 
hydrogen sulphide precipitates it as a yellow powder, insol- 
uble in dilute acids. It is employed as a pigment. 
Almost all the alloys of cadmium have a low fusion tem- 
perature. Cadmium amalgam forms, freshly prepared, a white 
plastic mass, which speedily becomes hard. It is used in fill- 
ing teeth. 
The chemical energy of cadmium is less than that of zinc ; this shows 
itself, among other things, in that cadmium may be displaced from its 
salts by zinc. We saw that, with the elements of the group of potas- 
sium and calcium, the chemical energy was inversely proportional to 
the increasing atomic weight; caesium is more energetic than rubidium, 
barium more than calcium. It is worrhy of remark that almost 
throughout, the elements belonging to the second sub-groups of the 
seven main groups of the periodic system exhibit a similar diminution 
in chemical energy with rising atomic weight. Copper displaces silver ; 
phosphorus is more energetic than arsenic and antimony ; sulphur 
more energetic than selenium and silver ; chlorine sets free or dis- 
places bromine and iodine. 314 
INORGANIC CHEMISTRY. 
The relations of affinity find full expression in the thermo chemical 
phenomena whereby the double periodicity of the great periods and 
the relations of the two sub-groups, Ca Sr Ba and Zn Cd Hg to mag- 
nesium distinctly appear; with the basic character increasing from 
magnesium to barium corresponds the increase of the heat liberation in 
the formation of their compounds ; thus, for example, of the chlorides, 
hydrates and sulphydrates: — 
(Mg, Cl2) = 151.0. 
(Ca,Cl2) = 170.2. 
(Sr,Cl2) =184.5. 
(Ba,Cl2) = 194.2. 
(Mg,0.H20) = 148.9. 
(Ca,0,Aq) = 149.4. 
(Sr,0,Aq) = 157.7. 
(Ba,0,Aq) = 158.2. 
(Mg,S,Aq) - 
(Ca,S,Aq) = 98.3. 
(Sr,S,Aq) = 106 6. 
(Ba,S, Aq) = 107.1. 
That the increase is so slight with the hydrates is explained, probably 
by the decreasing solubility of the same from Ba to Mg, which would 
correspond with an absorption of heat (heat of precipitation). 
The series Mg, Zn, Cd, Hg deports itself differently. In this, corres- 
ponding with the diminution of basicity, the heat disengagement 
becomes successively less:— 
(Mg,Cl2) = 151.0. 
(Zn,Cl2) = 97.2. 
(Cd,Cl2) = 93.2. 
(Hg,Cl2) = 63.1. 
(Mg,0) = 145.0. 
(Zn,0) = 86.4. 
(Cd,0) = 66.4. 
(Hg,0) = 30.6. 
(Mg.S) = 
(Zn,S) = 41.5. 
(Cd,S) = 33.9. 
(Hg,S) = 16.8. 
Comparing these numbers with the quantity of heat which is disen- 
gaged in the formation of aqueous hydrochloric acid (H,Cl,Aq = 39.3), 
we find explained the behavior of the metals toward this acid. All 
metals liberating a greater quantity of heat than 39.3 C., in the forma- 
tion of their chlorides (calculated for 1 equivalent of metal) are in con- 
dition to decompose the acid. To this class belong the majority of the 
metals; only mercury, copper, silver, gold, lead, thallium, and some 
others, set free a less amount of heat, and hence are incapable of de- 
composing dilute hydrochloric acid (see p. 259). 
The slight quantity of heat developed in the formation of hydrogen 
sulphide (S,H2 = 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 precipi- 
tated by hydrogen sulphide from their chlorides, etc. 
If in thermo-chemical equation 
(Me, Cl2.Aq) + (S, H2Aq) = Me, S + 2(H, 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 sul- 
phides) ; in the opposite case the sulphide is decomposed by the dilute 
hydrochloric acid. 
To the group of zinc and cadmium belongs, also, Mercury, 
according to the magnitude of its atomic weight. The relation- 
ship of these three heavy metals exhibits itself in the many 
similarities of the free elements, and also in their compounds. CADMIUM. 
315 
Occupying a similar position in the three great periods (234) 
they are distinguished in a physical point of view by their 
ready fusibility and volatility, which, like the specific grav- 
ities, increase with rising atomic weight (just as with the 
metals of the potassium group, p. 2(^j). 
Zn 
Cd 
Hg 
Atomic weight, 
65 
111.6 
200 
Fusing point, 
412° 
315° 
- 40° 
Boiling point, 
towards 1000° 
860° 
360° 
Specific gravity, 
7.1 
8.6 
13.6 
Also the gradation in the heat of formation of their com- 
pounds (p. 314) clearly indicates that mercury must be 
arranged in a group with cadmium and zinc. 
Like zinc and cadmium, mercury yields compounds of the 
form HgX>, in which it appears diatomic, and which, in 
many respects, are similar to the corresponding compounds of 
zinc and cadmium. Thus mercuric sulphate affords double 
salts with the alkaline sulphates, which crystallize with six 
molecules of H20 (S04Hg, S04K2 -j- 6H20) and are isomor- 
phous wTith the double sulphates of the metals of the magne- 
sium group (p. 307). The similarity, however, limits itself 
only to a few compounds. As in each group the properties 
of it experience, with increasing atomic weight, a gradual 
alteration, so with mercury (with high atomic weight of 200) 
does this become the more evident, as the middle member 
(p. 234), belonging to the third great period, is not known. 
Mercury differs essentially from zinc and cadmium in that, in 
ii 
addition to the compounds of the form HgX2 (mercuric 
compounds), it is also capable of yielding such of the form 
HgX (mercurous compounds), in which it seems to be mon- 
atomic. Here we meet an instance, frequently to be observed, 
that one and the same metal (as with the 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 
is seen that the compounds of one and the same metal, accord- 
ing to different forms or types, frequently are more essentially 
distinguished from one another than the compounds of differ- 
ent elements according to the same type. Thus, the mercu- 
ric compounds (HgX2) are similar to those of zinc and cad- 316 
INORGANIC CHEMISTRY. 
mium, after the same form, while the mercurous compounds 
i i 
HgX exhibit great resemblance to the cuprous (CuX) and 
i 
silver (AgX) compounds, constituted according to a similar 
type. .♦ . 
It shows that the similarity of the compounds is not only 
influenced by the nature of the metals, but frequently, in high 
degree, by the forms or types according to which they are 
constituted (p. 324). 
As above viewed, mercui’y in its ic compounds is a dyad, in the ous 
a monad. According to the theory of constant valence, the mercury 
atom in the om compounds is, however, also diatomic. We suppose 
that the molecules of the same are twice as large, and that in them 
every two Hg atoms form a diatomic group, as seen from the 
following:— 
Hgx 
I >0 
Hg/ 
Mercurous oxide. 
Hg-Cl 
I 
Hg-Cl 
Chloride. 
Hg — N 0 3 
I 
Hg-NOs 
Nitrate. 
Hg\ 
I > 
Hg/ 
Sulphide. 
An experimental decision upon the above has not yet been given 
(p. 318). 
MERCURY. 
Hg _ 200. 
Mercury (Hydrargyrum) occurs in nature principally as 
Cinnabar, more rarely native in form of little drops scattered 
through rocks. Its most important localities are Almaden in 
Spain, New Almaden in California, Idria in Illyria, Mexico, 
Peru, China and Japan. 
The metallurgical separation of mercury is very sifnple. 
Cinnabar is roasted in reverberatory furnaces, whereby the 
sulphur burns to dioxide; the mercury vapors are condensed 
in large chambers. Or, it is distilled with lime or iron from 
iron retorts. Commercial mercury usually contains, dissolved, 
a slight quantity of other metals. To purify it, it is poured 
in a thin stream into a deep layer of sulphuric or dilute 
nitric acid, by which the accompaning tin and lead are more 
easily dissolved than the mercury. Finally the metal is 
distilled out of a small glass retort and passed through 
chamois skin. 
Mercury is the only metal which is liquid at ordinary 
temperatures. At 0° its specific gravity equals 13.59 ; it MERCURY. 
317 
solidifies at 4(J° and crystallizes in regular octahedra ; it 
evaporates somewhat at medium temperatures and boils at 
360°. Its vapors are very poisonous. The specific gravity 
of mercury vapors is 100 (H = 1) or 6.97 (air = 1). 
Therefore its molecular weight is 200 ; as its atomic weight is 
also 200 the molecule is composed, like that of cadmium, of 
only one atom. At ordinary temperatures mercury is not 
altered in the air; near the boiling point, however, it gradually 
oxidizes to red mercuric oxide. Hydrochloric and cold 
sulphuric acids do not act upon mercury; by hot sulphuric 
acid it is changed, with evolution of sulphur dioxide, into 
mercury sulphate. Even dilute nitric acid will readily dis- 
solve it. It combines at ordinary temperatures with the 
halogens and sulphur. 
Mercury dissolves almost all metals (not iron) forming 
amalgams. With potassium and sodium it unites upon 
gentle warming, with production of heat and light. When 
the quantity of potassium and sodium exceeds 3% the alloy is 
solid and crystalline; by less amount it remains liquid. Tin 
amalgam is employed for coating mirrors. 
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 monatomic ; 
we, however, do not know, whether their molecules are not to 
be expressed by the double formula Hg2X2 (p. 315). In 
many respects the ous compounds are similar to the cuprous 
and silver derivatives. The halogen compounds are insoluble, 
and darken on exposure to light. 
In the ic derivatives—HgX2—mercury is diatomic and is 
very much like zinc and cadmium. Thus mercuric sulphate 
forms double salts with the alkaline sulphates, which crystal- 
lize with 6H,0, and are isomorphous with the double sulphates 
of the metals of the magnesium group. The ic compounds 
almost always form, if the substance reacting with the mer- 
cury is in excess; when the opposite is the case, mercurous 
salts result. The ic derivatives, by the addition of mercury 
pass into the ous, e. g., Hg(N03)2 + Hg = Hg2(NOs)2. Oxi- 
dizing agents convert the ous into the ic compounds ; the latter 
are, on the opposite hand, converted by reducing substances to 
the first. 318 
INORGANIC CHEMISTRY. 
The heat of formation of some mercuric compounds corresponds to 
the symbols :— 
That of the corresponding mercurous salts :— 
(Hg, 0) = 30.6 (Hg, Cl2) = 63.1 (Hg,I2) = 34.3 (Hg,S) = 16.8. 
(tlg2,0) = 42.2 (Hg2.Cl2) = 82.5 (Hg2,I2) = 48.3 (Hg2,S) - 
MERCUROUS COMPOUNDS. 
Mercurous Chloride—HgCl or Hg2Cl2—calomel, is an 
amorphous, white precipitate, produced by the addition of 
hydrochloric acid or soluble chlorides to the solution of mer- 
curous salts. Generally it is formed by the sublimation of 
HgCl2 with mercury or a mixture of HgS04, mercury and 
sodium chloride is sublimed:— 
HgS04 + 2NaCl + Hg = Na2S04 + Hg2Cl2. 
It then sublimes as a radiating, crystalline mass, quadratic 
prisms, of specific gravity 7.2. Calomel is insoluble in water 
and dilute acids ; it gradually decomposes when exposed to 
the light, with separation of mercury. Heated, it sublimes 
without fusing. By the action of strong acids it is converted 
into mercuric salts and free mercury. When ammonium 
hydrate is poured over calomel, it blackens and reacts accord- 
ing to the equation — 
The latter compound is viewed as ammonium chloride, in 
which 2H are replaced by Hg2. 
Hg2Cl2 + 2NH3 =NH4C1 + NH2Hg2Cl 
The vapor density of calomel vapors is 117.7 (H = 1), the molecular 
weight, therefore, 235.4, and corresponds to the formula HgCl (235.4). 
It appeal's, however, that its vapors consist of a mixture of mercury 
and mercuric chloride. Such a mixture must have the same density 
as HgCl:— 
HgCl + HgCl = Hg + HgCl 2. 
1 vol. 1 vol. 1 vol. 1 vol. 
The question, whether the mercurous compounds contain one or two 
atoms of mercury, whether, for example, the formula Hg2Cl2 or HgCl 
properly belongs to calomel, is, therefore, not decided by the deter- 
mination of its vapor density. 
Mercurous Iodide—Hgl or Hg2I2—is prepared by rubbing 
together 8 parts of mercury with 5 parts I, or by precipitating 
mercurous nitrate with potassium iodide. It is a greenish 
powder, insoluble in water and alcohol. Light changes it to 
Hgl2 and Hg. MERCURY. 
319 
Mercurous Oxide—Hg20—is black in color and formed by 
the action of potassium or sodium hydrate upon mercurous 
salts. In the light it decomposes into HgO and Hg. 
Mercurous Nitrate—HgN03 or Hg2(N03)2—is produced 
by allowing somewhat djlute nitric acid to act upon excess of 
mercury. It crystallizes in large monoclinic tables. It dis- 
solves readily in water acidulated with nitric aoid; by pure 
water it is decomposed into the acid salt which passes into 
✓ AIT 
solution and becomes the basic salt—Hg2 , which sepa- 
rates as a yellow powder. 
The nitric acid solution of mercurous nitrate oxidizes w’hen 
exposed to the air, and gradually becomes mercuric nitrate; 
this may be prevented by adding metallic mercury to the so- 
lution, whereby the resultant ic salt is again changed to the 
ous state — 
Hg(N03)2 + Hg — Hg2(N03)2. 
Mercurous Sulphate—Hg(S04)—results upon gentle warm- 
ing of excess of mercury with sulphuric acid; it separates as 
a crystalline precipitate, difficultly soluble in Avater, if sul- 
phuric acid be added to a mercurous nitrate solution. It 
fuses upon application of heat, and decomposes into S02, 02 
and Hg. 
Mercurous Sulphide—Hg2S—is precipitated by potassium 
hydrosulphide, as a black compound, from the dilute solution 
of mercurous nitrate. When gently warmed it decomposes 
into HgS and mercury. 
MERCURIC COMPOUNDS. 
Mercuric Chloride—HgCl2—Corrosive sublimate—is pro- 
duced when mercuric oxide is dissolved in HC1, or metallic 
mercury in aqua regia. On a large scale it is obtained by the 
sublimation of a mixture of mercuric sulphate with sodium 
chloride:— 
HgS04 + 2NaCl = HgCl2 + Na2S04 
It crystallizes from water in fine rhombic prisms, and dis- 
solves at medium temperatures in 15 parts, at 100°, in 3 parts 
water ; in alcohol it is still more soluble. Its specific gravity 
is 5.4. It fuses at 260-270°, and boils about 300°. The 
vapor density is 135.5 (H = 1), corresponding to the molecu- 
lar formula HgCl2 = 271. 320 
INORGANIC CHEMISTRY. 
By reducing substances, like S02 and SnCl2, it is changed 
to insoluble mercurous chloride:— 
2HgCl2 + S02 + 2H20 = Hg2Cl2 + H2S04 + 2HC1. 
By stannous chloride, mercurous chloride is at first precipi- 
tated : 2HgCl2 + SnCl2 = Hg2Cl2 -f 2SnCb, which afterwards, 
by excess of the first, is reduced to metallic quicksilver: 
Hg2Cl2 -f SnCl2 = 2IIg + SnCl4. 
Mercuric chloride is greatly inclined to form double salts 
with metallic chlorides, e. g., HgCl2 KC1 + H20. When am- 
monium hydrate is added to its solution, a white, heavy pre- 
cipitate, called ivhite precipitate, NH2HgCl, is throwTn down. 
This compound is regarded as a derivative of ammonium 
chloride, in which two atoms of II are replaced by diatomic 
mercury, and it has been called Mercur-ammonium Chloride. 
It forms with ammonium chloride the compound NH2HgCl 
Is H4C1, the structure of which is expressed by the formula:—■ 
H /NH3C1 
llg\NH3Cl 
Similar mercur-ammonium derivatives are numerous. 
Mercuric Iodide—Hgl2—is formed by the direct union of mer- 
cury with iodine. When potassium iodide is added to a solution 
of mercuric chloride, Hgl2 separates as a yellow precipitate, 
which immediately becomes red. In HgCl2 and KI solution 
Hgl2 is readily soluble ; from alcohol it crystallizes in bright 
red quadratic rhombohedra. Upon warming Hgl2 to 150° it 
suddenly becomes yellow, fuses and sublimes in yellow, shining, 
rhombic needles. On touching these with some solid they 
become red with separation of heat, and are changed into an 
aggregate of quadratic octahedra. Therefore, mercuric iodide 
is dimorphous. 
Mercuric Oxide—HgO—is obtained by continued heating 
of metallic mercury near the boiling point, or by ignition of 
mercurous or mercuric nitrate, and forms a red crystalline 
powder, of specific gravity 11.2. When sodium hydrate is 
added to a mercuric chloride solution, mercuric oxide sepa- 
rates as a yellowT, amorphous precipitate. Both modifications 
become black when heated, yellowish-red, however, on cooling. 
At about 400° mercuric oxide breaks up into mercury and 
oxygen. 
• <J 
Mercuric oxide combines directly with ammonia, to the MERCURY. 
321 
compound 2IIgO.NH3, which on being heated explodes with 
violence. 
Mercuric Nitrate—Hg(N03)a—is with difficulty obtained 
pure, as it is inclined to form basic salts. A solution of it 
may be made by dissolving mercury or mercuric oxide in ex- 
cess of hot nitric acid. On diluting the solution with water 
the basic salt, IIg(N03)a,2Hg0, separates. Boiling with wrater 
changes this to pure HgO. 
Mercuric Sulphate—HgS04—is produced by warming mer- 
cury or its oxide with excess of concentrated sulphuric acid, 
and is an insoluble white crystalline mass, which becomes 
yellow on heating. With a little water it forms the hydrate 
HgS04 -f- H20, which with more water is decomposed into 
sulphuric acid and into the yellow insoluble basic salt 
HgS04.2Hg0 (Turpetum minerale). 
With the alkaline sulphates mercuric sulphate forms double 
salts, e. g., HgS04, K2S04 -J- 6H20, which are isomorphous 
writh the corresponding double salts of the magnesium group 
(p. 307). 
Mercuric Sulphide—HgS—occurs in nature as.cinnabar, 
in radiating crystalline masses, or in hexagonal prisms of red 
color. It is obtained by rubbing together mercury and 
flowers of sulphur with water, or by precipitation of a solution 
of a mercuric salt with hydrogen sulphide it is got as a black 
amorphous mass. If the black sulphide be heated with exclu- 
sion of air it sublimes as a dark red mass of radiating crystal- 
line structure, which is perfectly similar to natural cinnabar. 
A similar conversion of the black modification into the red is 
attained by continued heating of the same to 50° with a 
solution of potassium or ammonium sulphide. The thus 
obtained red mercury sulphide is employed as artificial cinna- 
bar in painting. 
The mercury compounds can readily be recognized by the 
following reactions. On fusion with dry sodium carbonate, 
mercury escapes, which (if the operation be executed in a 
small tube) condenses upon the wall in metallic drops. Tin, 
copper and zinc throw out metallic mercury from its solutions. 
If a pure piece of sheet copper be dipped into the same, mer- 
cury is deposited as a gray coating, which on being rubbed 
acquires a metallic lustre. The mercurous compounds are 
distinguished from the mercuric by their precipitation with 
hydrochloric acid. 322 
INORGANIC CHEMISTRY. 
As regards their atomic weights, copper, silver and gold 
bear the same relation to the alkali group, especially to 
sodium, as zinc, cadmium and mercury to magnesium :— 
COPPER, SILVER AND GOLD. 
Na = 23 
Cu = 63.3 
Ag = 107.6 
Au = 197 
Mg = 24 
Zn — 65 
Cd = 111.6 
Hg = 200 
They occupy an entirely analogous position in the three great 
periods of the periodic system of the elements (p. 234), and 
constitute the transition from the elements of group VIII, 
especially from nickel, palladium and platinum, to the less 
basic elements of group II—zinc, cadmium and mercury :— 
Ni = 58.6 
Pd = 106 
Pt =196.7 
Cu = 63.3 
Ag = 107.6 
Au = 197 
Zn = 65 
Cd = 111.6 
Hg = 200 
This medium position of the three elements to be treated dis- 
tinctly shows itself in their entire physical deportment. 
While the elements of group VIII, with the last members, 
Ni, Pd and Pt, are with difficulty fusible, and not volatile, 
Cu, Ag and Au, as regards their fusibility and volatility, 
form the transition to the readily fusible and volatile ele- 
ments, Zn, Cd and Hg. They take a medium position, also, 
with reference to their coefficients of expansion, their atomic 
volumes, and other physical properties. It is noteworthy that 
the ability to conduct heat and electricity attains its maximum 
in Cu, Ag and Au. 
As the physical properties of the free elements, so, also, 
those of their compounds, and, especially, those depending 
upon the valence of the elements, are determined, in great 
degree, by the position of the latter in the periodic system. 
In accordance with the double periodicity of the system 
of the elements, Cu, Ag and Au, although really belonging 
to the three great groups, attach themselves to group I, 
and especially sodium, just as the immediately following 
elements, Zn, Cd and Hg, arrange themselves with group 
II and magnesium. Corresponding with this, Cu, Ag and 
Au, like Na, yield compounds of the form MeX, in which 
they appear monatomic. Some of these are isomorphous ; 
thus NaCl, CuCl and AgCl crystallize in forms of the regular 
system. So, also, is silver sulphate, Ag2S04, isomorphous with COrPER, SILVER AND GOLD. 
323 
sodium sulphate, S04Na2; likewise some of the other salts of 
these two metals, Cu and Ag, like the alkalies, form so-called 
sub- or quadrant oxides, Xa40, Cu40, Ag40. 
But to these few, so to say, external properties is almost 
limited the similarity of Cu, Ag and Au with sodium. Just as 
the heavy metals, Zn, Cd and Hg, as regards many properties, 
differ from the light metal magnesium, so do the high specific 
gravity metals, Cu, Ag and Au, and in a marked degree, dis- 
tinguish themselves from the light metal sodium. They show 
all the properties belonging to the heavy metals, which, chiefly, 
are marked by the insolubility of the oxides, sulphides and 
many salts. Their differing from sodium is explained by the 
fact that they really belong to the three great periods, and 
only in a slight degree arrange themselves with the alkalies. 
Particularly irregular is the deportment of gold, with the 
high atomic weight, 197, which, in this respect, agrees with 
mercury (p. 314). 
In the compounds constituted according to the form, 
I 
MeX, in which they appear monatomic, Cu, Ag and Au 
exhibit, as regards their physical and chemical properties, 
great resemblance. The chlorides CuCl, AgCl and AuCl, are 
colorless and insoluble in water, soluble, however, in hydro- 
chloric acid, ammonia, the alkaline hyposulphites, etc., and 
furnish perfectly similar double compounds. While silver only 
I 
enters compounds of the form AgX, copper and gold are 
capable of yielding another form; indeed, copper forms, be- 
I II 
sides cuprous, CuX, also cupric, CuX2, derivatives, in 
which it appears diatomic. The latter are much more stable 
than the former, and embrace the most usual copper salts. 
I 
Gold, however, besides furnishing ous, AuX, compounds, 
III 
has ic derivatives, AuX3, in which it appears triatomic. 
While in their ous forms, Cu and Au are analogous to 
silver (and in less degree, Na) the cupric derivatives show a 
great resemblance to those of the metals of the magnesium 
group and other metals, in their diatomic combinations. 
Thus, the sulphates of zinc, magnesium, cupric oxide (CuO), 
ferrous oxide (FeO), nickel ous oxide (NiO), cobaltous oxide 
(CoO), manganous oxide (MnO), are similarly constituted, 
resemble each other, are isomorphous, and form, with the alka- 324 
INORGANIC CHEMISTRY. 
line sulphates, entirely analogous double salts (p. 307). In the 
II 
same way the carbonates (MeC03), the chlorates and bromates 
II 
(MeCl206 + 6H20) and others, are exactly similarly consti- 
tuted and isomorphous. In its ic derivatives, gold exhibits 
III 
some similarity to the aluminium compounds (A1X3), to those 
III 
of indium (InX3) and other metals, in their triatomic combi- 
nations. Here we see, as already with mercury (p. 316), that 
the similarity of the 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 according to different types, each 
series is usually more or less similar to the compounds of 
other metals of like type. In this manner, the resemblance 
of the compounds of the following types shows itself:— 
Na20 
Sodium oxide. 
Ag20 
Silver oxide. 
Cu20 
Cuprous oxide. 
Au20 
Aurous oxide. 
T120 
Thallous oxide. 
MgO 
Magnesium oxide. 
ZnO 
Zinc oxide. 
CuO 
Cupric oxide. 
FeO 
Ferrous oxide. 
HgO 
Mercuric oxide. 
A1203 
Aluminium oxide. 
Fe203 
Ferric oxide. 
Au 2 0 3 
Auric oxide. 
ti203 
Thallic oxide. 
The character of their derivatives varying with the degree 
of combination or valence, shows itself distinctly, as we have 
seen, with chromium, manganese and iron. In the monatomic 
combinations the heavy metals also exhibit a strong, positive 
basic character. Thus silver oxide (Ag20) and thallous oxide 
(T120) 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. In the triatomic 
combination appears the metalloidal character of the metals, 
the acidic of their oxides. In the hydroxyl derivatives of 
aluminium, indium and gold, Al(OH)3, In(OH)3 Au(OH)3, 
hydrogen may be replaced by the alkalies. Their higher 
forms of oxidation show, like those of the metalloids, a pro- 
nounced acid-like character, (as Pb02, Pt02, Cr03, Fe03) 
which is only lessened by a high atomic weight of the metal 
(as in Pb02 and Pt02). In a less, if not an unimportant de- 
gree, is the character of the compound influenced by the posi- 
tion of the elements in the periodic system. According to this 
the properties of the metallic compounds are not only influ- COPPER 
325 
enced by the nature of the metals, but to a high degree by the 
combination forms. These forms of the elements, and particu- 
larly those of the metals, are, however, if not entirely, yet to a 
considerable degree, regulated by the periodic system, as has 
already been declared (see p. 240). 
This connection of Cu, Ag and Au in an analogous group expresses 
itself also in the heat resulting from the formation of their compounds 
of the form MeX :— 
(Na, Cl) = 97.6 
(Cu, Cl) = 32.8 
(Ag, Cl) = 29.8 
(Au, Cl) = 5.8 
(Na2, 0) = 100.2 
(Cu2, 0) = 40 8 
(Ag2, O) = 5.9 
(Au2, 0) = — 
(Na2, S) = 88.0. 
(Cu2, S) = 20.2. 
(Aga, S) = 6.3. 
(Au2, S) = —. 
Here occur relations exactly the same as in the elements of the 
zinc group (p. 314), and from the thermal data perfectly analogous 
conclusions as regards affinitive relations may be deduced. Thus, for 
example, copper is able to decompose concentrated but not dilute 
hydrochloric acid. The heat of formation of some cupric compounds 
equals :— 
(Cu, 0) = 37.1 (Ca, Cl2) = 51.6 (Cu, Cl2f Aq) = 62.7 
(Cu, S, 04) = 182. 
COPPER. 
Cu = 63.3. 
In native condition, copper is found frequently crystallized 
in cubes and octahedra, in large quantities, in America, China, 
Japan, also in Sweden and in the Urals. The most important 
and most widely distributed of its ores are: cuprite (Cu.,0), 
malachite and azurite (basic carbonates), chalcocite (Cu2S), 
and especially chalcopyrite or speckled copper ore (CuFeS2). 
Metallurgy of Copper.—The extraction of copper from its oxygen 
ores is very simple; by ignition of the same with charcoal, metallic 
copper is melted out. The divided ores are first roasted in the air, by 
which means copper sulphide is partly converted into oxide. The 
mass is afterward roasted with sand, silica fluxes and carbon, 
whereby iron sulphide is converted into oxide and passes into the slag. 
By several repetitions of this process the so-called copper stone—a 
mixture of copper sulphide with oxide—is obtained. This is re- 
peatedly roasted and heated, whereby, by the action of copper oxide 
upon the sulphide, metallic copper is obtained : — 
The molten copper is again fused with charcoal, to free it from the 
oxide. 
2CuO + CuS = 3Cu + S02. 326 
INORGANIC CHEMISTRY. 
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 ; in thin 
leaflets it transmits a green light. It is rather soft and 
ductile, and possesses a specific gravity 8.9. It fuses about 
1054° and vaporizes in the oxy-hydrogen flame. In dry air 
it remains unaltered, in moist it is gradually coated with a 
green layer of copper carbonate. On heating it oxidizes to 
black cupric oxide. 
Copper is not changed by dilute hydrochloric or sulphuric 
acids; if, moistened with these, it be exposed to the air, it 
absorbs oxygen and gradually dissolves. It is similarly dis- 
solved by ammonium hydrate. By hot concentrated sulphuric 
acid it is changed to copper sulphate, with evolution of sulphur 
dioxide. In dilute nitric acid it dissolves in the cold, with 
evolution of nitrogen oxide. From the aqueous solutions of 
copper salts, zinc, iron and also phosphorus precipitate 
metallic copper. 
Copper forms two series of compounds, known as cuprous 
and cupric. In the ic compounds copper is diatomic:— 
These are more stable than the ous derivatives; to them 
belong the ordinary copper salts. In many respects they 
resemble the other dyad metals, namely, those of the magne- 
sium group and ous compounds of iron (FeO), manganese 
(MnO), cobalt and nickel (see p. 323). 
The cuprous compounds are the reverse, very unstable, and 
pass rapidly in the air, with absorption of oxygen, into cupric 
compounds. They show some similarity to the mercurous 
derivatives (p. 318), and possess an analogous composition:— 
CuO 
CuC12 
Cu(OH)2 
S04Cu 
CuCl Cul Cu20 Cu2S. 
Oxygen compounds of the same are, however, not known. 
Corresponding to the above formula, copper in the ous compounds, 
like silver, appears to be monatomic. However, it has not been deter- 
mined whether really these formulas express the true molecular size, 
as, owing to the non-volatility of the copper compound, we possess 
no means of fixing the size of the molecules. It is generally supposed 
that the molecules of cuprous compounds correspond to the double 
formula, that also, in them the copper atom is diatomic and that they COPPER. 
327 
contain a diatomic group, formed of two copper atoms, as seen from 
the following formulas :— 
CuCl 
Cu2Cl2 or | 
CuCl 
Cux 
Cu20 or | /O 
CuX 
Cuprous chloride. 
Cuprous oxide. 
Owing to the isomorphism with the compounds of silver we use exclu- 
sively the simpler formulas. 
CUPBOUS COMPOUNDS, 
Cuprous Oxide—Cu20—occurs as cuprite crystallized in 
regular octahedra. Artificially it is obtained by boiling a 
solution of copper sulphate and grape sugar with potassium 
hydrate, when it separates as a crystalline, bright red powder. 
In the air it does not change. In ammonium hydrate it is 
readily soluble. 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, the half of the copper separating as metal:— 
Cu20 + S04H2 = CuS04 + Cu + h2o. 
The hydrate Cu2(OH)2 is precipitated by the alkalies as a 
yellow powder from hydrochloric acid solutions of Cu2Cl2. 
In the air it oxidizes to cupric hydrate. 
Cuprous Chloride—CuCl or Cu2Cl2—is produced by the 
combustion of metallic copper in chlorine gas (together with 
CuCl2), upon conducting HC1 over copper at a moderate 
glowing heat, by boiling the solution of cupric chloride with 
copper (CuCl2 -f- Cu = Cu2Cl2) and by the action of many re- 
ducing substances upon cupric chloride. It is most conveni- 
ently made if through the concentrated copper sulphate solu- 
tion and sodium chloride, sulphur dioxide be passed, when it 
separates as a white, shining powder, consisting of small tetra- 
hedra. In the air it rapidly becomes green, owing to oxygen 
/Cl 
absorption, basic cupric chloride being formed qjj. Cu- 
prous chloride is readily soluble in concentrated hydrochloric 
acid and in ammonium hydrate ; both solutions possess the 
characteristic property of absorbing carbon monoxide. 
Cuprous Iodide—Cul or Cu2I2—is precipitated from soluble 
cupric salts by potassium iodide:— 
CuS04 + 2KI = Cul + K2S04 + I. 328 
INORGANIC CHEMISTRY. 
By extracting the co-precipitated iodine by means of ether 
it is obtained as a gray pywder insoluble in acids. 
Cuprous Sulphide—Cu2S—occurs as chalcocite crystallized 
in rhombic forms. Produced by burning copper in vapor of 
sulphur and by heating the sulphide in a current of hydrogen ; 
after fusion it solidifies in crystals of the regular system. Com- 
bined with silver sulphide, it constitutes the mineral strorn- 
eyerite | S or Cu2S. Ag2S, isomorphous with chalcocite. 
Copper Hydride—CuH or Cu2H2—belongs to the deriva- 
tives of monatomic copper. It separates as a yellow amor- 
phous precipitate, rapidly becoming brown, if a solution of 
copper sulphate be warmed with hypophosphorous acid. At 
60° it decomposes into copper and hydrogen. 
With hydrochloric acid it forms cuprous chloride :— 
Copper suboxide Cu40 or tetrantoxide, corresponds to potassium 
suboxide. On adding an alkaline stannous chloride solution to one 
of copper sulphate, there separates at first cupric hydrate, which is 
further'reduced to cuprous hydrate and then to suboxide. The la!ter 
is an olive-green powder, which easily oxidizes and is decomposed by 
H2S04 in CuS04 and 3Cu. 
CuH + HC1 = CuCl + H2. 
CUPRIC COMPOUNDS. 
The cupric salts when hydrous are generally colored blue or 
green ; when dry they are colorless. 
Cupric Hydrate—Cu(OH)2—separates as a voluminous 
bluish precipitate when sodium or potassium hydrate is added 
to soluble copper salts. When heated, even under water, it 
loses water and is changed to black cupric oxide. 
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 temper- 
ature settles together and acquires a metallic lustre. By 
heating with organic substances their carbon is converted into 
carbon dioxide, and the hydrogen into water, the cupric salt 
being reduced to metal; upon this rests the application of 
cupric oxide in the analysis of such compounds. 
Copper oxide and hydrate dissolve in ammonium hydrate, 
with dark blue color. The solution possesses the power of 
dissolving wood fibre (cotton-wool, linen, filter paper, etc.)— 
Schiveizer's reagent. COPPER. 
329 
Cupric Chloride—CuCl2—is formed by the solution of 
cupric oxide or carbonate in hydrochloric acid. From 
aqueous solution it crystallizes with 2 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 decom- 
posed into chlorine and cuprous chloride. With potassium 
and ammonium chlorides it yields beautifully crystallized 
double salts. Cupric bromide is like the chloride; the 
iodide is not known, as in its formation it at once breaks up 
into cuprous iodide and iodine. 
Copper Sulphate—CuS04 -f- 5H20—cupric sulphate, 
copper vitriol—may be obtained by the solution of copper in 
concentrated sulphuric acid. On a large scale it is produced 
by roasting chalcocite. It forms large blue crystals of the 
triclinic system, which on exposure become somewhat weath- 
ered. At 1003 the salt loses 4 molecules of water; the fifth 
separates above 200°. The anhydrous sulphate is colorless, 
absorbs water very energetically, and again becomes the blue 
hydrous compound. 
Although copper sulphate only crystallizes with 5 molecules 
of H20, it is capable, like the sulphates of the magnesium 
group, of forming double salts with potassium and ammo- 
nium sulphates, which crystallize with 6H20, and are isomor- 
phous with double salts of the metals of the magnesium group. 
Copper sulphate is employed in electro-plating. When its 
solution is decomposed by the galvanic current copper sepa- 
rates at the negative pole and deposits in a regular layer upon 
the conducting objects connected with the electrode. 
Ammonium hydrate added to a copper sulphate solution in 
sufficient quantity to dissolve the cupric hydrate at first pro- 
duced, changes the color of the liquid to a dark blue, from 
which alcohol precipitates a dark blue crystalline mass with 
the composition CuS04, 4NH3 -j- H20. Heated to 150° this 
compound loses water and 2 molecules of NH3 and becomes 
CuS04.2NH3. It is supposed that these compounds are 
ammonium salts in which a part of the hydrogen is replaced 
by copper; they have been designated cuprammonium com- 
pounds, e. g.:— 
/NH8 
S04( -Cu 
xnh/ 330 
INORGANIC CHEMISTRY. 
The other soluble copper salts give similar compounds with 
ammonium hydrate. 
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. When heated it leaves behind 
cupric oxide. 
Copper Carbonates. The neutral salt (CuC03) is not 
known. When sodium carbonate is added to a warm solution 
of a copper salt the basic carbonate separates as a green 
precipitate, CuC03, Cu(OH)2, or occurs in 
nature as malachite, which is especially abundant in Siberia. 
Another basic salt—2C03Cu.Cu(0H)2* is the beautiful blue 
azurite. 
Copper Arsenite—(As03)2Cu—separates as a beautiful, 
bright green precipitate, upon the addition of sodium arsenite 
to a copper solution. Formerly, under the name of Scheele’s 
green,- it was extensively employed as a color, 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, produced 
by hydrogen sulphide in copper solutions. 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, Cu2S. 
Alloys of Copper. Pure copper is very ductile, and may be 
readily rolled and drawn out into a fine wire. It cannot be 
well poured into moulds, because, upon cooling, it contracts 
unequally and does not fill out the moulds. For such pur- 
poses, alloys of copper are employed, which, in addition, 
possess other technically valuable properties. The most 
important of these are :— 
Brass, consisting of two to three parts copper and one part 
ziuc. It has a yellow color, and is considerably harder than 
pure copper. Ordinarily, one to two per cent, of lead are 
.0—CuOH 
*C0( 
x°\ 
)Cu 
y0/ 
po/ 
X0—CuOH SILVER. 
331 
added to the brass, which facilitates its working upon the 
turning-lathe. Tombac contains 15% zinc, and has a gold- 
like color. The alloy of 1 part zinc and 5.5 parts copper 
answers for the manufacture of spurious gold leaf. The 
alloys of copper with tin are called bronze. Most of the mod- 
ern bronzes also contain zinc and lead ; those from Japan, 
gold and silver. The cannon bronze contains 90% copper and 
10% tin ; bell metal has 20%-25% of tin. 
Argentan is an alloy of copper, zinc and nickel (see 
latter.) The so-called Talmi gold consists of 90-95% Cu and 
5-10% aluminium. The German copper coins consist of 95% 
Cu 4% Sn and 1% Zn. 
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 hydrate, with a blue color. When a 
pure piece of iron is introduced into a copper solution, it 
becomes covered wTith 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 and green lines. 
SILVER 
Silver frequently occurs native. Its most important ores 
are Ag2S and various compounds with sulphur, arsenic, 
antimony, copper and other metals. Of rarer occurrence 
are combinations with chlorine (hornsilver AgCl), bromine 
and iodine. Slight quantities of silver sulphide are present 
in almost every galenite (PbS). The principal localities for 
silver ores are America (Chili, Mexico, California), Saxony 
(Freiberg), Hungary, the Altai and Nertschinsk. 
Ag - 107.6. 
Metallurgy of Silver.—The separation of the metal from its ores is 
tolerably complicated and variously effected ; its elaborate description 
belongs to the province of metallurgy At present, the ores containing 
silver and copper are, in Saxony and the Hartz, roasted in a divided 
state and fused with slags rich in silicic acid. In this way, as with 
copper, a copper stone consisting of iron, copper and silver sulphides 
is obtained. This is then oxidized in a furnace ; from the resulting 
mixture of ferric and cupric oxides and silver sulphate (S04Ag2), the 332 
INORGANIC CHEMISTRY. 
latter is extracted by water. From the aqueous solution the silver 
is precipitated by copper. 
Formerly, in Saxony, the separation of the silver was executed 
according to the so-calied amalgamation process. According to this 
the mixture of sulphides is roasted with sodium chloride, whereby 
silver chloride is produced. The divided material is then mixed 
in rotating vessels with iron scraps and water. The iron causes 
the precipitation of the metallic silver from its chloride : — 
To free the metal from various impurities it is dissolved in mercury 
and the liquid amalgam glowed ; mercury distills off and silver re- 
mains. Owing to scarcity of combustible 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 pre- 
sence of water. In this way cuprous chloride is produced, which, with 
silver sulphide, is transposed into silver chloride and cuprous sul- 
phide :— 
2Ag61 + Fe = FeCl2 + 2Ag. 
2CuCl + Ag2S = Cu2S + 2AgCl. 
To get silver from galenite, proceed as follows :—First, metallic 
lead is obtained. In this way all the silver in the ore passes into the 
lead and may be obtained with profit from the latter, even if it does 
not constitute more than per cent of it. To this end, the metallic 
lead is fused and allowed to cool slowly; pure lead first crystallizes 
out, which can be removed by sieves, while a readily fusible alloy of 
lead with more silver remains behind. This method of Pattiiison’s is 
repeated until the residual liquid lead contains 1 per cent, of silver. 
The lead, rich in silver, is subjected to cupellation: it is fused in a 
reverberatory furnace with air access. The bottom of the oven con- 
sists of some porous substance. In this process the lead is changed to 
readily fusible oxide, which partly flows out of side openings from the 
hearth, or is, in part, absorbed by the porous bed; the unoxidized 
silver remains in the cupel in metallic condition. 
The ordinarily occurring silver (work silver) is not pure, 
but invariably contains, in greater or less quantity, copper 
and traces of other metals. To prepare chemically pure 
metal, the work silver is dissolved in nitric acid, and from 
the solution of the nitrates thus obtained, hydrochloric acid 
precipitates the silver as chloride :— 
AgN03 + HC1 = AgCl + HN03. 
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 :— 
Silver is a pure white, brilliant metal, of specific gravity 
10.5. It is tolerably soft and very ductile, and can be drawn 
2AgCl -f~ Zn = ZnCl2 -f~ 2 Ag. SILVER. 
333 
out to a thin wire. It crystallizes in regular octahedra. It 
fuses about 954°, and in the oxy-hydrogen flame is converted 
into a greenish vapor. 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 combining chemically with it; the 
absorbed oxygen escapes again when the metal cools. 
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 sulphuric acid, it dis- 
solves to sulphate:— 
Nitric acid is the best solvent for silver; cold and in dilute 
state it 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% silver and 10% copper; the 
English shillings contain 95% silver. 
2Ag+ 2H2S04 = Ag2S04 + S02 + 2H20. 
Oxygen forms three compounds with silver, but only the 
oxide affords corresponding salts. 
Silver Oxide—Ag20—is thrown out of silver nitrate solu- 
tions as a dark-brown amorphous precipitate, by sodium or 
potassium hydrate. It is somewhat soluble in water, and 
blues red litmus paper. In this, and by the neutral reac- 
tion of the nitrate, the strong, basic, alkaline nature of silver 
and its oxide exhibits itself; the soluble salts of all other 
heavy metals show an acid reaction. When heated, the oxide 
decomposes into metal and oxygen ; at 100° it is reduced by 
hydrogen. . The hydrated oxide is not known ; the moist 
oxide reacts, however, mostly like the hydrate. 
On dissolving precipitated silver oxide in ammonium hy- 
drate, there separates, on evaporation, black crystals (Ag2.0 
2NH3), which, in a dry condition, explode by the slightest 
disturbance. (Fulminating silver.) 
Silver Suboxide—Ag40—corresponding to potassium suboxide, 
is produced by the heating of silver citrate in an hydrogen stream, and 
forms a black, very unstable powder, which decomposes readily into 
silver oxide and silver. 
Silver Superoxide—AgO or Ag202—is formed by passing ozone 
over silver or its oxide, and by the decomposition of the nitrate by 
the galvanic current. It consists of black, shining octahedra, and at 
100° decomposes into Ag20 and oxygen. 334 
INORGANIC CHEMISTRY. 
The salt-like compounds of silver are, corresponding to the oxide 
Ag20, exclusively constructed after the form AgX and termed argentic. 
They are analogously constructed to the cuprous and mercurous deriva- 
tives, and show, especially as regards physical and chemical properties, 
a great resemblance to the former. It would, therefore, be more 
correct to designate them argentous. Compounds of the diatomic 
form AgX2, as of copper and mercury, are not known for silver. If, 
however, the mercurous and cuprous compounds are expressed by 
double formulas (p. 327)— 
CuOl Cu\ 
I I o 
CuCl Cu/ 
HgCl Hg\ 
I X 0 
HgCl Hg/ 
and 
for which their chemical deportment argues, those of silver must 
have analogous formulas— 
AgCl Ag\ 
I I 0 
AgCl Ag/ 
AgN03 
AgNOs 
ascribed them. Then the silver atom will appear diatomic and a 
complete parallelism with copper would be re-established. The non- 
existence of the silver compounds AgX2 is explained by their slight 
stability, just as lead, the perfect analogue of tin (SnX2 and SnX4), 
yields almost exclusively only compounds of the form PbX2. 
Silver Chloride—AgCl—exists in nature as hornsilver. 
When hydrochloric acid is added to solutions of silver salts, 
a white, curdy precipitate separates ; the same fuses at 260° 
to a liquid, which solidifies to a horn-like mass. The chloride 
is insoluble in acids; it dissolves somewhat in sodium chloride, 
readily in ammonium hydrate, potassium cyanide and sodium 
hyposulphite. From ammoniacal solutions it crystallizes in 
large, regular octahedra. Dry silver chloride absorbs 10% of 
ammonia gas, forming a white compound—2AgCl 3NH3—with 
it, which at 38° gives up its ammonia. 
Silver Bromide—AgBr—is precipitated from silver salts 
by hydrobromic acid or soluble bromides. It has a bright 
yellow color, and dissolves with more difficulty than the 
chloride in ammonium hydrate ; in other respects it is per- 
fectly similar to the latter. Heated in chlorine gas it is con- 
verted into chloride. 
Silver Iodide—Agl—is distinguished from the chloride 
and bromide by its yellow color and its insolubility in ammo- 
nia. In hydriodic acid it dissolves readily to Agl. HI, which 
upon evaporation of the solution separates in shining scales. 
Heated in chlorine or bromine gas, it is converted into chlo- SILVER. 
335 
ride or bromide ; conversely are, however, chloride and bro- 
mide of silver, by action of hydriodic acid, converted into 
silver iodide. 
Sunlight, and also other chemically active rays (magne- 
sium light, phosphorus) color silver chloride, bromide and 
iodide, at first violet, then dark black, whereby they are pro- 
bably converted into compounds of the form Ag2X. In such 
an altered condition, they are capable of fixing finely 
divided silver; on this depends their application in photo- 
graphy. 
In photographic work a negative is first prepared. A glass plate is 
covered with collodion (a solution of pyroxylin in an ethereal solution 
of alcohol) holding in solution halogen salts of potassium or cadmium. 
After the evaporation of ether the glass plate is covered with a dry 
collodion layer containing the haloid salts. Now the plate is im- 
mersed in a solution of silver nitrate, whereby haloids of silver are pre- 
cipitated upon the surface. The plate thus prepared is exposed to light 
in the camera obscura, and then, after the action, dipped into a solu- 
tion of pyrogallic acid or ferrous sulphate. These reducing substances 
separate metallic silver in a finely divided state, which is precipitated 
upon the places where the light has acted. Now the plate is intro- 
duced into a solution of potassium cyanide, which dissolves the silver 
salts not affected by the light, while the metallic, unaltered silver 
remains. The negative thus formed is covered at the places upon 
which the light shone, with a dark silver layer, while the places corres- 
ponding to the shadows of the received image are transparent. The 
copying of the glass negative on paper is executed in a perfectly 
similar manner. 
Silver Cyanide—AgCN—is precipitated from silver solu- 
tions by potassium cyanide, as a white, curdy mass, not 
affected by light. It readily dissolves in ammonium hydrate 
and potassium cyanide, forming with the latter the crystalline 
compound AgCN.KCN. The solution of potassium cyanide 
is employed in the galvanic silvering of metals. 
Silver Nitrate—AgN03—is obtained by dissolving pure 
silver in somewhat dilute nitric acid, and crystallizes from its 
aqueous solution in large rhombic tables, isomorphous with 
potassium saltpetre. At ordinary temperatures it is soluble in 
one part water or in four parts alcohol. It fuses about 200°, 
and on further heating decomposes into oxygen and silver 
nitrite AgN02, which at higher temperatures decomposes 
into nitrogen, oxygen and silver. In sunlight it turns black, 
with separation of metallic silver. Organic substances also INORGANIC CHEMISTRY. 
reduce it to metal. Silver nitrate is employed in the cauter- 
ization of 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 thus 
converted into oxide and the unaltered nitrate extracted with 
water. 
Silver Nitrite—AgN02—is precipitated from silver solu- 
tions by potassium nitrite. It crystallizes in needles, diffi- 
cultly soluble in water. 
Silver Sulphate—Ag2S04—is obtained by the solution 
of silver in hot sulphuric acid, and crystallizes in small rhom- 
bic prisms that are difficultly soluble in water. It is isomor- 
phous with anhydrous sodium sulphate. 
Silver Sulphite, Ag2S03, is precipitated as a white, curdy 
mass, if sulphurous acid be added to the solution of the nitrate. 
It blackens in the air and decomposes at 100°. 
S ilver Sulphide, Ag2S, occurs in regular octahedra, as ar- 
gentite. Hydrogen sulphide precipitates black amorphous 
sulphide from silver solutions. By careful ignition in the 
air it is oxidized to silver sulphate. In water and ammonium 
hydrate it is insoluble and dissolves with difficulty in boiling 
nitric acid. 
Silvering.—When silver contains more than 15 per cent, copper it 
has a yellowish color. To impart a pure white color to objects made of 
such silver they are heated, by air access, to redness. The copper is 
thus superficially oxidized, and may be removed by dilute sulphuric 
acid. The surface of pure silver is then polished. 
The silvering of metals and alloys (new silver, argentan) is 
executed in a dry or wet way. In the first, the objects to be silvered 
are coated with liquid silver amalgam, with a brush, and then 
heated in an oven ; the mercury is volatilized, and the silver surface 
then polished 
At present, the galvanic process has almost completely superseded 
the other processes. It depends on the electrolysis of the solution of 
the double cj'anide of silver and potassium, whereby the silver is 
thrown out upon the electro-negative pole and deposits upon the 
metallic surface, in connection with that electrode. 
To silver glass, cover it with a mixture of an ainmoniacal silver solu- 
tion with reducing organic substances like aldehyde, lactic and tar- 
taric acids. Under definite conditions, the reduced silver deposits 
upon the glass as a regular metallic mirror. GOLD. 
337 
Hydrochloric acid throws down a white, curdy precipitate 
from acid silver solutions. Ammonium hydrate dissolves the 
compound readily. Zinc, iron, copper and mercury throw out 
metallic silver from solutions of silver salts, and from in- 
soluble compounds, like the chloride. 
Recognition of Silver Compounds. 
GOLD 
Gold (durum) occurs usually in the native state, and 
is found disseminated in veins in some of the oldest rocks. 
By the breaking and disintegration of these, gold sands are 
formed. In slight quantity it is found in the sand of almost 
every river. Combined with tellurium it forms sylvanite, 
found in Transylvania and California. Contained in minute 
quantity in the most varieties of pyrites and in many lead 
ores. For the separation of the gold grains, the sand or 
pulverized rocks are washed with running water, which 
removes the lighter particles and the specifically very 
heavy gold remains. 
Native gold almost invariably has accompanying it silver, 
copper and various other metals. To free it from these it is 
boiled with nitric or concentrated sulphuric acid. The 
removal of the silver by the latter acid is only complete if that 
metal predominates; in the reverse case a portion of it will 
remain with the gold. Therefore, to separate pure gold from 
alloys poor in silver they must first be fused with about three- 
fourths their weight of the latter. Gold may be separated 
from copper and lead by cupellation (p. 332). 
Pure gold is tolerably soft (almost like lead), and has a 
specific gravity 19.32. It is the most ductile of all metals and 
may be drawn out into extremely fine Avire and beaten into 
thin leaves, which transmit green light. About 1100° it 
melts to a greenish liquid. Not altered by oxygen even upon 
glowing, and is not attacked by acids. Only in a mixture of 
nitric and hydrochloric acids (aqua regia) which yields free 
Au = 196.2 * 
*The atomic weight of gold was last (1850) found, by Levol, to be 
196.2. In accordance with its position in the periodic system it is, 
however, very probable that its atomic weight is somewhat greater than 
that of platinum (196.7). 338 
INORGANIC CHEMISTRY. 
chlorine, is it dissolved to gold chloride AuC13. Free chlorine 
behaves similarly. Most metals and many reducing agents 
(ferrous sulphate, oxalic acid) precipitate gold from its 
solutions as a dark brown powder. 
Gold being very soft and readily extracted, is, for use, gener- 
ally 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 English gold coins contain 90% gold and 10% copper. 
For ornamental objects a 14-karat gold is generally 
employed. This contains about 58% pure gold. 
Gold, according to its atomic weight, belongs to the group 
of copper and silver, and, upon the other hand, forms the 
transition from platinum to mercury, by which double rela- 
tion its character is, in an important measure, determined 
(p. 322). Having the high atomic weight 196-197, it ex- 
hibits, like the other elements of high atomic weight belonging 
to the 10th horizontal series of the periodic system—mercury, 
thallium, lead and bismuth—a deportment more varying, from 
the lower analogues;—like the elements of the 1st horizontal 
series (lithium to fluorine) which have the lowest atomic 
weights, and also differ somewhat from the other members of 
their groups. 
Like silver and copper, gold yields compounds of the form 
AuX—aurous, analogous to the cuprous and argentous. In 
addition, it has those of the form AuX3, auric derivatives, in 
which it is triatomic. These show the typical character of 
the triatomic combination form, which expresses itself in the 
acidity of the hydrates (p. 324) ; auric hydrate, Au(OII)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 (PtX4), mercury 
(HgX2), thallium (T1X3), and lead (PbX4). 
AUROUS COMPOUNDS. 
Aurous Chloride—AuCl—is produced by heating auric 
chloride, AuC13, to 150°, and forms a white powder insoluble 
in water. When ignited, it decomposes into gold and chlo- 
rine : boiled with water it decomposes into the trichloride and 
gold. GOLD. 
339 
Aurous Iodide—Aul—separates as a yellow powder, if po- 
tassium iodide be added to a solution of auric chloride. 
AuC13 + 3KI = Aul + I2 + 3KC1. 
When warmed it breaks up into gold and iodine. 
When auric oxide or sulphide is dissolved in potassium 
cyanide, upon evaporation, large colorless prisms of the double 
cyanide, AuCN. KCN, crystallize out. The galvanic current 
and many metals precipitate gold from this compound; hence 
it serves for electrolytic gilding, which, at present, has almost 
entirely superseded the gilding in the dry way (see p. 336). 
Aurous Oxide—Au20—is formed by the action of potas- 
sium hydrate upon aurous chloride. It forms a dark violet 
powder which at 100° decomposes into gold and oxygen. By 
the action of hydrochloric acid it is changed to AuC13 and 
gold. 
Of the oxygen derivatives of monatomic gold but a few 
double salts are known. 
AURIC COMPOUNDS. 
Auric Chloride—AuC13—results by the solution of gold 
in aqua regia, and by the action of chlorine upon the metal. 
By evaporation of the solution it is obtained as a reddish- 
brown, crystalline mass, which rapidly deliquesces in the air. 
It is readily soluble in alcohol and ether. 
With many metallic chlorides gold chloride forms beauti- 
fully crystallized double salts, e. g., AuC13, KC1 + 2IH20 and 
AuC13 NH4C1 + H20. When 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 (Au203). This is a brown powder 
wThich, near 250°, decomposes partly into gold and oxygen. 
If the precipitate containing the magnesia be treated, not 
with concentrated, but with dilute nitric acid, Auric Hydrate 
—Au(OHq3—remains as a yellowish-red powder. Both the 
oxide and hydrate are insoluble in water and acids; they 
possess, however, acid properties, and dissolve in alkalies. 
Therefore the hydrate is also called auric acid. Its salts, the 
aurates, are constituted according to the formula MeAu02, 
and are derived from the meta-acid HAuO, = HO.AuO. 340 
INORGANIC CHEMISTRY. 
Potassium Aurate—KAu02 -f- 3II20—crystallizes in bright 
yellow needles, from a potassium hydrate solution of auric 
oxide. These are readily soluble in water; the solution 
reacts alkaline. From this solution the corresponding aurates 
are precipitated by many metallic salts, e. g.— 
The precipitate produced by magnesia in a solution of auric 
chloride (see above) consists of magnesium aurate (Au02)2Mg. 
Oxygen salts of auric oxide are not known. 
KAu02 -(- AgN03 = AgAu02 + KNOs. 
Auric Sulphide—Au2S3—is precipitated as a blackish-brown 
compound, from gold solutions, by hydrogen sulphide. It 
dissolves in alkaline sulphides with formation of sulphosalts. 
Stannous chloride (SnCl2) added to an auric chloride solu- 
tion causes, under definite conditions, a purple-brown precipi- 
tate, purple of Cassius, which is employed in glass and porce- 
lain painting. It probably consists of a mixture of aurous 
stannate and stannous oxide. 
On pouring ammonium hydrate over auric oxide a brown 
compound is produced—fulminating gold. 
To group III of the periodic system (p. 237) belong the 
triatomic elements affording: derivatives chiefly of the forms 
MeX3:— 
METALS OF GROUP III. 
B = 11 A1 = 27.3 
Sc = 45 Y= 89 La =139 Yb = 173 
Among themselves, these show perfectly similar relations, 
as do the elements of group II (p. 290). In boron, with the 
lowest atomic weight, the metallic basic character is very 
much diminished or does not appear. According to its exclu- 
sively acidic hydroxide B(OH)3, it approaches the metalloids, 
hence treated with them, although it does not afford a volatile 
hydride (p. 231). 
Aluminium is a perfect metal; its hydrate AlfOH)3 exhibits 
a predominating basic character, and with acids yields salts. 
Its relations to boron are like those of silicon to carbon, or 
of magnesium to beryllium. The connection of aluminium and 
boron with one group plainly shows itself in the entire char- 
acter of the free elements, as also of their compounds. Thus 
Ga =70 In= 113 
T1 = 204. GOLD. 
341 
aluminium and boron are not dissolved by nitric acid, but, 
indeed, upon boiling with alkalies:— 
A1 + 3K0H = Al(OK), + 3H. 
The difference between the hydrates is only a gradual one. 
As boron hydrate B(OH)3 represents a feeble acid, so does 
aluminium hydrate present an acidic-like character, as it is 
capable (p. 324) of forming metallic salts with strong bases 
(chiefly the alkalies) ; owing to its higher atomic weight the 
basic character with aluminium exceeds the acidic. The 
similarity is also shown by the existence of perfectly analo- 
gous compounds ; thus, e. g., the chlorides BC13 and A1C13 can 
unite with PC15 and POCl3. 
To aluminium attach themselves, as first sub-group: scan- 
dium, yttrium, lanthanum and ytterbium. These constitute 
the third members of the great periods, and hence, exhibit a 
pronounced basic character. As light metals, they are, in 
their compounds, very similar to aluminium, so that they all 
are embraced in one group, which (corresponding to the 
earthy nature of their oxides) is designated the Group of 
Earth Metals. In peculiar relation to lanthanum stand cerium 
and didymium, with closely agreeing atomic weights and very 
similar properties. The apparent abnormal existence of the 
same is explained by the fact that the 5th but very imperfect 
period (series 7 and 8) in its middle members possesses a some- 
what varying function (p. 238). Here too may be classed, 
very probably, the but recently discovered, and as yet little 
characterized metals—erbium, terbium, phillipium, thulium 
and decipium. 
More distinctly characterized and accurately investigated 
is the second sub-group, consisting of the heavy metals, gal- 
lium, indium and thallium. These belong to the right side of 
the great period, possess, therefore, a less basic character, and 
bear the same relation to each other as Zn,*Cd and Hg. 
Formerly, aluminium was classed together with chromium, iron, 
manganese, cobalt and nickel, in one group, as they all afford sesqui- 
oxides, Me203, whose salts are very much alike. Particularly proof 
bearing for this classification was thought to be the existence of the 
perfectly similarly constituted alums : — 
(S04)3A12.S04K2 + 24H20 
Potassium aluminium sulphate. 
(S04)3Fe2.S04K2 + 24H20. 
Potassium irou alum. 
Yet in its entire behavior, aluminium is distinguished very essentially 
from the other here mentioned metals—by the acid nature of its 342 
INORGANIC CHEMISTRY. 
hydrate, Al(OH)3—and by its inability to form higher or lower com- 
bination forms, while the others yield basic monoxides, MeO, and 
acid-forming trioxides (Cr03, FeOs, Mn03). The similarity of the 
sesquioxide compounds, Me203 is, therefore, here, like those of the 
monoxide compounds, MeO, to be regarded as chiefly influenced by 
the similarity of the combination forms. 
The, at present, accepted triatomic nature of aluminium apparently 
contradicts the circumstance that not the simple formulas, A1C13, 
AlBr3, but the double ones, A12C16, Al2Brfi, fall to its halogen deriv- 
atives (the result of vapor density determinations). On the other hand, 
however, exist the so-called metallo-organic compounds of aluminium, 
whose molecules are constituted according to the formulas A1(CH3)3, 
A1(C2H5)3 ; from them, undoubtedly, follows the trivalence of alu- 
minium, as the compounds with hydro-carbons (like those with hydro- 
gen) afford the surest guide for the deduction of the valence (p. 239). 
The existence of the molecules A12C16, Br2Cl6, etc., does not prove 
anything against its being a triad, but must be explained by a poly- 
merization of the simple chemical molecules A1C13, BrCl3—as with 
arsenious oxide, As203, and antimony trioxide, Sb203, the molecules 
in vapor form correspond with the doubled formulas As406 = 
(As203.As203) and Sb406. 
ALUMINIUM. 
This is one of the most widely distributed elements. 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 
feldspar, mica, and also in most crystalline rocks. It occurs, 
too, united with fluorine and sodium, as cryolite, in large de- 
posits, in Iceland. 
Metallic aluminium is obtained by igniting the chloride, or 
better, the double chloride-of sodium and aluminium with 
metallic sodium:— 
A1 = 27.3. 
A1C13. NaCl + 3Na = A1 + 4NaCl. 
It is a silver white metal of strong lustre, is very ductile, 
and may be drawn out into thin wire and beaten into thin 
leaflets. Its specific gravity is 2.56; it belongs, consequently, to 
the light metals and possesses, therefore, all the properties 
opposed to those of the heavy metals (see p. 305). It fuses at 
a red heat but will not vaporize. At ordinary temperatures, 
as also by heat, it changes very little in the air; if, however, 
thin leaves be heated in a stream of oxygen they burn with a ALUMINIUM. 
343 
bright light. Nitric acid does not affect aluminium ; sul- 
phuric acid only dissolves it on boiling, while in hydrochloric 
acid it is readily soluble, even in the cold. It dissolves 
in potassium and sodium hydrate, with evolution of H, to form 
aluminates:— 
Owing to its stability in air and beautiful lustre, aluminium 
is sometimes employed for vessels and ornaments. The alloy 
of copper with 10-12 per cent, aluminium is distinguished by 
its great hardness and durability. It may be poured into 
moulds, and possesses a gold-like color and lustre. Under the 
name of aluminium bronze, it is used for the composition of 
various articles, like pocket wTatches, spoons, etc. It forms 
only compounds of the form A1X3 or A12X6 (p. 340). Its 
salts soluble in water have an acid reaction, and sweet, 
astringent taste. 
A1 + 3K0H = K3AIO3 + 3H 
The heat of formation of some of the aluminium compounds 
equals :— 
(AlatCle) =■ 321.8. Al2,Br6 = 239.3. A12,I6 = 140.6. 
(A12,C16,Aq) = 475.5. (Al2,Br6.Aq) =409.9. (Al2,I6,Aq)= 818.6. 
(A12,03, 3H20) = 388.8. 
The heat evolved in the formation of a quantity of aluminium hy- 
drate, corresponding to one atom of oxygen, is 129.6 ; as that of water 
is far less (H2,0 = 69.0), it must be decomposed by aluminium, with 
liberation of hydrogen (p. 267). If this does not transpire under ordi- 
nary conditions, the reason is to be sought for in the insolubility of 
aluminium hydrate. Indeed, the reaction occurs if there be added to 
water aluminium chloride, or another salt, in which the aluminium 
oxide is soluble. Conversely, from the high heat of formation is ex-- 
plained why aluminium oxide is not reduced by carbon. 
Aluminium Chloride, A1C13, or A12C16, is produced by the 
action of chlorine upon warmed or heated aluminium ; also 
by heating a mixture of aluminium oxide and carbon in a 
current of chlorine:— 
Chlorine and carbon alone do not act upon the oxide; by 
their mutual action, however, the reaction occurs in conse- 
quence of the affinity of carbon to oxygen, and of chlorine for 
aluminium. Similar deportment is shown by the oxides of 
boron and silicon. 
A1203 + 3C + 6C1 = A12C16 + 3C0. 344 
INORGANIC CHEMISTRY. 
By sublimation aluminium chloride may be got in white, 
liexagoual leaflets. It readily sublimes, but under elevated 
pressure fuses. Its vapor density is 133.9 (H = 1), from 
which the molecular formula, A12C16 = 367.8, is derived. 
From the specific heat of free aluminium it follows that its 
atomic weight is 27.3. 
Aluminium chloride absorbs moisture from the air and 
deliquesces. From concentrated hydrochloric acid solution it 
crystallizes with 6 molecules of water. On evaporating the 
aqueous solution, the chloride decomposes into aluminium 
oxide and hydrogen chloride:— 
It forms double chlorides with many metallic chlorides, viz.: 
AlCl3.NaCl, AICI3.KCI. The solutions of these may be evap- 
orated to dryness without decomposition. Further, it unites 
with many halogen derivatives of the metalloids:— 
A12C16 + 3H20 = A1203 + 6HC1. 
A1C13. PCI,, AICI3. POCl3, A1C18. SCI*. 
Aluminium Bromide—Al2Br6—is obtained like the chloride, 
and is composed of shining leaflets which fuse at 90° and boil at 265- 
270°. Its vapor density is 267.4 (H = 1), corresponding to the for- 
mula Al2Br6. It behaves like the chloride. 
Aluminium Iodide—A12I6—is formed by the direct union of 
aluminium filings with iodine, with gentle application of heat. It is a 
white, crystalline mass, fusing at 185°, and boiling about 400°. 
Aluminium Fluoride—A1F13 or A12F16—obtained by conduct- 
ing hydrogen fluoride over heated aluminium oxide or hydrate, sub- 
limes at a red heat, in colorless rhombohedra. It is insoluble in water, 
unaltered by acids, and is very stable. With alkaline fluorides it 
yields insoluble double fluorides. The compound—A1F13 3NaFl—oc- 
curs in Greenland, in large deposits, as Cryolite, and is employed in 
the soda manufacture (p. 283). 
Aluminium Oxide—A1203—is found crystallized in hexag- 
onal prisms in nature, colored by other admixtures, as ruby, 
sapphire and corundum. Impure corundum, containing alumi- 
nium 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. Artificially, 
aluminium oxide may be obtained by igniting the hydrate, as 
a white amorphous powder, which, in the oxy-hydrogen flame, 
fuses to a transparent glass. A mixture of aluminium fluo- 
ride and boron trioxide heated to Avhite heat has the boron ALUMINIUM 
fluoride volatilized, and crystallized aluminium oxide re- 
mains :— 
A12F16 + B203 = A1203 + 2BF13 
The crystallized or strongly ignited aluminium oxide is al- 
most insoluble in acids; to decompose it, it is fused with 
caustic alkalies or with primary potassium sulphate—HKS04. 
Aluminium Hydrates—The normal hydrate, Al(OH)3 or 
Al2(OH)6, occurs in nature as Hydrargillite. The hydrate 
A1202(0H)2 is diaspore. Bauxite is a mixture of the hydrate 
A120(0H)4, with ferric oxide. The normal hydrate is artifi- 
cially obtained as a white voluminous precipitate, by adding 
ammonium hydrate or carbonate (in latter case carbon diox- 
ide escapes, p. 347) to a soluble aluminium salt. Freshly 
precipitated, it dissolves in acids and in potassium and sodium 
hydrates. By long standing under water, or after drying, it is, 
without any alteration in composition, difficultly soluble in 
acids. Carefully heated, the normal hydrate first passes into 
AlO.OH. 
The freshly precipitated hydrate dissolves readily in a solution of 
aluminium chloride or acetate. On dialysing this solution the alumi- 
nium salt or crystalloid diffuses, and in the dialyser remains the pure 
aqueous solution of the hydrate. This has a faint alkaline reac- 
tion and is coagulated by slight quantities of acid, alkalies and many 
salts; the soluble hydrate is converted into the insoluble gelatinous 
modification. 
Gelatinous aluminium hydrate possesses the property of precipitating 
many dyestuffs from their solutions, forming with them colored insol- 
uble compounds (lakes). On this is based the application of alumi- 
nium hydrate as a mordant in dyeing. Generally the acetate is used 
for this purpose. The fibre of goods saturated with this salt is 
heated with steam, which causes the decomposition of the weak 
acetate ; acetic acid escapes, while the separated aluminium hydrate 
sets itself upon the fibre of the material. If the latter now be intro- 
duced into the solution of coloring matter it is fixed by the aluminium 
hydrate upon the fibre. At present, sodium aluminate is employed 
instead of the acetate. 
Aluminium hydrate has a feeble acid character, and can 
form salt-like compounds with strong bases. On carefully 
evaporating its solution in sodium and potassium hydrate, or 
upon addition of alcohol, white amorphous compounds of 
KA10-2, NaAlOj and (NaO)3Al are obtained. The potassium 
compound can be obtained in crystalline form. These deriv- 
atives, known as aluminates, are not very stable, and are even 346 
INORGANIC CHEMISTRY. 
decomposed by carbon dioxide, with elimination of aluminium 
hydrate:— 
The thus produced aluminium hydrate, in distinction to 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 com- 
merce. 
2A102Na + C02 = A1203 + C03Na2. 
On adding calcium hydrate to the solution of potassium or 
sodium aluminate, insoluble calcium aluminate is precipi- 
tated :— 
Barium and strontium aluminates are, however, readily solu- 
ble in water. Similar aluminates frequently occur as crys- 
tallized minerals, in nature. Thus the spinels consist chiefly 
of magnesium aluminate, A10.0/Mgi chrysoberyl is beryl- 
lium aluminate, AlO 0/^e ’ gahnite is zinc aluminate, 
A10.0\v 
A10.0/^n' 
2A102Na + Ca(OH)2 = (A102)2Ca + 2NaOH, 
Technically, alumina is obtained from cryolite, bauxite and 
other alumina-containing minerals. The pulverized bauxite 
is heated with dry sodium carbonate, in furnaces, and the 
formed sodium aluminate extracted with water. From the 
clear solution carbon dioxide precipitates the hydrate, while 
sodium carbonate remains dissolved, and is afterward recov- 
ered. The dried aluminium hydrate comes as a white powder 
in trade. 
The gelatinous, readily soluble (colloidal) aluminium hy- 
drate (see above) precipitated from acid solutions by alkalies, 
has lately been prepared upon a large scale, according to the 
method of Lowig, by treating the sodium aluminate solution 
with milk of lime; calcium aluminate precipitates, while in 
solution remains sodium hydrate:— 
2A102Na + Ca(OH)2 = (A102)2Ca + 2NaOH. 
Calcium aluminate is dissolved in hydrochloric acid:— 
and to the solution the necessary quantity of calcium alumi- 
nate added, which precipitates aluminium oxide:— 
(A102)2Ca + 8HC1 = 2A1C13 + CaCl2 + 4H20. 
2A1C13 + 3(A102)2Ca = 4A1203 + 3CaCl2. ALUMINIUM. 
347 
According to this procedure, the sodium hydrate formed in 
the first reaction is obtained together with the alumina. 
The basic character of aluminium hydrate exceeds the acid ; 
it is, however, so feeble, that it is not capable of forming salts 
with weak acids, like carbon dioxide, sulphurous acid and 
hydrogen sulphide. When sodium carbonate is added to 
aluminium salt solutions, aluminium hydrate is precipitated, 
while C02 is set free:— 
A12C16 + 3Na2C03 + 3H20 = Al2(OH)6 + 6NaCl + 3C02. 
The alkaline sulphides behave similarly:— 
Aluminium Sulphate—A12(S04)3—crystallizes from aqueous 
solution with 18 molecules of H20 in thin leaflets with pearly 
lustre. These dissolve readily in water; when heated, they 
melt and lose all their water of crystallization. The sulphate 
is obtained by dissolving the hydrate in sulphuric acid, or by 
the decomposition of pure clay with the same acid ; the residual 
silicic acid is removed by filtration, and the solution of the 
sulphate evaporated. When a quantity of ammonium, in- 
sufficient 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:— 
A12C16 + 3(NH4)2S + 6H20 = Al2(OH)6 + 6NH4C1 + 3H2S. 
*.<£!>*+7h2°- 
Aluminium sulphate can combine with the alkaline sul- 
phates to form double salts, termed alums :— 
(S04)3A12.S04K2 + 24H20 or (S04)2A1K + 12II20. 
Their constitution is expressed by the following formula :— 
04 /S04K | 24JT Q 
04S^A12\S04K t-J4±12U 
or 
O.S V” 
S04Q' + 12H20. 
In this compound the potassium may be replaced by sodium, 
ammonium, rubidium, csesium, and also by thallium. Iron, 
chromium and manganese form like derivatives :— 
Fe2(S04)3.K2S04 + 24H20 
Potassium Iron Alum. 
Mn2(S04)s.Na2S04 + 24H20. 
Sodium Manganese Alum. 
All these alums crystallize in regular octahedra or cubes, and 
can form isomorphous mixtures. 348 
INORGANIC CHEMISTRY. 
The most important of them is Potassium Aluminium Sul- 
phate, or ordinary alum, A1K(S04)2 + 12H20. From water 
it crystallizes in large, transparent oetahedra, soluble in 8 
parts water of ordinary temperature, or in i part boiling 
water. The solution has an acid reaction and a sweetish, 
astringent taste. When placed over sulphuric acid, alum 
loses 9 (or 18) molecules of H20. It melts, when heated in 
its water of crystallization, loses all the latter and becomes a 
white, voluminous mass—burned alum. Upon adding to a 
hot alum solution a little sodium or potassium carbonate, the 
hydrate first produced dissolves, and when the liquid cools, 
the alum crystallizes out in cubes, as cubical alum. The addi- 
tion of more sodium carbonate causes the precipitation of the 
basic salt—A1K(S04)2.A1(0H)3. Alunite, found in large 
quantities near Rome and in Hungary, has a similar com- 
position. 
Commercially, alum is obtained according to various methods: 1. 
From alunite, in which it exists already, by heating and extracting with 
water. In this way alum dissolves, while the hydrate 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 alum shale, a clay containing pyrite and peat. This is 
roasted and moistened with water exposed for a long time to the action 
of the air. By this means FeS2 is converted into 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 directly treated 
with sulphuric acid, and to the solution of aluminium sulphate potassium 
or ammonium sulphate are added. Bauxite and cryolite are admirable 
material for the preparation of alum. The working of cryolite for 
alumina and soda is described on p. 283, and that of bauxite p. 345. 
Ammonium Alum—(S04)2 A1 NH4 -f 12H20—crystallizes 
like potassium alum, in large crystals, and at present, owing to 
its cheapness, is almost exclusively applied technically. 
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 applicable. Further- 
more, at present, in practice we find alum more and more 
supplanted by aluminium sulphate and sodium aluminate, 
Avhich can be procured perfectly free from iron. 
Aluminium Phosphate—A1P04—is thrown out of aluminium 
solutions by sodium phosphate, as a white gelatinous precipi- 
tate ; this is readily soluble in acids, excepting acetic. ALUMINIUM. 
349 
Aluminium Silicates. The most important of the aluminium 
double silicates, so widely distributed in nature, is ordinary 
feldspar—Orthoclase—AlKSi308—and the various micas, 
which, with quartz, compose granite. By disintegration of 
these, under the influence of water and C02 of the air, 
alkaline silicates are dissolved and carried away by water, 
while the insoluble aluminium silicate, clay, remains. Per- 
fectly pure clay is white, and is called kaolin, or porcelain 
clay; its composition mostly corresponds to the formula, 
Al2(Si02)3, A1205H4. 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 earthen ware, from the red brick to porcelain. 
To produce porcelain a very fine mixture of kaolin, feldspar and 
quartz is employed. On strong ignition, feldspar fuses, fills the pores 
of the clay and thus furnishes a fused transparent mass—porcelain. 
When it is not so strongly glowed, it remains porous—faience— 
serving for finer clay vessels. To render these impervious to water, 
they are covered with glazing. This consists of various readily fusible 
silicates. Rough earthenware vessels are constructed from impure 
clay, and they are ordinarily glazed by throwing salt, at the time of 
burning, into the ovens. The hot steam decomposes this into hydro- 
chloric acid and sodium hydrate, which forms with the clay, on its sur- 
face, an easily fusible silicate. 
Ultramarine. The rare mineral Lapis lazuli, which 
formerly, under the name of Ultramarine, was employed as a 
very valuable blue color, is a compound of aluminium sodium 
silicate with sodium polysulphides. At present ultramarine 
is prepared artificially, in large quantities, by heating a mixture 
of clay, dry soda (or sodium sulphate), sulphur and wood 
ashes, away from air. Green ultramarine is the product. This 
is then washed with water, dried, mixed with powdered sul- 
phur and gently heated with air contact until the desired blue 
color has appeared—blue ultramarine. The cause of the blue 
coloration is generally assumed to be due to the existence of 
a complicated sulphur compound, whose nature is not yet 
explained. On pouring hydrochloric acid over the blue pro- 
duct, the color disappears with liberation of sulphur and hydro- 
gen sulphide—which wTould point to the existence of a poly- 
sulphide. At present violet and red ultramarines are pre- 
pared. 350 
INORGANIC CHEMISTRY. 
In some very rare minerals, like cerite, gadolinite, euxenite, orthite, 
occurring principally in Sweden and Greenland, is found a series of 
metals which in their entire deportment closely resemble aluminium 
(p. 341). These are yttrium, cerium, lanthanum, didymium, and the 
mote recent scandium, ytterbium, erbium, terbium, thulium (philli- 
pium, decipium). These mostly form difficultly soluble oxalates; 
they are precipitated from solution by oxalic acid. They also afford 
difficultly soluble sulphates and double sulphates, of which all the 
potassium double salts are constituted, according to the formula 
Me2(S04)3, 3K2S04. The different decomposability of their 
nitrates upon application of heat affords an excellent means for their 
isolation and separation. 
Most accurately investigated are lanthanum, cerium and didymium, 
whose atomic weights, by the determination of the specific heat of the 
free metals are at least very approximately established. That of 
yttrium (89) appears, from the isomorphism of its sulphate with that of 
didymium to be positively established. Erbium, formerly regarded as 
an elementary substance, consists, according to Nilson and Cleve, of 
six different earths: scandium, ytterbium, thulium, erbium, terbium, 
and a metal designated X. Only the first two are as yet accurately 
characterized, and especially does scandia offer particular interest, as 
the metal scandium contained in it, with the atomic weight 44, fills out 
an until now vacant place in tbe first large period. All its properties 
have essentially shown themselves the same as those of the element 
ekaboron, theoretically deduced from the periodic system. 
RARE METALS. 
Scandium—Sc=44—chiefly contained in euxenite and gadolinite, 
has not yet been obtained in a free condition. Its oxide, Sc203. is 
obtained by igniting the hydrate or nitrate, and is a white, infusible 
powder (like magnesia and oxide of beryllium). Its specific gravity 
equals 3.86 the specific heat 0,1530. The hydrate Sc (OH)s is pre- 
cipitated as a gelatinous mass from its sails by the alkalies, and insolu- 
ble in an excess of the latter. The nitrate crystallizes in little prisms, 
and is decomposed with difficulty upon heating. The potassium double 
sulphate Sc2 (So4)3, 3K2 So4, is soluble in warm water, but not in a 
solution of potassium sulphate. The chloride affords a characteristic 
spark spectrum. 
Yttrium—Y—-89, has long been known in its compounds, but has 
never been investigated in pure condition. Occurs chiefly in gadoli- 
nite (upwards of 35 per cent.) Its potassium double sulphate is sol- 
uble in a potassium sulphate solution, whereby it can readily be 
separated from cerium, lanthanum, and didymium. Its nitrate is 
much more difficult to decompose than that of scandium and ytterbi- 
um. The chloride YC13+7H20 forms large prisms and gives a 
spark spectrum. 
Lanthanum—La=139— separated from its chloride by electroly- 
sis, resembles ironas regards color .and lustre, oxidizes in the air, and 
burns in a flame with a bright light. Its specific gravity equals 6.16, GALLIUM GROUP. 
351 
the specific heat 0.0448. The hydrate La(OH)3 is precipitated as a 
gelatinous mass, and reacts alkaline. 
Cerium—Ce = 140—occurs in cerite (60 per cent.), and is also 
obtained by the electrolysis of the chloride. Very similar to lanthanum, 
but at ordinary temperatures is more stable than the latter; 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, the specific 
heat 0.0448. Besides the salts of the sesquioxide Ce203 it forms some 
from the dioxide CeOs. The first are colorless, while the latter are 
colored yellow or brown ; from the first, on addition of hypochlorites, 
red ceric hydrate Ce(OH)4 is precipitated. 
Didymium—Di = 142 or 146.5—in free condition very closely 
resembles lanthanum, but shows a somewhat yellowish color; it oxi- 
dizes in the air and burns in the flame with a brilliant light. Its 
specific gravity is 6.54, the specific heat 0.0456. Its hydrate Di(OH)3 
and salts are rose red or violet in color. 
. All three metals, didymium, cerium and lanthanum, as chlorides, 
yield spark spectra. Didymium is distinguished by the ability of its 
salts, in solid or dissolved form, of absorbing definite light rays and 
yielding a very characteristic absorption spectrum. 
Ytterbium—Yb = 173.—Its oxide, Yb >03, is obtained from the 
so-called erbium earth (from euxenite and" gadolinite) by repeated 
partial heating of the mixed nitrates, whereby the scandium nitrate first 
decomposes. It is a white infusible powder, of specific gravity 9.17 ; 
its specific heat is 0.0616. The salts of ytterbium are colorless, and 
show no absorption spectrum. 
Erbium Oxide—Er303 (Er = 166)—is red in color. Its salts, 
like those of didymium yield an absorption spectrum. 
The same relation that Cu, Ag and Au bear to Na, and Zn, 
Cd and Hg to Magnesium, do the three heavy metals, gallium, 
indium, and thallium, bear to aluminium: 
GALLIUM GROUP, 
Cu 63.3 
Ag 107.6 
Au 197.0 
Zn 65.0 
Cd 111.6 
Hg 200.0 
Ga 69.8 
In 113.4 
T1 203.6 
Sn 118.0 
Pb 206.0 
As 75.0 
Sb 122.0 
Bi 210.0 
They constitute the corresponding members of the three 
great periods, and as second sub-group attach themselves to 
aluminium, while the first more basic group is formed by some 
cerite metals (p. 341). The entire character of the three ele- 
ments under consideration is influenced by their position in 
the periodic system, as in all respects, regular relations 
appear, e. g., seen in the specific gravities, fusing points and 
other physical properties of the free metals. 352 
INORGANIC CHEMISTRY. 
Ga 
In 
T1 
Atomic Weight, 
Specific Gravity, 
F using Point, 
69.8 
5.9 
29.5° 
113.4 
7.4 
176° 
203.6 
11 8 
290° 
Belonging to Group III of the periodic system, Ga, In 
and T1 yield compounds of the triatomic form, which in many 
respects are analogous to those of aluminium. 
Like other elements with high atomic weight (Au, Hg, Pb), 
thallium exhibits great variations from the group properties 
(p. 315) which are seen in the fact that in addition to the 
compounds of the form T1X3 the metal is also capable of afford- 
ing those of the form T1X. Regarding thallium as a member 
of the last great period (Pb, Au, Hg, Tl, Pb, Bi) we discover 
that its combination forms, like those of other elements of this 
series, are constructed upon a remarkable regularity, and as 
well according, indeed, with reference to the highest as to 
the lowest forms :— 
PtCla 
PtCl4 
AuCl 
AuCls 
HgCl 
HgCl2 
T1C1 
T1C1, 
PbCl2 
PbCl4 
BiCl, 
BiX5. 
1. GALLIUM 
Ga - 69.8. 
Gallium was discovered in 1875, by Lecoq de Boisbaudran, by 
means of the spectroscope, in zinc blende from Pierrefitte. Already, 
in the year 1870, Mendelejeff, upon the basis of the table of the 
periodic system devised by him, predicted the existence of an element 
(standing between aluminium and indium, with an atomic weight of 
nearly 64), which he named Eka-aluminium, the properties of which 
were, to an important degree, determined by its position in the periodic 
system. The agreement of the most at that time known properties of 
gallium with those of the theoretical eka-aluminium, made it exceed- 
ingly probable that gallium represented this theoretically established 
element. At present the complete confirmation is afforded by the fact 
that the experimentally determined atomic weight has shown itself like 
the theoretical. 
As yet gallium has only been found in very small quantity, and is 
but imperfectly investigated. It is characterized by a spectrum con- 
sisting of two violet lines. Separated by the electrolysis from an 
ammoniacal solution of its sulphate, it is a white, hard metal, of specific 
gravity 5.9, with a fusing point 29 5°. It is only superficially oxidized 
in the air, not altered by water, and up to a red heat is not volatile. 
Like aluminium, it is scarcely attacked by nitric acid, but dissolves 
readily in hydrochloric acid. INDIUM. 
353 
The oxide Ga203 is thrown out of its solutions as a white precipitate, 
readily soluble in potassium hydrate, but little soluble in hydrate of 
ammonium. The chloride GaCl3 is very deliquescent, and is decom- 
posed by much water- The nitrate and sulphate are readily crystallized; 
the latter forms a double salt with ammonium sulphate—similar to 
the alums:— 
Hydrogen sulphide only precipitates gallium from acetate solutions. 
(S04)3 Ga2. S04(NH4)2 + 24H20. 
INDIUM. 
Owing to its resembling zinc, indium was regarded as a dyad metal, 
and its compounds composed according to the formula, InX2 ; 
from this its atomic weight was found to be 75.4. From the specific 
heat, however, it follows that the atomic weight is one and a half times 
as large, (p. 250). Therefore it is triatomic and its derivatives are 
constituted according to the form, InX3. It belongs to the group 
of aluminium, to which, in its derivatives, it manifests some similarity. 
It was discovered in 1863, by Reich and Richter, by the aid of 
spectrum analysis. Its spectrum is characterized by a very bright 
indigo blue line, hence its name. It only occurs in very minute quan- 
tities in some zinc blendes from Frieberg and the Hartz. 
It is a silver white, soft and tenacious metal, of specific gravity, 7.42. 
It fuses at 176° and distills at a white heat. At ordinary temperatures 
it is not altered in the air ; heated, it burns with a blue flame to indium 
oxide. In hydrochloric and sulphuric acids it is difficultly soluble, but 
readily dissolved by nitric acid. 
In — 113. 
Indium Chloride—InCl3—results from the action of chlorine on 
metallic indium, or upon a glowing mixture of indium oxide and 
carbon. It sublimes in white, shining leaflets, which deliquesce in the 
air. 
Indium Oxide—ln203—is a yellow powder resulting from the 
ignition of the hydrate. 
Indium Hydrate—In(OH)3—is precipitated as a gelatinous mass, 
by alkalies, from indium solutions. It is soluble in sodium and potas- 
sium hydrates. 
Indium Nitrate—In(N03)3—crystallizes with three molecules 
of water, in white deliquescent needles. 
Indium Sulphate—In2(S04)3—is a gelatinous mass, with three 
molecules of water, remaining after the evaporation of a solution of 
indium in sulphuric acid. With ammonium sulphate it forms an 
alum. 
Indium Sulphide—In2S3—is precipitated as a yellow colored 
compound, from indium solutions by hydrogen sulphide. 354 
INORGANIC CHEMISTRY, 
THALLIUM. 
Is rather widely distributed in nature, but in very small 
quantity. It is present in some pyrites. On roasting these, 
for the production of sulphuric acid, according to the chamber 
process, it deposits as soot in the chimney and in the chamber 
sludge, and was discovered in the latter, in 1861, almost simul- 
taneously, by Crookes and Lamy, by means of the spectroscope. 
To get the thallium the chimney dust is boiled with water 
or sulphuric acid, and from the solution thallous chloride is 
precipitated by hydrochloric acid. Then the chloride is 
converted into sulphate, and from the latter, by means of zinc 
or the galvanic current, the metal is separated. Thallium 
is a white metal as soft as sodium, of specific gravity 11.9. It 
fuses at 285° and distills at a white heat. In moist air it oxidizes 
very rapidly. It does not decompose water at ordinary 
temperatures. It is best preserved under water, in a closed 
vessel. By air access it gradually dissolves in water to 
hydrate and carbonate. Heated in the air it burns with a 
beautiful green flame whose spectrum shows a very intense 
green line, hence the name thallium, from tfal/Wc, green. 
Thallium is readily soluble in sulphuric and nitric acids, little 
attacked by hydrochloric acid, owing to the insolubility of 
thallous chloride. 
Tl = 204. 
i 
Thallium forms two series of compounds: Thallous—T1X 
III 
and thallic—T1X3. The first are very similar to the compounds 
of the alkalies (and also those of silver). This is especially 
seen in the solubility of the hydrate and carbonate in water; 
their solutions have an alkaline reaction. Further, many 
thallous salts are isomorphous with those of potassium, and 
form perfectly similar double salts. The insolubility of its 
sulphur and halogen compounds would make monatomic thal- 
lium approach silver and lead. 
In the compounds constituted according to the form T1X3, 
thallium, like aluminium, is triatomic; otherwise it scarcely 
shows any similarity to it. 
(TlaO) = 42.2 (T1,C1) = 48.5 (Tl,Br) = 41.2 (Tl.I) = 30. 
The heat of union of some of the thallous compounds is :— 
(T1,0,H) = 56.9 (T12,S,04) = 210.9 (T1,N,03) = 58.1. THALLIUM. 
355 
The heat of solution of all these compounds is negative. From the 
cited numbers is explained the deportment of thallium to water and the 
acids. 
The heat of formation of the ic compounds in aqueous solution 
equals:— 
(Tl,Cl,Aq) = 89.0 
(Tl,Br3Aq) = 56.1 
(Tl,I3,Aq) = 10.5 
(T1,203,3H20) = 86.9. 
THALLOUS COMPOUNDS 
Thallous Oxide—T120—is formed by the oxidation of 
thallium in the air, or by heating the hydrate to 100°. It 
forms a black powder which dissolves in water to yield the 
hydrate. 
Thallous Hydrate—Tl(OH)—may be prepared by decom- 
posing thallium sulphate with an equivalent amount of barium 
hydrate, and crystallizes with one molecule of water in 
yellowish prisms. It dissolves readily in water and alcohol 
to strong alkaline solutions. 
Thallous Chloride—T1C1—is thrown down from solutions 
of thallous salts by hydrochloric acid as a white, curdy 
precipitate, which is very difficultly soluble in water. Like 
potassium chloride it forms an insoluble double salt—PtCl4, 
2T1C1—with platinic chloride. In the same manner thallous 
bromide is obtained as a white and thallous iodide as a yellow 
precipitate. 
Thallous Sulphate—T12S04:—crystallizes in rhombic prisms, 
isomorphous with potassium sulphate. At ordinary tempera- 
tures it dissolves in 20 parts of water. With the sulphates of 
the metals of the magnesium group, of ferrous oxide, of 
cupric oxide, etc. (p. 307), it affords double salts, e. g., 
MgS04Tl2S04 -f- 6ILO, which are perfectly similar and analo- 
gous to the corresponding double salts of potassium and 
ammonium. Further, with the sulphates of the sesquioxides 
of the iron group it forms Thallium Alums, e. g., A1T1(S04)2 
-j- 12H20, similar to potassium alum—A1K(S04)2 -f- 12H20. 
Thallous Carbonate—T12C03—is formed from the oxide by 
absorption of C02, crystallizes in needles, which dissolve at 
ordinary temperatures, in 20 parts of water, and has an alkaline 
reaction. 
Thallous Sulphide.—T12S—is precipitated from thallous 
salts by hydrogen sulphide, as a black compound, insoluble in 
water. 356 
INORGANIC CHEMISTRY. 
Thallic Chloride—T1C13—is formed by the action of chlor- 
ine upon T1C1, and is very soluble in water. At 100° it de- 
composes into T1C1 and Cl2. From its solutions the alkalies 
precipitate thallous hydrate, TIO.OII, a brown powder, which, 
at 100°, passes into black thallic oxide T1203. Further heat- 
ing decomposes the latter into thallous oxide and oxygen. 
The oxide and hydrate are soluble in hydrochloric, nitric, 
and sulphuric acids, forming T1(N03)3,T]2(S04)3,T1C13. 
On conducting chlorine through a solution of thallic hydrate 
in potassium hydrate, the same assumes an intense violet 
color, due probably to the formation of the potassium salt of 
thallic acid, the composition of which is yet unknown. 
The thallium compounds are poisonous. They are employed 
in making thallium glass, which refracts light more strongly 
than lead glass. The spectrum of the thallium flame shows a 
very bright green line. 
THALLIC COMPOUNDS. 
Tin and lead, with silicon and carbon, belong to one group 
in which the transition, corresponding with the increasing 
atomic weight, from metalloidal to metallic character, finds 
distinct expression: C = 12, Si = 28, —= 73, Sn = 117.8, 
Pb = 206.4. The differences between these elements are 
purely gradual; the apparent greater difference between met- 
alloidal silicon and metallic tin is explained by the fact that 
a member corresponding to arsenic of the nitrogen group, with 
atomic weight 73, is not known (p. 156). 
Tin, like carbon and silicon, is tetratomic, and yields per- 
fectly analogous compounds, SnCl4,Sn02. Corresponding to 
the latter are the hydrates Sn(OH)4, and SnO(OH)2, of which 
only the latter forms salts (p. 218). Yet stannic acid is but a 
weak acid, the alkali salts of which react alkaline and are not 
very stable. At the same time these hydrates have a basic 
character, hence form salts with acids. Therefore stannic acid 
anhydride is called stannic oxide, and stannic acid stannic 
hydrated oxide. The metallic salts of stannic acid (like 
Sn03,Xa2), are called stannates; those, with acids, are stannic 
salts. 
TIN AND LEAD. 
Thus, too, tin affords compounds of the form SnX2 (cor- 
responding to CO), in which it figures as a dyad. The mon- 357 
oxide, also called stannous oxide, possesses a decided basic 
character, and only yields salts with acids—stannous or 
sfanno-salts. 
A further advance in metallic basic character is exhibited 
by lead, which, like tin, forms two series of compounds of the 
forms PbXt and PbX2. While, however, with tin the com- 
pounds SnX4 are more stable than those of the form SnX2 
lead affords almost exclusively compounds of the form PbX2 
in which it figures as a dyad. 
The tetravalence of lead shows itself almost entire in its 
metallo-organic compounds: as (Pb(CH3)4,Pb(C2H5)4—p. 239); 
further in lead dioxide Pb02. The latter does not form corres- 
ponding salts with the acids, but yields up one atom of 
oxygen with formation of salts of the monoxide PbO. Warmed 
with hydrochloric acid, it liberates chlorine; hence behaves like 
the peroxides, and is commonly called lead peroxide (p. 256). 
A distinct gradation shows itself in the series Si02,Sn02, 
Pb02. The stability and acidity grow less successively, yet 
lead dioxide preserves the acidic character and forms salts with 
the alkalies (as Pb03K2) which are very similar to the salts of 
stannic acid; therefore lead dioxide Pb02 is to be regarded as 
the oxide of a plumbic acid H2Pb03. 
The connection of tin and lead in one group results from 
the heat of formation of their compounds:— 
(Sn.Cl2) = 80.8 
Pb.Ci2 = 82.7 
(Sn.O) = 68.0 
(Pb.O) =60.3 
(Sn.Cl4) = 127.0 
(Pb.Cl4) = 
(Sn.02) = 136.0 
(Pb.02) = 74.0 
From these numbers is seen that from tin chloride (as from 
other acid salts), the metal is separated by lead, while con- 
versely, from lead oxide in alkaline solution, lead is pre- 
cipitated by tin (p. 259). 
1. TIN 
In nature tin occurs principally as dinoxide (Cassiterite-tin 
stone) in England (Cornwall), Saxony and India. To pre- 
pare the metal the oxide is roasted, lixiviated and heated 
in a furnace with charcoal:— 
Sn = 118 
Thus obtained, it usually contains iron, arsenic and other 
metals; to purify it the metal is fused at a low temperature, 
Sn02 + 2C = Sn +2C0. 358 
INORGANIC CHEMISTRY. 
when the pure tin flows away, leaving the other metals. The 
tin obtained in the Indian isles (Malacca) is almost chemi- 
cally pure, while that of England contains traces of arsenic 
and copper. 
Tin is an almost silver white, strongly lustrous metal, with 
a specific gravity of 7.3. It possesses a crystalline structure ; 
on bending a rod of it, a peculiar noise (tin cry) is heard, 
due to the friction of the crystals. Upon etching a smooth tin 
surface with hydrochloric acid its crystalline structure is recog- 
nized by the appearance of remarkable striations. Perfectly 
pure compact tin, at low temperature, passes into an aggregate 
of small quadratic crystals. The metal is tolerably soft and 
very ductile, and may be rolled out into thin leaves. At 200° 
it becomes brittle, and may then be fractured. It fuses at 
228°, and distills at a white heat; heated in the air, it burns 
with an intense white light, forming tin dioxide. It does not 
oxidize in the air at low temperatures, and withstands the 
action of many bodies, hence employed in tinning copper and 
iron vessels for household use. 
Of the alloys of tin, besides bronze and soft solder, brittania 
metal is remarkable. It contains 9 parts tin and 1 part 
antimony, and frequently, also, 2-3% tin and 1% copper. 
In hot hydrochloric acid, tin dissolves to stannous chloride, 
with evolution of hydrogen gas :— 
By somewhat dilute nitric acid, it is oxidized to metastan- 
nic acid ; anhydrous nitric acid HN03 does not change it. 
When boiled with potassium or sodium hydrates, it dissolves, 
forming stannates:— 
Sn + 2HCl = Sn C12 + 2H. 
Tin forms two series of compounds ; the stannous and stannic, 
or compounds of stannic acid (p. 356). 
Sn + 2K0H + H20 = Sn03K2 + 2H2. 
1. STANNOUS COMPOUNDS. 
Tin Dichloride—Stannous chloride, SnCl2—results when tin 
dissolves in concentrated hydrochloric acid. Upon evaporation 
it crystallizes with two molecules of water (SnCl2 -j- 2H20) 
which it loses at 100° C. It is used in dyeing, as a mordant, 
under the name of Tin Salt. The anhydrous chloride, 
obtained by heating the metal in dry hydrochloric acid gas, 359 
fuses at 250° and distills without decomposition at a red heat. 
Its vapor density at 600-700° agrees with the formula Sn2CU; 
at 900° with SnCL. 
Water in part decomposes stannous chloride. Its solution 
is strongly reducing, and from air it absorbs oxygen with the 
separation of basic tin chloride:— 
3SnCl2 + 0 + H20 = + SnCl2. 
In presence of hydrochloric acid only stannic chloride is 
produced. From mercuric chloride solution stannous chloride 
precipitates mercurous chloride and metallic mercury (p. 
320). It unites with chlorine to form stannic chloride, and 
with many chlorides to yield double salts, e. g., 
SnCl2,2KCl and SnCla,2NH4Cl. 
Tin Mon oxide, SnO, or Staymous Oxide, is obtained by heating 
its hydrates, Sn02H2, in an atmosphere of carbon dioxide, and 
is a blackish-brown powder which burns when heated in the 
air, becoming stannic oxide. Sodium carbonate added to a 
solution of stannous chloride precipitates white 
Stannous Hydrate—stanno-hydrate—Sn(OH)2: 
SnCl2 + C03Na2 + H20 = Sn(OH)2 + 2NaCl + C02 
It is insoluble in ammonium hydrate, but is readily dissolved 
by potassium hydrate. Upon slow evaporation of the alkaline 
solution dark crystals of SnO separate ; on boiling the solution 
the hydrate decomposes into potassium stannate, K2Sn03, 
which remains dissolved, and metallic tin. By solution in 
acids the hydrate forms salts. 
Tin Dichloride—Stannous chloride—SnCl2—and stanno- 
sulphate—SnS04—are formed when tin is warmed with 
concentrated hydrochloric or sulphuric acids. The sul- 
phate separates by the evaporation of the solution, in small, 
granular crystals. 
Tin Monosulphide—Stannous sulphide—SnS—is precipi- 
tated from stannous solutions by hydrogen sulphide, as a 
dark brown amorphous precipitate. Obtained by fusing 
together tin and sulphur; it is a lead-gray crystalline mass. 
Dissolves in concentrated hydrochloric acid, with liberation 
of H2S to form stannous chloride; if boiled at same time with 
sulphur it dissolves as a sulpho-stannate:— 
SnS -f- S -f- K2S — K2SnS3. 360 
INORGANIC CHEMISTRY 
Tin Tetrachloride—Stannic Chloride—SnCl4—is produced 
by the action of chlorine upon heated tin or stannous chloride 
—SnCl2. It is a colorless liquid (Spiritus fumans Libavii), 
fuming strongly in moist air, of specific gravity 2.27, and 
boiling at 1503; its vapor density equals 130 (H = 1), cor- 
responding to the molecular formula, SnCl4 = 260. It 
attracts moisture from the air and is converted into a crystal- 
line mass (Butter of Tin), SnCl4 + 3H20, readily soluble in 
water. Boiling decomposes the solution into metastannic acid 
(H2Sn03) and hydrochloric acid :— 
2. STANNIC COMPOUNDS. 
SnCl4 + 3H20 = H2Sn03 + 4HC1. 
Tin chloride possesses a salt-like nature, combining with 
metallic chlorides to the so-called double salts, e. g., SnCl4. 
2KC1 and SnCl4.2NH4Cl; the latter compound is known as 
-pink salt in calico printing. With chlorides of the metalloids 
it also yields crystalline double salts, e. g., SnCl4.PCl5 and 
SnCl4.2SCl4. 
Tin Bromide—SnBr4—forms a white, crystalline mass. 
Tin Iodide—Snl4—orange red octahedra, fusing at 146° and boil- 
ing at 295°. It may be obtained by heating tin with iodine to 50°. 
Tin Fluoride—SnFl4—only known in combination with metallic 
fluorides (e. g., K2SnFl6), very similar and generally isomorphous 
with the salts of hydrofluosilicic acid (SiFl6K2). 
Tin Dioxide—Stannic oxide—Sn02—is found in nature, as 
tin stone, in quadratic crystals or thick brown masses, of 
specific gravity 6.8. It is prepared, artificially, by heating 
tin in the air, and forms then a white, amorphous powder. 
It may be obtained, crystallized, by conducting vapors of 
the tetrachloride and water through a glowing tube. The 
dioxide is infusible, and not soluble in acids or alkalies. 
Fused with sodium and potassium hydrate it yields stannates 
soluble'in water. 
On adding ammonium hydrate to the aqueous solution of 
tin tetrachloride or hydrochloric acid to the solution of potas- 
sium stannate (Sn03K2), a white precipitate of stannic acid 
will separate. This dissolves readily in concentrated nitric 
acid, hydrochloric acid and the alkalies. Preserved under 
water, or in vacuo, it becomes insoluble in acids and sodium 
hydrate. Both modifications appear to have the same compo- 
sition, H2Sn03, and the cause of the isomerism is not yet LEAD 
361 
explained. The insoluble modification is commonly called 
metastannic acid. It is also obtained as a white powder by 
warming tin with dilute nitric acid. On adding sodium 
hydrate to the insoluble stannic acid it is converted into 
sodium metastannate, insoluble in sodium hydrate, but readily 
dissolved by pure water. The salts of stannic oxide with 
acids, e. g., the sulphate, are not very stable, and washing with 
warm w ater decomposes them. More stable are the metal salts 
of stannic acid. The most important of these is sodium stan- 
nate—Na^SnO;)—which is employed in calico printing, under 
the name of preparing salts. Prepared upon a large scale 
by fusing tin stone with sodium hydrate. On evaporating the 
solution, it crystallizes in large, transparent, hexagonal 
crystals, containing three molecules of wTater. 
Tin Disulphide—Stannic Sulphide—SnS4—precipitated as 
an amorphous, yellow' powder by H2S from stannic solutions. 
If a mixture of tin filings, sulphur and ammonium chloride 
be heated, it is obtained in form of a brilliant crystalline 
mass, consisting of gold yellow scales. It is then called 
mosaic gold, and is applied in bronzing. The precipitated tin 
disulphide is dissolved by concentrated hydrochloric acid, con- 
verting it into stannic chloride; nitric acid converts it into 
metastannic acid. The sulphides and hydrosulphides of the 
alkalies dissolve tin disulphide as sulphostannates (see p. 214). 
Sodium sulphostannate SnS3Na2 -f- 2H20 crystallizes in color- 
less octahedrse. Acids decompose the sulphostannates, with 
the separation of tin disulphide. 
LEAD. 
Pb — 206.4. 
Lead (Plumbum) is found in nature principally as Galenite 
—PbS. Of the other more widely distributed lead ores may 
be mentioned, Cerussite—PbC03—Crocoisite (PbCr04) and 
Wulfenite (MoOJPb). For the preparation of lead, galenite 
is almost exclusively employed ; the process of its separation 
is very simple. The 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 :— 
PbS + 30 = PbO + S02 
and 
PbS + 0* = PbS04. 362 
INORGANIC CHEMISTRY 
Upon ignition these two substances react with the lead sul- 
phide, according to the following equations :— 
2PbO + PbS = 3Pb + S02 
and 
S04Pb + PbS =2Pb + 2S02. 
Metallic lead has a bluish white color, is very soft and 
tolerably ductile. Upon a freshly cut face it shows a bright 
lustre, but on exposure to air becomes dull by oxida- 
tion. Its specific gravity is 11.37. It fuses at 325° and 
volatilizes somewhat at a white heat. Heated in the air it 
burns to lead oxide. 
In contact with air and water, lead oxidizes to lead 
hydrate, Pb(OH)2, which is somewhat soluble in water. If 
the water contain carbonic acid and mineral salts, if only in 
slight quantity, as in mineral waters, no lead goes into solu- 
tion, 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 the water.) This 
behavior is very important for practical purposes, as lead 
pipes are frequently employed in conducting water for various 
objects. 
Sulphuric and hydrochloric acid have little effect on the 
metal, owing to the insolubility of its sulphate and chloride ; it 
dissolves readily in nitric acid to nitrate. Zinc, tin and iron 
precipitate it, as metal, from solution ; a strip of zinc immersed 
in a dilute lead acetate solution is covered with an arborescent 
mass, consisting of shining crystalline leaflets (lead tree). 
Alloys.—An alloy of equal parts lead and tin fuses at 186° 
and is used for soldering (soft solder). An alloy of 4-5 parts 
of lead and 1 part of antimony is very hard and answers for 
the manufacture of type (hard lead or type metal). 
The heat of formation of some of the lead compounds equals :—• 
(Pb,Cl2) = 82.7 (Pb.Br2) = 64.4 (Pb..I2) = 89.6 (Pb,0)=60.3. 
(Pb, S)=20.4 (Pb,S04) = 216.2 (Pb,N2.06) = 105.5 
From these numbers, calculating in the heat of solution, is explained 
the conduct of lead toward acids, as also the various other transposi- 
tions of its compounds (p. 2-59). 
Lead Oxide—PbO—is produced when lead is heated in 
air. After fusion it solidifies to a reddish-yellow mass, of 
rhombic scales (litharge). By careful roasting of lead, or by 
heating the hydrate or nitrate, a yellow amorphous powder, LEAD 
363 
called massicot, is obtained. Lead oxide has strong basic 
properties ; in the air it absorbs carbon dioxide and imparts 
an alkaline reaction to water as it dissolves as hydrate. Like 
other strong bases it saponifies fats. In hot potassium hydrate 
it dissolves, and on cooling crystallizes from solution in rhom- 
bic prisms. 
Lead Hydrate—Pb(OH)2—is thrown down as a white, 
voluminous precipitate by alkalies from solutions of lead salts. 
It imparts an alkaline reaction to water, as it is somewhat 
soluble, and absorbs carbon dioxide with formation of lead 
carbonate. When heated, it decomposes into lead oxide and 
water. 
Lead, or the amorphous oxide heated to 300-400°, for 
some time, in the air, absorbs oxygen and is converted into a 
bright red powder, called red lead, or minium. Its composi- 
tion most probably is expressed by the formula, Pb304; it 
is considered a compound of PbO with lead peroxide, (Pb3 
04 = 2 PbO-f Pb02). When treated with somewhat dilute 
nitric acid, lead nitrate is formed in solution, while a dark 
brown amorphous powder—lead peroxide Pb02—remains. 
The latter warmed with hydrochloric acid yields lead chloride, 
and chlorine:— 
Pb02 + 4HC1 = PbCl2 + 2H20 + Cl2 
Oxygen is disengaged when sulphuric acid acts upon it, and 
lead sulphate (PbS04) formed. When dry sulphur dioxide is 
conducted over it, lead sulphate, accompanied by glowing, is 
the result:— 
Pb02 + S02 = PbS04 
When ignited Pb02 breaks up into PbO and oxygen. 
As previously mentioned, lead dioxide, like that of tin, has an acid 
nature. When warmed with potassium hydrate, it dissolves, and on 
cooling, large crystals of potassium plumbate—K2Pb03+ 3H20—sepa- 
rate out; these are perfectly analogous to potassium stannate—K2Sn03 
-f- 3H,0. An alkaline lead oxide solution added to that of this salt 
produces a yellow precipitate (Pb304 -f- H20), which loses water upon 
gentle warming, and is converted into red lead. Therefore, the latter 
must be considered as the lead salt of a normal plumbic acid Pb(OH)4 
which corresponds to stannic Sn(0H)4 and silicic Si(0H)4 acids: — 
Pb304 = Pb2Pb04. 
Another oxide—Pb203—which is precipitated as a reddish-yellow 
powder, on the addition of sodium hypochlorite to an alkaline lead 
solution, is very probably the lead salt of metaplumbic acid: Pb2 03 — 364 
INORGANIC CHEMISTRY. 
PbPb03. Nitric acid decomposes it into lead nitrate and peroxide. 
It dissolves in hydrochloric acid without liberation of chlorine ; 
when the solution is heated this gas escapes. 
Lead Chloride—PbCl2—separates as a white precipitate, 
when hydrochloric acid is added to the solution of a lead salt. 
In cold water it is almost insoluble ; from hot water, of which 
it requires 30 parts for solution, it crystallizes in white, shiuing 
needles. At a red heat it fuses to a horn-like mass. 
Lead Iodide—Pbl 2—is thrown down as a yellow precipitate 
from lead solutions by potassium iodide; from a hot solution it 
crystallizes in shining, yellow hexagonal leaflets. 
Lead Nitrate—Pb(N03)2, obtained by solution of lead in 
nitric acid, crystallizes in regular octahedra (isomorphous 
with barium nitrate) and dissolves in 8 parts water. It melts 
at a red heat, and is decomposed into PbO, N02 and oxygen. 
Boiled with lead oxide and water, it forms the basic nitrate 
pu / N 03 
1 d\OH. 
Lead Sulphate—PbSO*, occurs in nature as Anglesite, in 
rhombic crystals, isomorphous with barium sulphate. It is 
precipitated as a white, crystalline mass by sulphuric acid 
from lead solutions. It dissolves with difficulty in water, 
more readily in concentrated sulphuric acid. Ignited with 
carbon, it is decomposed according to the following equation: 
PbS04 + C = PbS + 2C02 
Lead Carbonate—PbC03, occurs as Cerussite in nature. Is 
precipitated by ammonium carbonate from lead nitrate solu- 
tions. Potassium and sodium carbonates precipitate basic 
carbonates, the composition of which varies with the tempera- 
ture and concentration of the solution. A similar basic salt, 
which in composition generally corresponds to the formula:— 
CO;/PbOH 
2PbC03,Pb(0H)2 = 3^)Pb 
C°3\PbOH 
is prepared on a large scale by the action of carbon dioxide 
upon lead acetate. It bears the name of white lead. 
Formerly white lead was manufactured by what is termed 
the Dutch process. Rolled lead sheets were moistened in 
earthenware pots, with acetic acid, and these covered with 
manure and permitted to stand undisturbed for some time. 
In this way, by the action of acetic acid and air upon the lead BISMUTH. 
365 
a basic acetate is formed, which by the C02 evolved from the 
decaying manure is converted into basic lead carbonate. At 
present it is more rationally prepared, that is, by dissolving 
litharge in acetic acid the basic acetate is produced, and into 
this carbon dioxide is passed, for the formation of the carbonate. 
White lead is employed for the manufacture of white oil 
colors. As it is poisonous, and further blackened in air by 
hydrogen sulphide (formation of lead sulphide), it is at present 
being more and more replaced by zinc white and permanent 
white (BaS04). 
Lead Sulphide—PbS—occurs crystallized in metallic, shin- 
ing cubes and octahedra. Hydrogen sulphide precipitates it 
as an amorphous black powder. In dilute acids it is insoluble. 
The lead compounds are very poisonous. The soluble salts 
have a sweetish, astringent taste. They are readily recog- 
nized by the following reactions: sulphuric acid precipitates 
white lead sulphate, which is blackened by hydrogen 
sulphide; potassium iodide precipitates yellow lead iodide. 
BISMUTH. 
Bismuth constitutes a natural group with antimony, arsenic, 
phosphorus and nitrogen, and like these affords compounds of 
the forms BiX3 and BiX5. We observed that with increase 
of atomic weight the metalloidal character of the lower mem- 
bers becomes more metallic (see p. 139) ; the acid nature of 
the oxides passes into a basic. Antimony oxide (Sb203) is 
a base, while the higher oxide Sb205 represents an acid 
anhydride. In bismuth the metallic nature attains its full 
value. This is manifest in its inability to unite with hydro- 
gen. Bismuth trioxide is a base, and the pentoxide also possesses 
a but very feeble acid character, yielding with alkalies only 
indefinite compounds; it behaves more like a metallic per- 
oxide, and in its properties exhibits great similarity to lead 
peroxide. 
Bi = 210. 
Bismuth usually occurs native and in combination with 
sulphur, as bismuthinite. To obtain the metal the sulphide 
is roasted in the air and the resulting oxide reduced with 
charcoal. 
Bismuth is a reddish-white metal, of specific gravity 9.9. 366 
INORGANIC CHEMISTRY. 
It is brittle and may be easily pulverized. It crystallizes in 
rhombohedra. It fuses at 267° and distills at a white heat. 
At ordinary temperatures it does not change in the air; 
heated it burns to bismuth oxide—Bi203. In hydrochloric 
acid it is insoluble ; by boiling sulphuric acid it is dissolved 
to sulphate with evolution of sulphur dioxide. Nitric acid 
readily dissolves it in the cold. 
Water decomposes bismuth solutions in the same manner as 
those of antimony; insoluble basic salts are precipitated while 
acid salts pass into solution. 
Bismuthous Chloride—BiCl3—arises from the action of 
chlorine upon heated bismuth, and by the solution of the metal 
in aqua regia. It is a soft, white mass which readily fuses, 
sublimes and deliquesces in the air. Water renders its 
solution turbid, a white, crystalline precipitate of Bismuth 
Oxychloride—BiOCl—separating at the same time:— 
BiClj + H20 = BiOCl + 2HC1. 
The metalloidal character of bismuth is indicated by this 
reaction. 
The compounds BiBr3 and Bil3 are very similar to bismuth 
chloride. All three combine with many metallic chlorides to 
form double chlorides. 
Halogen derivatives of pentatomic 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 hydrate—Bi(OH)3—is not known in a free 
state. Potassium hydrate added to a bismuth solution pre- 
cipitates a white, amorphous metahydrate—BiO.OH. 
Chlorine conducted through a concentrated potassium hydrate 
solution in which there is suspended bismuth oxide, precipitates 
red bismuthic acid (BiO:iH or Bi2H407), which when gently 
heated becomes Bi205, bismuthic oxide. Strong ignition con- 
verts the latter into Bi203 and 02; by hydrochloric acid it 
is dissolved to bismuthic chloride, with liberation of chlorine. 
Bismuth Nitrate—Bi(N03)3—is obtained by the solution of 
bismuth in nitric acid, and crystallizes with 5 molecules of H20 
in large, transparent tables. In a little water it dissolves with- 
out any change ; much water renders it turbid, owing to the 
f n°3 fN03 
precipitation of white, curdy basic salts: Bi | NOsaud Bi j OH. CHROMIUM. 
367 
The precipitate is employed in medicine under the name of 
Magisterium bismuthi. 
Bismuth Sulphate—Bi2(S04)3—is formed when bismuth 
dissolves in sulphuric acid. It crystallizes in delicate needles. 
Bismuth Sulphide—Bi2S3—occurring as bismuthinite, is thrown 
down as a black precipitate from bismuth solutions by hydro- 
gen sulphide. Unlike antimony and arsenic sulphides, it does 
not form sulpho-salts. 
The alloys of bismuth are for the most part readily fusible. 
An alloy of 4 parts Bi, 1 part Cd, 1 part tin and 2 parts Pb, 
fuses at 65°. The alloy of 2 parts bismuth, 1 part lead and 1 
part tin (Rose’s metal) fuses at 94°. 
CHROMIUM GROUP. 
We observed that to the metalloidal elements, carbon, silicon 
and tin, a group of more metallic analogous elements, titanium, 
zirconium and thorium, attaches itself; further, that an analo- 
gous group of metallic elements—vanadium, niobium and 
tantalum—corresponds to the metalloidal group of phospho- 
rus (p. 137). In a perfectly similar relation to the elements 
of the sulphur group stand the metals chromium, molybde- 
num, tungsten, and somewhat removed, uranium (see Periodic 
System of the Elements). The resemblance of these elements 
to sulphur and its analogues is plainly manifest in their highest 
oxygen derivatives (see also Manganese). As the elements of 
the sulphur group in their highest oxygen compounds are 
hexatomic, so chromium, molybdenum, tungsten and uranium 
form acid oxides—Cr03, Mo03, W03, Ur03. Many of the 
salts corresponding to these are very similar to and isomorph- 
ous with the salts of sulphuric acid. The sodium chromate, 
like sodium sulphate, crystallizes with 10 molecules of water ; 
the potassium salts of both groups form isomorphous mixtures ; 
the magnesium salts, as also, that of tungstic acid, have the 
same constitution :— 
MgS04+ 7H20 and MgCr04 + 7H20. 
Corresponding to the acid oxides are the chlorine deriva- 
tives :— 
S02C12, Cr02Cl2, Mo02C12, MoOC14, W0C14 and WC16> 
which, as regards chemical deportment, are perfectly analogous 368 
INORGANIC CHEMISTRY. 
Besides these highest oxides the elements of the sulphur group 
form yet feebler acid oxides— 
IV 
so2, 
IV 
Se02 
and 
IV 
Te02. 
Of these the last approaches the bases. Their analogues in 
the chromium group : Cr02, Mo02, W02 in which the elements 
appear tetratomic, possess an undetermined, neither acid nor 
basic, character. 
The tetratomic nature of chromium also shows itself in the 
basic oxide Cr203 and its corresponding salts Cr2X6. These 
contain an hexatomic group, composed of two chromium atoms: 
IV IV 
= Cr - Cv~ 
IV IV 
Cl3Cr — CrCl3. 
The oxide derivatives of chromium, in their properties, are 
very similar to the compounds of aluminium, iron and manga- 
nese, constituted according to the same type. (See p. 341.) 
Finally, for chromium we have yet compounds of the form 
CrX2, in which it figures as a dyad. These so-called Chro- 
mous compounds are very much like the compounds of the 
metals of the magnesium group, especially the ferrous salts, 
(FeX2). They are very unstable, and by oxidation in the air 
readily pass into chromic compounds. 
The salts corresponding to the lowest grades of oxidation 
for molybdenum and tungsten are not known, as they mostly 
occur as hexads. Uranium, which has the highest atomic 
weight of the group, shows some variations from its analogues, 
which, as in similar cases, are explained by its high atomic 
weight. 
CHROMIUM. 
Chromium is found principally as chromite in nature. 
This is a combination of chromic oxide with ferrous oxide— 
Cr203Fe0—and occurs in North America, Sweden, Hungary, 
and particularly in the Ural in large quantities. Rarer is 
Crocoisite—PbCrO<. Chromic iron is almost exclusively used 
for the preparation of all other chromium derivatives, in that 
it is first converted into potassium chromate (see this) by 
fusion with potassium carbonate and nitrate. 
Metallic chromium may be isolated by very strong ignition 
Cr = 52.4. CHROMIUM. 
369 
of the oxide with charcoal. It is more conveniently obtained 
by igniting a mixture of chromium chloride, potassium chloride 
and sodium chloride with zinc, in a closed crucible. The 
separated chromium dissolves in the molten zinc, and when the 
latter is dissolved in nitric acid it remains behind as a gray, 
metallic, crystalline powder, of specific gravity 6.8. It is very 
hard (cuts glass), and extremely difficultly fusible. Heated in 
the air it slowly oxidizes to chromic oxide ; ignited in oxygen 
it burns with a bright light. In hydrochloric and warm dilute 
sulphuric acid it dissolves readily, wTith elimination of hydro- 
gen ; it is not altered by nitric acid. 
Chromium forms three series of compounds: chromous— 
CrX2, chromic—Cr2X8, and the derivatives of chromic acid 
called chromates. All chromium compounds are brightly 
colored, hence, too, the name chromium (from yptbiia, color). 
CHROMOUS COMPOUNDS. 
These are very unstable, and by oxidation readily pass into ic com- 
pounds. Like ferrous salts they are produced by the reduction of the 
higher oxides. The following may be mentioned :—Chromous Chloride 
CrCl2. This is obtained by heating chromic chloride Cr2Cl6 in a stream 
of hydrogen. It is a white crystalline powder, dissolving in water 
with a blue color; the solution attracts oxygen with avidity and 
becomes green colored. The alkalies precipitate brown chromous 
hydrate, Cr(OH)2, from it. This is oxidized by the oxygen of the 
water to Cr304 (Cr203Cr0), corresponding to magnetic iron Fe304. 
CHROMIC COMPOUNDS. 
Chromic Chloride.—Cr2Cl6, like A12C16 is obtained by 
ignition of the oxide and charcoal in a chlorine stream. At 
a red heat, in the latter it sublimes in shining violet leaflets, 
which, heated in the air, are transformed into chromic oxide. 
Pure chromic chloride only dissolves in water after very long 
continued boiling; if, however, it contains traces of CrCl2, it 
dissolves readily at ordinary temperatures. From the green 
solution on evaporation—green crystals of Cr2Cl6 -f 12H20 
separate—which deliquesce in the air. These crystals are 
obtained from solutions of chromic hydrate Cr2(OH)6in hydro- 
chloric acid. When they are dried intermediate oxychlorides 
Cr2Cl4(OH)2 and Cr2Cl2(OH)4, and at last Cr2(OH)6 result. 
Chromic Hydrate—Cr2(OH)6. The alkalies throw it down 
from chromic solutions as a voluminous, grayish blue, hydrous 370 
INORGANIC CHEMISTRY. 
precipitate. It dissolves readily in potassium and sodium 
hydrates, with an emerald green color. In an excess of am- 
monium hydrate in the cold, it dissolves with a beautiful pur- 
ple color. When the solution is boiled, chromic hydrate is 
again precipitated. Upon ignition it becomes chromic oxide. 
Chromic Oxide—Cr203—is a green, amorphous powder. 
It is formed also by the ignition of chromium trioxide:— 
or of ammonium bichromate :— 
2Cr03 = Cr203 + 30 
It is obtained in black, hexagonal crystals, if the vapors of 
the oxychloride be passed through a glowing tube:— 
(NH4)2Cr207 — Cr203 + 4H20 + N2. 
Ignited chromic oxide is insoluble in acids. Fused with sili- 
cates, it colors them emerald green, and serves, therefore, to 
color glass and porcelain. 
A particularly beautiful green-colored hydrate, finding use 
as a color, under the name of Guignet's Green is obtained by 
glowing a mixture of one part potassium bichromate with three 
parts boric acid; extracting the mass with Avater, Avhereby 
potassium borate dissolves, there remains a green powder, the 
composition of which corresponds Avith the formula:— 
2Cr02Cl2 — Cr203 -|- 2C12 0. 
Cr20(0H)4 = Cr203.2H20. 
Chromium Sulphate—Cr2(S04)3—is obtained by dissolving 
the oxide in sulphuric acid, and upon slow evaporation it 
crystallizes with 12 molecules of H20, in violet blue octahedra, 
which dissolve in Avater with a beautiful violet color. On 
Avarming the liquid its color changes to green, and upon evapo- 
ration an amorphous green mass is obtained, the green solution 
of which, upon protracted standing, becomes violet, and violet 
blue crystals of sulphate again separate. From the violet 
solution alcohol precipitates the violet salt crystalline, while 
in the green liquid no precipitation occurs. Therefore 
chromium sulphate exists in two modifications, the violet crys- 
talline and the green amorphous. 
Similarly, the other chromium compounds, Cr2Cl6, Cr2(N03)6, 
soluble in Avater, exist in tAvo different modifications. 
With the alkaline sulphates, chromium sulphate forms 
double salts—the chromium alums (p. 342). 
Potassium Chromium Alum—Cr2(S04)3K2S04 + 24H20— 
crystallizes in large, dark violet octahedra. It may be most CHROMIUM. 
371 
conveniently made by acting upon a solution of potassium 
bichromate mixed with sulphuric acid, with sulphur dioxide: 
K2Cr207 + H2S04 + 3S02 = Cr2(S04)s.K2S04 + H20. 
At 80° the violet solution of the salt becomes green, and on 
evaporation affords an amorphous green mass. 
As chromium hydrate has only a feeble basic nature salts 
with weak acids, like C02, S02, H2S (see Aluminium, p. 347) 
do not exist. Therefore the alkaline carbonates and sulphides 
precipitate chromium hydrate from solutions of chromium 
salts:— 
Cr2(S04)3 + 3Na2C03 + 3H20 = Cr2(OH)6 + 3Na2S04 + 3C02 
and 
Cr2(S04)3 + 3(NH4)2S + 6H20 = Cr2(OH)6 + 3(NH4)2S04 + 
3H2S. 
Ammonium sulphide produces a black precipitate—CrS — 
in solutions of chromous salts. 
In its highest oxygen derivative, Cr03, chromium possesses 
a complete metalloidal, acid-forming character. Chromic 
acid H,Cr04 is perfectly analogous to sulphuric acid H2S04, 
but has not been obtained free, as upon liberation from its 
salts it at once breaks up into the oxide and water :— 
DERIVATIVES OF CHROMIC ACID. 
H2Cr04 = Cr03 + H20. 
The chromates are often isomorphous with the sulphates (p. 
367) Polychromates also exist, and are derived from poly- 
chromic acids, formed by the condensation of several mole- 
cules of the normal acid (see disulphuric acid, p. 186) :— 
K2Cr04 
Potassium chromate 
K2Cr207 
Potassium dichromate 
K2CrsOj0, etc. 
Potassium trichromate 
The constitution of these salts is expressed by the following 
formulas:— 
.OK 
Cr0\ 
yO 
Cr0^ 
\0 
CrO/° 
XOK 
OK 
Cl'02\ 
)° 
Cr02( 
X°K 
/ C\\r 
Cr°2\OK 
X°K INORGANIC CHEMISTRY. 
The free polychromic acids are not known as separated from 
their salts, they immediately break up into the acid oxide and 
water:— 
Frequently, but incorrectly, the polychromates are called 
acid salts; true acid or primary salts, in which only one IT 
atom is replaced by metal, are unknown for chromic acid. 
The salts of normal chromic acid are mostly yellow colored, 
while the polychromates are red. The latter are formed from 
the former by the action of acids :— 
H2Cr3010 — 3Cr03 H20. 
Conversely, by the action of the alkalies, the polychromates 
pass into the normal salts :— 
2K2Cr04 + 2HN03 = K2Cr207 + 2KN0S + H20. 
Their formation may also be as follows: The chromic acid 
liberated from its salts by stronger acids breaks up into water 
and the acid oxide, which combines with the excess of the 
normal chromate:— 
K2Cr207 + 2K0H = 2K2Cr04 + H20 
When there is an excess of acid free oxide is formed. 
Cr04K2 + Cr03 = K2Cr207 
Chromium Trioxide—Chromic acid—Cr03. It consists of 
long, red, rhombic needles or prisms, obtained by adding 
sulphuric acid to the concentrated potassium dichromate solu- 
tion. The crystals deliquesce in the air and are readily 
soluble in water. When heated, they blacken, and at about 
250° decompose into chromic oxide and oxygen :— 
Chromium trioxide is a powerful oxidizing agent, and 
destroys organic matter; hence its solution cannot he filtered 
through paper. When alcohol is poured on the crystals, 
detonation takes place, the alcohol burns, and green chromic 
oxide remains. By the action of acids, e. g., sulphuric, the 
trioxide deports itself like a peroxide ; oxygen escapes and a 
salt results. With hydrochloric acid chlorine is evolved :— 
2Cr03 = Cr2Os +30. 
3Cr03 + 12HC1 = Cr2Cl6 + 6H20 + 3C12. 
Reducing substances, like sulphurous oxide and hydrogen 
sulphide, convert chromic acid into oxide. CHROMIUM. 
373 
Chromate of Potassium—K2Cr04—is obtained by adding 
potassium hydrate to potassium dichromate. It forms yellow, 
rhombic crystals, isomorphous with potassium sulphate (K2 
S04); from the solution of both salts isomorphous mixtures 
crystallize out. 
Bichromate of Potassium—K2Cr207—called acid potassium 
chromate, is manufactured on a large scale, and in commerce 
bears the name red chromate of potash. It is obtained by 
igniting pulverized chromite, Cr203Fe0, with potashes and 
nitre, whereby potassium chromate and ferric oxide are 
formed. The fusion is treated with water, and the resulting 
solution of potassium chromate, K2Cr04, mixed with the acetic 
or nitric acid (see above); from the concentrated solution 
potassium bichromate crystallizes. 
This method in practice is advantageously replaced by the following: 
the pulverized chromite is ignited, together with lime, in furnaces 
allowing air access. Calcium chromate (CaCr04) (together with ferric 
oxide) is produced. This dissolves in sulphuric acid to dichromate 
CaCr207 ; the latter is converted by potassium carbonate into potas- 
sium bichromate. 
Bichromate of potassium crystallizes in large, red, triclinic 
prisms, soluble at ordinary temperatures in 10 parts of water. 
When heated the salt fuses without change ; at higher heat it 
decomposes into potassium chromate, chromic oxide and 
oxygen:— 
When the salt is warmed with sulphuric acid, oxygen escapes 
and potassium chromium alum is produced :— 
2K2Cr207 = 2K2Cr04 + Cr203 + 03. 
K2Cr207 + 4H2S04 = Cr2(S04)3.K2S04 + 4H20 + 30. 
This reaction answers for the preparation of perfectly pure 
oxygen. Further, the mixture is made use of in laboratories, 
as an oxidizing agent. 
Chromate of Sodium—Na2Cr04 -f- 10H2O — forms deliques- 
cent crystals, and is analogous to Glauber’s salt (Na2S04 + 
10H2O). Barium and Strontium Chromates—BaCr04 and 
Sr(Jr04—are almost insoluble in water. Calcium Chromate 
—CaCr04—is difficultly soluble in water, and crystallizes 
like gypsum, with two molecules of water. The magnesium 
salt, MgCr04 -f- 7H20, dissolves readily and corresponds to 
Epsom salt. 374 
INORGANIC CHEMISTRY, 
The chromates of the heavy metals are insoluble in water, 
and are obtained by transposition. 
Chromate of Lead—PbCr04—is obtained by the precipita- 
tion of soluble lead salts with potassium chromate. It is a 
yelloAV, amorphous powder, which serves as a yellow paint. 
When heated it fuses and decomposes, with liberation of 
oxygen. It is employed in the analysis of carbon com- 
pounds. In nature lead chromate exists as crocoisite. 
The oxide Cr02, called peroxide, which, among others, is obtained 
by the careful ignition of chromium trioxide, is, most likely, a salt- 
like compound: Cr203Cr03 or Cr0.Cr03. Its hydrate is precipi- 
tated from chromium solutions upon the addition of potassium chro- 
mate. On warming the peroxide with hydrochloric acid, chlorine is 
evolved. 
Chromic Acid Chloranhydrides. This acid forms chloran- 
hydrides similar to those of sulphuric acid (p. 187). Corres- 
ponding to S02C12, we have chromyl chloride Cr02Cl2; and for 
f PI 
the first sulphuric acid chloranhydride, S02 < qjj is the salt 
Cr°2 { OK: 
CrO /C1 
U2 \C1 
CrO /C1 
°rUi! \OK 
CrO /0K 
2 0K 
Chromyl Chloride—Cr02Cl2—Chromium oxychloride is 
produced by heating a mixture of potassium bichromate and 
sodium chloride Avith sulphuric acid ; it distills over as a dark 
red liquid, of specific gravity 1.92, which fumes strongly in 
the air. Boils at 118° ; the vapor density equals 77.7 (H = 1), 
corresponding to the molecular formula Cr02Cl2 = 155. 
Chromyl chloride has a strong oxidizing action. With water 
it is decomposed according to the following equation :— 
Cr02Cl2 + H20 = Cr03 + 2HC1. 
If potassium dichromate be heated with concentrated hydro- 
chloric acid, there crystallizes from the cold solution the salt 
/Cl 
Cr02 called potassium chlorochromate. Heated to 100°, it 
gives up chlorine. By water it is decomposed into potassium 
chloride and chromium trioxide. 
The chromium compounds can readily be recognized by 
their color. The following reaction is very characteristic for 
chromic acid. On adding to a solution of chromium trioxide MOLYBDENUM. 
375 
or the acidified solution of a chromate, some hydrogen per- 
oxide, the red liquid is colored blue. The nature of the com- 
pound causing this coloration is not known ; usually it is 
assumed to be a higher oxide of chromium. On shaking the 
blue solution with ether, the latter withdraws the blue com- 
pound, and is, in consequence, beautifully colored. The 
ethereal solution is somewhat more stable than the aqueous. 
Both, gradually, are decolorized, with liberation of oxygen. 
MOLYBDENUM. 
Molybdenum occurs rather rarely in nature ; usually as molybdenite 
(MoS2) and wulfenite (Mo04Pb). Free molybdenum is obtained 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, and fuses at a 
higher temperature than platinum. Heated in the air it oxidizes to 
molybdenum trioxide. It is soluble in concentrated sulphuric and 
nitric acids. By the latter it is also converted into insoluble Mo03. 
Like chromium, molybdenum forms compounds of the forms MoX2, 
MoX4 and MoX6 ; besides which, also, are known derivatives in 
which it appears to act as a pentad and also a triad. 
Molybdenum Dichloride—MoC12—resulting from the trichlo- 
ride M0CI3, when heated in a stream of carbon dioxide (together with 
M0CI4) is a bright yellow, non-volatile powder. By potassium hydrate 
it is converted into the hydrate Mo(OH2), a black powder. 
Molybdenum Trichloride—MoC13 or Mo2C16—produced by 
gentle heating (at 250°) ofMoCl5 in a current of H or CO2- A reddish- 
brown powder which, upon strong ignition, yields a dark blue vapor. 
Soluble in concentrated sulphuric acid, with beautiful blue color By 
potassium hydrate it is converted into the hydrate Mo(OH)g or 
Mo2(OH)6, which forms salts with acids. Upon glowing the hydrate 
there results the black oxide Mo203. Strong heating of the trichloride 
in a current of C02 leaves MoC12 and it sublimes. 
Molybdenum Tetrachloride—M0CI4—a brown, crystalline 
powder, which appears by evaporation to break up into MoC15 and 
MoC13. With ammonium hydrate it yields a hydrate, forming salts 
with acids. The brown solution of the salts readily assumes a blue 
color by oxidation in the air. The ignition of the hydrate leaves the 
dioxide Mo02, which is converted by nitric acid into the trioxide 
Mo03. Molybdenum disulphide, MoS2, is got by ignition of the tri- 
sulphide MoS3 away from air as a shining black powder, which occurs 
native as molybdenite, in hexagonal crystals, with a specific gravity of 
4.5. 
Molybdenum Pentachloride—MoC13—is prepared by heating 
MoS2, or molybdenum in chlorine gas. It is a metallic, shining, black, 
crystalline mass, fusing at 194° and distilling at 268° ; the vapor density 
Mo = 95.8. 376 
INORGANIC CHEMISTRY. 
equals 136.5, corresponding to the molecular formula MoCl5 = 273. 
It fumes and deliquesces in the air, and dissolves with brown color in 
water, with hissing. Soluble in alcohol and ether with a dark green 
color. 
The hexachloride MoC16 is not known, but the oxychlorides 
MoOC14 and Mo02C12 are. The first results from the ignition of 
Mo02 and carbon in a stream of chlorine, and is a green crystalline 
mass subliming under 100° and yielding a dark red vapor. Bromine 
forms perfectly analogous compounds with molybdenum. 
Molybdenum Trioxide—Mo03—results by roasting metallic 
molybdenum or the sulphide in the air. It is a white, amorphous 
mass, which turns yellow on heating ; it fuses at a red heat and then 
sublimes. It is insoluble in water and acids ; but dissolves readily in 
the alkalies and ammonium hydrate. When fused with the alkaline 
hydrates or carbonates salts are produced which are partly derived 
from the normal acid, H2Mo04, and partly from the polyacids, and 
correspond to the polychromates :— 
The ammonium salt—(NH4)2Mo04—is obtained by dissolving the 
trioxide in concentrated ammonium hydrate. 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 
(NH4)6Mo7024 + 4H20 crystallizes out. Both salts are decomposed 
by heat, leaving molybdenum trioxide. 
Hydrochloric acid added to a concentrated solution of a molybdate 
causes the precipitation of molybdic acid—H2Mo04. It is a white, 
crystalline compound, readily dissolved by an excess of acid. Zinc 
added to this solution causes it to become blue and then green, in 
consequence of the formation of lower oxides (like Mo308 = 2Mo03, 
Mo02). 
Molybdic acid can also form polyacids with phosphoric and arsenic 
acids, e. g., H3P04, 10MoO3. These complex phosphomolybdic 
acids are distinguished by the fact that they form insoluble salts with 
the metals of the potassium group, with ammonia and with organic 
bases in dilute acids. On adding to the nitric acid solution of ammo- 
nium molybdate a solution containing phosphoric (or arsenic) acid, 
there is produced a yellow crystalline precipitate of ammonium 
phospho-molybdate—(NH4)3PO4,10MoO3 + 1jH20. This reaction 
serves for the detection and separation of phosphoric acid. 
K2Mo04, K2Mo207, K2Mo3O10, Na2Mo4013, K6Mo702, etc. 
Molybdenum Trisulphide—MoS3—is thrown down as a brown 
precipitate from acidulated molybdenum solutions by hydrogen sulphide. 
It dissolves in alkaline sulphides forming sulpho-salts. Ignited away 
from air it is converted into molybdenum disulphide, MoS2, which 
occurs native as molybdenite. TUNGSTEN. 
3. TUNGSTEN. 
Tungsten is found in nature in the tungstates: as wolframite, 
FeW04, as scheelite, CaW04, and as stolzite, PbW04. 
W — 184. 
The metal is obtained, like molybdenum, by the ignition of the 
oxides in a stream of hydrogen. It has a grayish-yellow color, is very 
hard and difficultly fusible ; specific gravity 16.6. It becomes trioxide 
when ignited in the air. 
Wolfram forms the following chlorides WC12, WC14, WC15 and 
WC16. 
The Dichloride—WC12—arises by strong ignition of WC16 and 
WCI4 in a C02 stream, and forms a bright gray, non-volatile mass, 
that yields the brown oxide, W02, with water. 
The Tetrachloride—WC14—obtained by gentle ignition of WCI6 
and WC15, in an H or C02 stream, is grayish-brown and upon subli- 
mation decomposes into WC12 and volatile WC15. With water it 
forms a brown oxide. 
The Pentachloride—WC1S—is formed by the distillation ofWCl6 
in an II or C02 stream, and forms shilling, black, needle-like crystals. 
It fuses at 248° and boils at 275°, forming a dark brown vapor, with 
the vapor density 180.4 (WC15 = 360.8). With water it gives an olive 
green solution and a blue oxide, W205 ; in carbon disulphide it dis- 
solves with a deep blue color. 
Tungsten Hexachloride — WC16—arises when heating the 
metal or a mixture of wolframite with carbon in a current of chlorine. 
It forms a dark violet, crystalline mass, fusing at 245° and boiling at 
346°. The vapor density equals 198 (WC16 = 396.4). In CS2 it dis- 
solves with a reddish-brown color ; with water it forms W03. 
The Oxychloride—WC140—consists of red crystals, fusing at 
210° and boiling at 227°; its vapor density equals 170 (WC140 = 
341.6). The Dioxychloride—VVC1202—sublimes in bright yellow, 
shining leaflets. 
Tungsten Trioxide—WO3—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 hydrates. From the 
cold solution, however, tungstic acid, WO(OHi.t, is precipitated, which 
on standing over sulphuric acid becomes, W02(0H)2 and at 100° 
passes into di-tungstic acid, W207H2 = W205(0H)2. 
The salts of tungstic acid are perfectly analogous to the molybdates 
and are derived from the normal acid or the polyacids. The normal 
sodium salt, Na2 WO4 + 2H.,0, and the so-called meta-tungstate of 
sodium, Na2W4013 -f- 10H2O, are practically applied. Materials satu- 
rated with their solutions do not fall into a flame, but smoulder away 
slowly. 
Like molybdic acid, it combines with phosphoric and arsenic acids. 
The metal is used in the manufacture of tungsten steel ; a slight 
quantity of it increases the hardness of the latter very considerably. 378 
INORGANIC CHEMISTRY 
4. URANIUM 
In nature it occurs chiefly as uraninite, a compound of uranic and 
uranous oxides—U02, 2U03 = U3Og. 
The metal is obtained by heating uranous chloride with sodium. It 
has a steel-gray color and a specific gravity of 18.3. When heated in 
the air it burns to uranous-uranic oxide. There are two series of 
uranium compounds. In the one, the metal is a tetrad.UX4 ; these 
uranous or urano compounds are very unstable, and readily pass into 
the uranic or derivatives of hexatomic uranium. Uranous oxide is of 
a basic nature, and only forms salts with acids. 
The compounds of hexatomic uranium are called the uranic com- 
pounds. U03 and U02(0H)2 have a predominant basic character, 
and are called uranates, but are also capable of forming salts with 
bases. In the salts formed from acids like U02(N03)2 and U02S04 
the group U02 plays the rSle of a metal; it is called uranyl, and its 
salts are termed uranyl salts. They may also be regarded as basic 
salts. 
U = 240. 
URANOUS COMPOUNDS. 
Uranous Chloride—UCl4—is obtained by heating metallic ura- 
nium in a stream of chlorine, or uranous oxide in hydrochloric acid. 
It forms metallic, shining, dark green octahedra. In the air it deli- 
quesces, and dissolves in water with hissing. On evaporation of the 
solution, uranous hydrate remains. 
Uranous Oxide—U02—is formed when the other oxides are 
heated in a current of hydrogen. It is a black powder, which, when 
heated in the air, becomes uranous-uranic oxide U022U03. 
Uranous oxide dissolves with a green color in hydrochloric and con- 
centrated sulphuric acids. Uranous sulphate, U(S04)2 + 8H20, 
consists of green crystals. From the salts the alkalies precipitate a 
brown powder U(0H)4. 
HEXATOMIC URANIUM COMPOUNDS. 
Uranium Hexachloride—UC16—could not be obtained, but 
the oxychloride U02C12 (Uranyl chloride) exists—a yellow crystalline 
mass, deliquescing in the air—and is obtained by heating U02 in dry 
chlorine gas. 
Uranic Oxide U03 or Uranyl oxide—is a yellow powder, 
and is obtained by heating uranyl nitrate to 250°. Warmed with nitric 
acid it becomes iiranylhydrate or uranic acid, U02(0H)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 adding MANGANESE. 
379 
sulphuric acid to the solution, there crystallizes, on evaporation, 
Uranylsulphate—U02,S04 -f- 6H20—in lemon yellow needles. 
If sodium or potassium hydrate be added to the solutions of uranyl 
salts, yellow precipitates of the uranates—U207K2 and U207Na2 are 
obtained. These are soluble in acids. In commerce the sodium salt 
is known as uranium yellow, and is employed for the yellow coloration 
of glass (uranium glass) and porcelain. 
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 
IV 
as uranous uranate—2U03.U02=(U02,02)2 U. 
Many uranium salts exhibit magnificent fluorescence. The oxide 
colors glass fluxes a beautiful greenish yellow (uranium glass). 
Uranous oxide—U02—imparts a beautiful black color to glass and 
porcelain. 
Besides these compounds, in which it appears tetratomic and hexa- 
tomic, uranium forms, like molybdenum and tungsten, also a penta- 
chloride UC15. The same results by leading chlorine gas over a 
moderately heated mixture of carbon with one of the uranium oxides, 
and forms dark needles, which, in direct light, are metallic green, but 
in transmitted, ruby red. In the air it deliquesces to a yellowish green 
liquid; upon heating it is dissociated into UC14 and Cl (at 120°-235°). 
There is also a tetroxide, U04, which, like the trioxide, U03, 
forms salts with the bases. 
MANGANESE 
In proportion to its atomic weight, manganese stands in 
the same relation to the elements of the chlorine group as 
chromium to the elements of the sulphur group. The rela- 
tionship manifests itself distinctly in the higher states of 
oxidation. Permanganic oxide, Mn207, and acid, HMn04, 
are perfectly analogous to Cl207(or I207) and HC104. The 
permanganates and the perchlorates are very similar, and for 
the most part isomorphous. In them manganese, like the halo- 
gens in their highest state of oxidation, appears as a heptad. 
To this, however, is limited the similarity of manganese with 
the halogens. In its remaining compounds it shows great 
resemblance to the elements standing in the same horizontal 
series of the periodic system, viz., with iron and chromium 
(p. 341). Like these two elements it forms three series of 
compounds. 
1. In the manganous derivatives—MnX2—the metal is 
diatomic. These salts are the more stable, and comprise the 
Mil = 54.8. 380 
INORGANIC CHEMISTRY. 
most common manganese compounds. They are very much 
alike, and usually isomorphous with the ous salts of iron and 
chromium, and the salts of metals of the magnesium group. 
2. The manganic compounds—Mn2X6—are similar to and 
isomorphous with the ferric, chromic and aluminium deriva- 
tives ; however, they are less stable, and easily reduced to the 
manganous state. Their composition is due to the tetratomic 
nature of manganese (p. 368). 
3. The derivatives of manganic acid—H2Mn04 = Mn02 
(OH)2 in which manganese is hexatomic, are analogous to those 
of ferric (H2Fe04) and chromic (H2Cr04), also, of course, to 
those of sulphuric acid (H2S04). 
Consequently, in manganese we plainly observe how the 
similarity of the elements in their compounds is influenced by 
the valence (see p. 324). In its ous condition manganese, 
like the elements of the magnesium group, has a rather strong 
basic character, which in the ic state diminishes considerably. 
Hexatomic manganese has a metalloidal, acidic character, and 
manganic acid approaches sulphur. By further increase of 
oxygen manganese finally (in permanganic acid) acquires the 
metalloid character of the halogens. A similar behavior, as 
we have already noticed, is shown by many other metals, 
especially chromium and iron. 
On the other hand, by addition of hydrogen or groups 
(CH3.C2H5), the metalloids and the weak basic metals acquire a strong 
basic, alkaline character. The groups, NH4 (ammonium), P(CH3)4 
(tetramethylphosphonium), S(C2H5)3 (triethylsulphine), Sn(C2H5)3 
(tin triethyl), etc., are of metallic nature, because their hydrates, 
P(CH3)4.OH, S(C2H5)3.OH, Sn(C2H5)3.OH, are perfectly similar 
to the alkaline hydrates (KOH,NaOH). 
Manganese is widely distributed in nature. Native, it is 
found in meteorites. Its most important ores are pyrolusite, 
Mn02, hausmannite, Mn304, braunite, Mn203, manganite, 
Mn203.H20, and rhodochrosite, MnCOs. 
Metallic manganese is obtained by igniting the oxides with 
charcoal. It has a gravish-white color, is very hard, and 
difficultly fusible ; specific gravity 7.2. In moist air it oxi- 
dizes readily. It decomposes water on boiling, and when dis- 
solved in acids forms manganous salts. 
The heat of formation of the most important manganese compounds 
corresponds to the symbols :— 
(Mn,0,H20) = 94 7 
(Mn,02,H20) = 116.2 
(Mn,Cl2) = 111.9 
(Mn,S,04) = 249.8 
(Mn,Cl,,Aq) = 128.0. 
(Mn,04,K) = 194.8. MANGANESE. 
381 
Manganous Oxide—MnO—results from ignition of the car- 
bonate, 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 Mn304. 
Manganous Hydrate—Mn(OH)2—is a voluminous, reddish- 
white precipitate, formed by the alkalies in manganous solu- 
tions. Exposed to the air, it oxidizes to manganic hydrate 
Mna(OH)6. 
Manganous salts mostly have a pale, reddish color, and are 
formed by the solution of manganese or manganic oxides in 
acids. 
Manganous Chloride—MnCl2—crystallizes with four mole- 
cules of water in reddish tables. On drying, it is decomposed 
with separation of hydrochloric acid. Anhydrous manga- 
nous chloride is made by heating the double salt MnCl22NH4Cl 
-j- H20 (see Magnesium chloride) or by heating manganese 
oxides in hydrochloric acid ; it is a crystalline, reddish mass, 
which deliquesces in the air. 
Manganous Sulphate—MnS04—crystallizes below +6° with 
7 molecules of H20 (like magnesium and ferrous sulphates), at 
ordinary temperatures with 5H20 (like copper sulphate); the 
last molecule of water does not escape until 200°. It forms 
double salts with the alkaline sulphates, e. g., MnS04. K2S04 
+ 5H20. 
Manganous Carbonate—MnC03—exists in nature as rhodo- 
cbrosite, and is precipitated by alkaline carbonates from man- 
ganous solutions, as a white powder, which, on exposure turns 
brown. 
Manganous Sulphide—MnS—is found in nature as alaban- 
dite or manganese blende. Alkaline sulphides precipitate a 
llesh-colored sulphide from manganese solutions. In the air 
it becomes brown. 
MANGANOUS COMPOUNDS. 
Manganic Oxide—Mn203—is a black powder produced by 
the ignition of the manganese oxides in a current of oxygen 
gas. 
Manganic Hydrate—Mn2(OH)6—is precipitated as a black- 
ish-brown powder, from manganic solutions, by ammonium 
MANGANIC SALTS. 382 
INORGANIC CHEMISTRY. 
hydrate. Warmed with hydrochloric acid the oxide and 
hydrate yield manganous chloride with evolution of chlorine: 
Mn203 + 6HC1 = 2MnCI2 + 3H20 + Cl2. 
Manganous-manganic Oxide—Mn304 = Mn0,Mn203. It 
constitutes the mineral hausmannite crystallized in dark brown 
quadratic octahedra, and is formed by the ignition of all 
oxides in the air; it is isomorphous with magnetite, Fe304. It 
reacts with hydrochloric acid, according to the equation :— 
Manganic oxide, like the other sesquioxides, is only a very 
feeble base; its salts are very unstable, and with separation 
of oxygen readily become manganous salts. 
Mn304 + 8HC1 = 3MnCl2 + 4H20 + Cl2. 
Manganic Sulphate—Mn2(S04)3—is a dark green powder. 
It is produced when concentrated sulphuric acid is poured 
over hydrated manganic peroxide. It deliquesces in the air 
and dissolves in cold water with a dark red color. It forms 
alums with potassium and ammonium sulphates—e. g., 
Mn2(S04)3.K2S04 + 24H20. 
When the solutions are heated both salts decompose and 
yield up oxygen. 
Manganese Dioxide—Mn02—peroxide. This is the mineral 
pyrolusite occurring in dark-gray radiating masses, or in 
almost black rhombic prisms, which possess metallic lustre. 
When gently heated it is converted into oxide, by strong 
glowing, into manganous-manganic oxide :— 
3Mn02 = Mn304 + 20. 
It is used for making oxygen. When warmed with hydro- 
chloric acid chlorine escapes :— 
Artificially the dioxide may be obtained in the form of 
hydrates—Mn02,H20,Mn02,2H20—if a hypochlorite be added 
to the solution of the manganese salt, or if chlorine be con - 
ducted through a solution of manganese containing sodium 
carbonate. In cold hydrochloric acid the precipitated dioxide 
dissolves without liberating chlorine, as MnCl4 probably is 
formed, which upon application of heat breaks up into MnCl2 
and Cl2. This deportment would indicate that manganese is 
tetratomie in the dioxide. 
Mn02 + 4HC1 = MnCl2 + 2H20 + Cl2 MANGANESE. 
Manganese peroxide (as also Mn203 and Mn304), serves chiefly 
for the manufacture of chlorine gas, and it is, therefore, technically im- 
portant to estimate the quantity of chlorine which a given oxide of 
manganese is able to set free. This is done by boiling the oxide with 
hydrochloric acid, conducting the liberated chlorine into a potassium 
iodide solution, and determining the separated equivalent amount of 
iodine by means of sodium hypochlorite. Or the oxide is heated in a 
flask with oxalic and sulphuric acids, when the first is oxidized to 
carbon dioxide. From its quantity we can calculate the quantity of 
active or available oxygen in the manganese oxide. 
In the chlorine preparation the manganese is found as manganous 
chloride in the residue. With the relatively high value of pyrolusite, 
it is important for trade that the peroxide be recovered from the 
residue. This regeneration is executed, at present, upon a grand 
scale, according to Weldon’s method, in the following manner. The 
manganous chloride, containing excess of hydrochloric acid, is neu- 
tralized with lime, the clear liquid brought into a high iron cylinder 
(the oxidizer), milk of lime added and air forced in. The mixture 
becomes warm, and so-called calcium manganite MnOsCa — Mn02 
CaO is precipitated as a white mud:— 
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 + CaO. 
MnCl2 -f- 2CaO -f- 0 = Mn03Ca + CaCl2 
COMPOUNDS OF MANGANIC AND PERMANGANIC 
ACIDS. 
When oxygen compounds of manganese are heated in the 
air, in contact with potassium hydrate, or better, with oxidizing 
substances, like nitre or potassium perchlorate, a dark gray 
amorphous mass is produced, and this dissolves in cold water 
with a dark green color. When this solution is evaporated 
under the air-pump, dark green metallic rhombic prisms of 
'potassium manganate—K2Mn04—crystallize out. This salt is 
isomorphous with potassium sulphate and chromate. In 
potassium or sodium hydrate, potassium manganate dissolves 
without alteration ; it is, however, decomposed by water, 
hydrated manganese dioxide separating, and the green solution 
of the manganate changing into a dark red solution of the 
permanganate: KMu04— 
A similar conversion of the green manganate to red per- 
manganate occurs more rapidly by the influence of acids :— 
3K2Mn04 + 3H20 = 2KMn04 + Mn02. H20 + 4K0H. 
3K2Mn04 + 4HN03 =2KMn04 + Mn02 + 4KN03 + 2H20. 384 
INORGANIC CHEMISTRY. 
Owing to this ready alteration in color the solution of the 
manganate, is called mineral chameleon. 
Potassium Permanganate—KMn04—is best prepared by 
conducting C02 into the manganate solution until the green 
color has passed into a red. On concentration the salt crystal- 
lizes in dark red rhombic prisms isomorphous with potassium 
perchlorate, KC104. It is soluble in 12 parts of water at 
ordinary temperatures. 
The permanganate solution is strongly oxidizing, as it con- 
verts lower oxygen compounds into higher; in this manner 
the permanganate is reduced to a colorless manganous salt. 
When a permanganate solution is added to an acidulated 
ferrous solution, the former is decolorized, and there results a 
feeble yellow solution of ferric and manganic salts :— 
2KMn04 + 10FeS04 + 8H2S04 = 2MnS04 + 5Fe2(S04)3 + 
8H20 + K2S04. 
Hence the permanganate solution serves in volumetric 
analysis for the quantitative estimation of ferrous salts. 
In the same manner, the permanganate oxidizes and 
destroys many organic substances, therefore the solution can- 
not be filtered through paper ; it serves as a disinfectant. 
The permanganate is also reduced by hydrogen peroxide 
(p. 93); the reaction proceeds according to the following 
equation:— 
Mn207K20 + 5H202 = 2MnO + K20 -|- 5H20 +502; 
the formation of oxides requires, for completion of the reaction 
the presence of acids (sulphuric acid). 
The remaining permanganates are also similar to and 
isomorphous with the perchlorates. The sodium salt is very 
soluble in water, and crystallizes but poorly. 
Strongly cooled sulphuric acid added to dry permanganate 
causes the separation of Manganese Heptoxide—Mn207—an 
oily, dark-colored liquid. By careful warming it is converted 
into dark violet vapors, Avhich explode on rapid application of 
heat. Manganese heptoxide has a violent oxidizing action; 
paper, alcohol and other organic matter are inflamed by 
mere contact with it. METALS OF GROUP VIII. 
385 
METALS OF GROUP VIIT. 
Of all known elements, there remain yet for consideration 
those standing in the VIII column of the periodic system. 
They arrange themselves into the following three groups, 
the members of which possess almost agreeing atomic weights 
and specific gravities :— 
Fe = 56.0 
Ru = 108.0 
Os = 198.0 * 
Co = 58.6 
Rh = 104 0 
Ir = 192.8 
Ni = 58.6 
Pd = 106.7 
Pt = 196.7 
These elements represent the middle members of the three 
great periods, for which, in the two small periods, no analogues 
exist (pp. 234, 236). 
Both as regards atomic weights and in their entire chemical 
character, the elements of these three groups constitute a 
transition from the preceding members of the great periods 
(Mn and Cr Mo W) to the next following members (Cu Ag 
Au and Zn Cd Ag). The elements standing side by side 
(heterologous) and belonging to the same periods are very 
similar in their physical properties, and show, e. g., very close 
specific gravities. Therefore these are included usually in 
groups and distinguished : (1) the iron group (Fe Co Ni), with 
the specific gravity 7.8-8.6 ; (2) group of the light platinum 
group metals (Ru Rh Pd), with specific gravity 11.8-12.1, 
and (3) the group of the heavy platinum metals (Os Ir Pt), 
of specific gravity 21.1-22.4. 
On the other hand, the homologous elements (Fe Ru Os, 
Co Rh Ir and Ni Pd Pt), according to their chemical prop- 
erties, show a like similarity, as all the other homologous 
groups, and therefore may be considered in such groups. This 
resemblance shows itself chiefly in their combination forms, 
and, of course, too, in the properties of the compounds (p. 324). 
We observed that the metals of group VI (chromium, molybde- 
num, wolfram) and of group VII (manganese) form the highest 
oxides (Me03 and Me307) of an acidic nature. Agreeing with 
*As already mentioned (p. 238), osmium must, corresponding with 
its position in the periodic system, have a lower atomic weight than 
that experimentally determined (198). As, however, in its entire 
deportment, it bears the same relation to Ir and Pt as Ru to Rh and 
Pd and Fe to Co and Ni, therefore it is extremely probable that its 
atomic weight is not correctly determined, and will show itself less, 
somewhat, than that of iridium (192.7). 386 
INORGANIC CHEMISTRY. 
this, there exist for the adjacent elements of group VIII (iron, 
ruthenium and osmium) salts— 
Fe04K2, Ru04K2) 0s04K2, 
which are derived from the unstable trioxides Fe03, Ru03 and 
0s03. This acid-forming function disappears in the follow- 
ing members, Co Rh Ir and Ni Pd Pt, the chemical valence 
of which rapidly diminishes, and which attach themselves to 
Cu Ag and Au. 
Consequently the whole physical and chemical deportment 
of the 9 elements to be considered is governed by their posi- 
tion in the periodic system. 
As previously mentioned pp. 236, 240 and at other places, 
the valences of the elements in their highest salt-forming 
oxides present themselves as periodic functions of the atomic 
weights. A similar dependence is also seen in the lowest salt- 
forming oxides, as is visible from the following tabulation of 
the highest and lowest salt-forming oxides of the middle 
members of the great periods :— 
v205 
VI 
CrO 3 
VII 
Mn207 
VI 
Fe03 
IV 
Co203 
II 
NiO 
11 
CuO 
11 
ZnO 
III 
Ga20 3 
VI 
hi 
v203 
11 
CrO 
MnO 
11 
FeO 
11 
CoO 
Cu20 
II 
Nb205 
III 
(Nb203) 
VI 
Moo3 
MoO 
VI 
Ru03 
II 
RuO 
Rh203 
RhO 
II 
PdO 
Ag20 
n 
CdO 
in 
ln203 
IV 
SnO 2 
11 
SnO 
Ta205 
VI 
W03 
VI 
0s03 
IV 
Ir02 
IV 
Pto2 
III 
Au203 
II 
HgO 
III 
ti2o8 
IV 
Pb02 
III 
(Ta203) 
ii 
(WO) • 
ii 
OsO 
II 
IrO 
ii 
PtO 
I 
Au20 
Hg20 
I 
T120 
II 
PbO 
METALS OF THE IRON GROUP. 
The metals of this group, iron, cobalt and nickel, form a 
gradual transition from manganese to copper. Their magnetic 
properties mostly distinguish them from the other elements. 
Iron forms three series of compounds after the forms, Fe03. 
Fe203 and FeO. 
In its highest combinations iron has an acidic character, 
and the derivatives of ferric acid (H2Fe04) are entirely simi- 
lar to those of chromic and manganic acids ; they are, however, 
less stable than the latter. Their analogues with cobalt and 
nickel are not known. IRON. 
387 
The ferric compounds—Fe2X6—containing the hexatomic 
VI 
group Fe2 (p. 368) are much like the aluminium, chromic and 
permanganic derivatives. They are usually isomorphous. 
Among iron salts they are characterized by their relative 
stability. The highest oxides of cobalt are much less stable, 
and only a few double salts of this form are known, while 
the higher salts with nickel are not known. 
Thirdly, iron, cobalt and nickel afford ous compounds, 
(FeX2, CoX2, NiXa) in which they appear to be dyads. They 
are very much like 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 copper 
and zinc. 
1. IRON. 
This metal, which is practically of such great importance, is 
very widely distributed in nature. Native, it is found on the 
earth’s surface almost exclusively in meteorites ; however, it 
is present in great masses in other worlds which (like the 
sun) are surrounded by an atmosphere of hydrogen. 
The most important iron ores are: magnetite 
hematite (Fe203), brown iron ore and limonite (hydrate of 
the oxide), siderite (FeC03). These are the almost exclusive 
materials for the manufacture of iron ; the sulphur ores, like 
pyrite, are not adapted to this end. 
In commerce there are three varieties of iron, cast iron, 
steel and wrought iron, which are principally chemically dis- 
tinguished by their varying quantities of carbon. 
Cast iron contains 3-6% carbon, partly chemically com- 
bined and partly mechanically mixed, in the form of graphite. 
Molten cast iron rapidly cooled yields the so-called white iron, 
in which the greater portion of the carbon is chemically com- 
bined with the iron. It has a whitish color, upon fracture 
exhibits a granular crystalline structure, and is very hard 
and brittle. Its specific gravity is 7.1. At about 1200° 
it fuses to a pasty mass. The chemically combined carbon in 
it can easily be removed by oxidation, for which reason it is 
suitable for the making of steel or wrought iron. 
Fe = 56. 388 
INORGANIC CHEMISTRY. 
Manganiferous cast iron has a leafy structure, and is called 
Spiegeleisen. It serves principally for the manufacture of 
Bessemer steel. By slow’ cooling of molten cast iron, the 
greater part of the carbon within separates, in the form of 
small leaflets of graphite. The thus produced gray cast iron 
has a darker gray color, is not so hard and brittle, fuses more 
readily (about 1150°) than white cast iron, and serves for the 
manufacture of castings. Owing to their brittleness, neither 
variety can be forged or welded. 
Steel contains 0.8-1.8 of carbon, which is entirely chemi- 
cally combined with the iron. It has a steel gray color, and 
a fine-grained structure; its specific gravity equals 7.6-8,0. 
It is more difficultly fusible (about 1400°) than cast iron, but 
easier than wrought iron. When gloiviug steel is rapidly 
cooled, it becomes very hard and brittle. In this process 
more carbon is chemically combined. Slowly cooled it is soft 
and malleable, and may be forged and welded. The weld- 
ability diminishes with the addition of carbon. 
Wrought Iron contains the least amount of carbon, 0.2- 
0.6 %. It possesses a bright gray color, has a specific gravity 
of 7.6, is rather soft and tough, and, at a red heat, may be 
readily forged, rolled and welded. The rolled iron possesses 
a radiating layer, while the forged is fine grained ; the former 
is more compact and tenacious. Wrought iron fuses at a 
bright white heat (1500°), while cast iron and steel begin to 
melt at 1200°. 
Metallurgy of Iron.—The extraction of iron from its oxygen 
ores is based upon the reduction of the same by carbon at a glowing 
heat. According to the oldest course, the ores were heated with carbon 
in wind furnaces ; in this way, by the excess of air the greater portion 
of the carbon was consumed and at once we got an iron poor in carbon, 
wrought iron, a spongy mass, which was then forged under the hammer. 
Since the beginning of the previous century, the present methods were 
adopted, according to which cast iron is first prepared from the 
ores, and this afterwards converted into steel or wrought iron. 
The smelting of the ores is executed in large, walled blast furnaces, 
which permit an uninterrupted process. These are filled, from open 
ings above, with alternating layers of coal, broken ore and fluxes 
containing silica and lime ; these latter facilitate the melting together 
of the reduced iron. The necessary air for the process is blown into 
the contracted portion of the furnace by means of a blast engine. By 
the combustion of the coal, carbon monoxide is produced, which 
reduces the iron oxide to metal:— 
Fe203 + 3C0 = 2Fe + SCO,. IRON. 
389 
In sinking in the furnace the reduced iron comes in contact with the 
coal, takes up carbon and forms cast iron, which on further sinking 
fuses and flows into the hearth of the furnace. By protracted and strong 
heating, whereby the at first chemically united carbon passes into the 
graphitic, the formation of the gray cast iron is accelerated. The 
earthy impurities of the ores combine with the fluxes, forming readily 
fusible slag, which envelops the fused iron and protects it from further 
oxidation. 
To convert the cast iron thus produced into steel or wrought iron, 
carbon must be withdrawn from it. In making the wrought iron the 
cast iron is fused in open hearths (refining process), or in reverbera- 
tory furnaces with air access, and the mass stirred thoroughly 
until it has become semi-pasty (puddling process). In this way almost 
all the carbon is burned to carbon monoxide and the other admixtures, 
like silicon, sulphur, and phosphorus, present in small quantities, 
oxidized. The wrought iron is then worked up by rolling, or under 
the iron hammers (bar iron). 
Formerly, steel was manufactured from wrought iron (not cast iron), 
by cementation. The iron bars were exposed, together with fine char- 
coal, for some time, to a red heat, whereby the iron takes up carbon from 
the surface. The bars are then reforged, again heated with fine char- 
coal, and the process repeated until the mass has become as homoge- 
neous as possible (cementation steel). It can be obtained more uni- 
form by fusion in crucibles (cast steel). 
At present, steel is, chiefly, prepared directly from cast iron, by the 
method invented by Bessemer, somewhere in 1850. It consists in 
blowing air, under high pressure, into the molten iron, until the neces- 
sary amount of carbon has been consumed (Bessemer steel). 
Puddle steel is obtained from cast iron through the puddling pro- 
cess, by the less extensive removal of carbon ; another variety, by the 
fusion of cast iron with iron ore and pyrolusite. 
Induced by the different more recent processes for manufacturing 
steel and wrought iron, whereby the same are obtained in a liquid 
condition, bearing in mind that steel differs from wrought iron essen- 
tially in its hardness, that the so-called Bessemer steel, however, is 
not tempered—there has been accepted, lately, a new division and 
nomenclature of the ideas of steel and wrought iron (which, in con- 
trast to cast iron, are difficultly fusible and workable). At present 
we distinguish— 
(1) Weld iron as a non-fused, non-tempered mass, formerly 
wrought iron ; (2) weld steel, not fused, tempered, formerly puddle 
steel ; (3) ingot iron, fused, not tempered, formerly Bessemer steel; 
(4) ingot steel, has been fused and tempered. 
The ordinarily occurring iron, even the purest wire, always 
contains foreign ingredients, principally carbon, and in minute 
quantities silicon, sulphur, phosphorus and nitrogen, nickel, 
cobalt, titanium and others. The quantity of manganese is 390 
INORGANIC CHEMISTRY. 
purposely increased (to 30%), as the iron by this means 
acquires technically valuable properties; it becomes more 
compact and solid. When iron containing carbon is dissolved 
in hydrochloric acid the chemically combined carbon unites 
with hydrogen, to form hydro-carbons, while the mechanically 
admixed graphite remains behind. The whole quantity of 
carbon is determined by solution of the iron in bromine or 
cupric chloride, when all the carbon remains behind. 
To prepare chemically pure iron, pure oxide or the oxalate 
is heated in a current of hydrogen— 
Fe203 + 3H2 = 2Fe + 3H20 ; 
the iron then remains as a fine black powder. In case the 
reduction occurs at a red heat the powder glows in the air and 
burns (pyrophoric iron). The strongly ignited powder is not 
inflammable. Iron obtained by the electrolysis of ferrous 
sulphate contains some hydrogen. 
Chemically pure iron has a grayish-white color, is tolerably 
soft, and changes but slowly in the air. Its specific gravity is 
7.78. It melts in an oxy-hydrogen flame (above 1500°). 
Ordinary iron rusts rapidly in moist air, as it covers itself 
with a thin layer of ferric hydrate. Glowed in the air it is 
coated with a layer of ferrous-ferric oxide (Fe304) which is 
readily loosened. In oxygen it burns with an intense light. 
In contact with a magnet iron becomes magnetic ; but only 
steel retains the magnetism, while cast iron and wrought iron, 
after the removal of the magnet, soon lose the property. 
Iron decomposes water at a red heat, with the formation of 
ferrous-ferric oxide, and the liberation of hydrogen:— 
3Fe -\- 4H20 — -f- 4H2. 
The metal dissolves without trouble in hydrochloric and 
sulphuric acids, with evolution of hydrogen ; the latter has a 
peculiar odor, due to hydrocarbons simultaneously produced. 
In nitric acid, iron dissolves with separation of nitrogen oxide. 
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 due to the production of oxide 
upon its surface. IRON. 
391 
These are formed by the solution of iron in acid, and may 
also be obtained by the reduction of ferric salts :— 
FERROUS COMPOUNDS. 
In the hydrous state they are commonly colored green; 
in the air they oxidize to ferric salts :— 
Fe2Cl+ Zn = (FeCl2)2 +ZnCl2. 
Ferrous Chloride—FeCl2—crystallizes from hydrochloric 
acid solutions in green monoclinic prisms, with four molecules 
of water. They deliquesce in the air, and oxidize. When 
dried they sustain a partial decomposition. The anhydrous 
salt is formed by conducting hydrogen chloride over heated 
iron. It is a white mass, which on application of heat fuses 
and at a red heat sublimes in white, six-sided leaflets. 
It forms double salts with the alkaline chlorides, e. g.:— 
2FeO+ 0 = Fe,03. 
Ferrous Iodide—Fel2—is obtained by warming iron with 
iodine and water. It also crystallizes with four molecules of 
water. 
FeCl2. 2KC1 + 2H20. 
Ferrous Oxide—FeO—is a black powder, resulting from the 
reduction of ferric oxide with carbon. When warmed it 
oxidizes readily. Ferrous Hydrate—Fe (OH)2—is thrown out 
of ferrous solutions, as a white precipitate, by the alkalies. 
Exposed to the air, it oxidizes, becoming green at first, then 
reddish-brown. It is somewhat soluble in water, and has an 
alkaline reaction. 
Ferrous Sulphate—FeS04—crystallizes with 7 molecules of 
H.20 in large, greenish, monoclinic prisms, and is generally 
called green vitriol. In dry air, these disintegrate some; in 
moist air they become coated, by oxidation, with a brown layer 
of basic ferric sulphate. At 100° they lose 6 molecules of H20, 
and become a white powder. The last molecule of water es- 
capes at 300°. Therefore, ferrous sulphate behaves just like the 
sulphates of the metals of the magnesium group. Like them, 
it yields double salts with potassium and ammonium sulphates, 
which have six molecules of H20. They are more stable than 
ferrous sulphate, and oxidize but very slowly in the air. 
Ferrous sulphate is obtained by dissolving iron in sulphuric 
acid; commercially it is obtained from pyrites (FeS2). On 392 
INORGANIC CHEMISTRY. 
roasting, these lose one molecule of sulphur, and are converted 
into ferrous sulphide (FeS), which, in presence of water, 
absorbs oxygen from the air, and forms sulphate, which may 
then be extracted by water. 
Iron vitriol has an extended practical application ; among 
other uses, it is employed in the preparation of ink, and in 
dyeing. 
When heated it decomposes according to the following 
equation:— 
2FeS04 = Fe203 + S03 + S02. 
On this is based the production of fuming Nordhausen 
sulphuric acid and of colcothar. 
Ferrous Carbonate.—FeC03—exists in nature as siderite, 
crystallized in yellow-colored rhombohedra, isomorphous with 
calcite and smithsonite. Sodium carbonate added to ferrous 
solutions precipitates a white voluminous carbonate, which, 
in the air, rapidly oxidizes to ferric hydrate. Ferrous carbo- 
nate is somewhat soluble in water containing carbon dioxide, 
hence present in many natural waters. 
Ferrous Phosphate—Fe3(P04)2 + 8ILO—occurs crystal- 
lized in bluish leaflets as Vivianite. Ferrous phosphate, 
precipitated by sodium phosphate, is a white amorphous 
powder, which oxidizes in the air. 
Ferrous Sulphide—FeS—is obtained by fusing together 
iron and sulphur as a dark gray, metallic mass. It is made 
use of in laboratories for the preparation of hydrogen sulphide. 
If an intimate mixture of iron filings and sulphur be mois- 
tened with water, the union will occur even at ordinary tem- 
peratures. Black ferrous sulphide is precipitated from ferrous 
solutions by alkaline sulphides. In moist condition in the air 
it oxidizes to ferrous sulphate. The same reagents also 
precipitate ferrous sulphide from ferric salts, but they first 
suffer reduction :— 
Fe2Cl6 + 3(NH4)2S = 2FeS + 6NH4C1 + S. 
FERRIC COMPOUNDS. 
Ferric Oxide—Sesquioxide of Iron—Fe203—exists as hema- 
tite in nature. It may be formed by heating the iron oxygen 
compounds in the air, and on a large scale is obtained by the 
ignition of green vitriol. It then is a dark red powder IRON, 
393 
(colcothar or caput mortuum) used as a paint and for polish- 
ing glass. 
Ferric Hydrate — Fe.2(OH)6. As a voluminous, reddish- 
brown mass, it is precipitated by alkalies from ferric solutions. 
On boiling with water, it becomes more compact, gives up 
water and is converted into the hydrate, Fe20(0H)4. Many 
iron ores, like bog-iron ore, Fe20(0H)4, pyrosiderite, Fe202 
(OH)2, and brown hematite, Fe403(0H)6, are similarly 
derived. 
Freshly precipitated ferric hydrate is soluble in a solution 
of ferric chloride or acetate. When such a solution is sub- 
jected to dialysis, the iron salt diffuses, and there remains a 
pure, aqueous solution of ferric hydrate. From this, by slight 
quantities of alkalies and acids, all the ferric hydrate is pre- 
cipitated as a jelly. 
• Ferrous-Ferric Oxide—Fe304=Fe0.Fe203—occurs crys- 
tallized in regular octahedra, in nature, as magnetite. It is 
abundant in Sweden, Norway, and in the Urals. It may be 
artificially obtained by conducting steam over ignited iron. 
Magnetite constitutes the natural magnets. 
Ferric hydrate, like other sesquioxides, is only a weak base, 
which does not yield salts with weak acids, like carbonic or 
sulphurous (p. 347). 
Ferric salts arise by the solution of ferric oxide in acids, or 
by the oxidation of ferrous salts in the presence of free 
acids:— 
2FeS04 + H2S04 + 0 = Fe2(S04)8 + H20. 
They generally have a yellow-brown color. By reduction, 
they pass into ferrous salts :— 
Fe2Cl6 + H2S = 2FeCl2 + 2HC1 + S. 
Ferric Chloride—Fe2Cl6. In aqueous solution it is ob- 
tained by conducting chlorine into a solution of ferrous chlo- 
ride :— 
2FeCl2 + Cl2 = Fe2Cl6. 
Upon evaporation, the hydrate—Fe2Cl6 + 6H20—remains 
as a yellow crystalline mass, readily soluble in water, alcohol 
and ether. It is partially decomposed when heated ; hydro- 
gen chloride escapes and a mixture of chloride and oxide 
remains. 394 
INORGANIC CHEMISTRY. 
Anhydrous ferric chloride is made by heating iron in a 
current of chlorine gas ; it sublimes in brownish-green, metal- 
lic, shining, six-sided prisms and scales, which deliquesce in 
the air. The specific gravity of their vapor is 162.5 ( H = 1) 
corresponding to the molecular formula Fe2Cl6 = 325. 
Ferric Sulphate—Fe2(S04)3—is obtained by dissolving the 
oxide in sulphuric acid, and remains, on evaporating the 
solution, as a white mass, which dissolves in water gradually, 
with a reddish-brown color. With alkaline sulphates it forms 
alums (p. 341) e. g.:— 
Fe2(S04)3, K2S04 + 24H20. 
Potassium Iron Alum. 
Ferric Phosphate—Fe2(P04)2—is a white precipitate, thrown 
out of ferric solutions by sodium phosphate. It is insoluble 
in water and acetic acid. 
Ferric Sulphide—FeS2—occurs in nature as pyrites, crystal- 
lized in yellow, metallic, shining, regular cubes. It is em- 
ployed in the manufacture of sulphuric acid and green vitriol. 
It may be artificially prepared in various ways. 
COMPOUNDS OF FERRIC ACID. 
On fusing iron filings with nitre, or by conducting chlorine 
into potassium hydrate, in which there is suspended ferric hy- 
drate, Potassium Ferrate K2Fe04 is produced, which crystallizes 
from the alkaline solution in dark red prisms; they are iso- 
morphous with those of the chromate and sulphate. They are 
readily soluble in water; the dark red liquid soon decom- 
posing with separation of ferric hydrate and oxygen. The free 
acid is not known, as it at once decomposes when liberated 
from its salts. 
CYANOGEN DERI Y ATI YE S OE IRON. 
With the cyanogen group iron forms very characteristic, 
and for commerce, important compounds. When potassium 
cyanide is added to aqueous solutions of the ferrous or ferric 
salts, the cyanides Fe(CN)2 and Fe2(CN)6 are thrown down as 
white precipitates, which in the air soon break up. In excess 
of potassium cyanide they dissolve to the double cyanides 
Fe(CN)2.4KCN and FeJ(CN)6.6KCN. Acids added to these IRON 
395 
solutions induce the separation of the hydrogen compounds, 
H4FeCy6 = FeCy2.4HCy and Fe2Cy,2H6= Fe2Cy6.6HCy.* 
These are of acid nature, and form salts by exchanging 
their hydrogen for metals. In these salts, and also in the 
free acids, neither the iron nor the cyanogen group can be 
detected by the usual reagents (iron is not precipitated, e. g., 
by the alkalies). It is supposed that in these double cyanides 
compound groups of peculiar structure are present, and that 
they conduct themselves analogously with the halogens. The 
group FeCy6 in the ous compounds is called ferrocyanogen, that 
of Fe2Cyi2 in the ic, ferricyanogen. The ferro conduct them- 
selves towards the ferri compounds like the ferrous to the 
ferric salts ; by oxidation the first are converted into the latter, 
and by reducing agents the latter into the former. 
2FeCy6K4 -4- Cl2 — K6Fe2Cyi2 “l- 2KC1 and 
Fe2Cy12K6 + 2K0H -(- H2 = 2K4FeCy6 -(- 2H20. 
Cobalt, manganese, chromium and the platinum metals 
form similar cyanides. 
Potassium Ferrocyanide—Yellow Prussiate of Potash—K4 
FeCy6, is produced by the action of potassium cyanide upon 
iron compounds, or upon free iron (in this case the oxygen of 
the air or of water takes part). Commercially it is prepared 
by igniting carbonized nitrogenous animal matter (blood, 
horn, hoofs, leather, offal, etc.,) with potashes and iron. 
In this operation first the carbon and nitrogen of the organic 
matter combine 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 cyan- 
ide can be extracted from the fusion). Upon now treating the 
fusion with water, there arises, through the alternating action 
of the potassium cyanide and sulphide, ferrocyanide of 
potassium, which is purified by recrystallization:— 
It crystallizes from water in large, yellow, quadratic prisms, 
having three molecules of water and soluble in 3-4 parts 
H20. At 100° the crystals lose all their water and become 
a white powder. At a red heat the ferrocyanide breaks up 
into cyanide, nitrogen and iron carbide (FeC2). When the 
salt is warmed with dilute sulphuric acid, half of the cyanogen 
FeS + 6KCy = K4FeCy6 + K2S. 
* The cyanogen group, CN, is usually designated by the letiers Cy. 396 
INORGANIC CHEMISTRY. 
escapes as hydrogen cyanide; by concentrated sulphuric acid 
it is decomposed, according to the following equation: 
K4FeCy6 + 6H2S04 + 6H20 = FeS04+ 2K2S04 +3S04(NH4)a 
+ 6C0. 
When strong hydrochloric acid is added to a concentrated 
potassium ferrocyanide solution, there separates hydrogen fer- 
rocyanide, H4FeCy6, as a white crystalline powder, which 
rapidly turns blue in the air. It has the nature of an acid. 
Its salts with the alkalies and alkaline earths are very soluble 
in water. The sodium salt crystallizes with difficulty. The 
salts of the heavy metals are insoluble, and are obtained by 
double transposition. Ferrocyanide added to the solution of 
a ferric salt precipitates a dark blue cyanide (FeCy6)3, (Fe2)2 
called Prussian Blue:— 
3K4FeCy6 + 2Fe2Cl6 = (FeCy6)3(Fe2)2 + 12KC1. 
This is the ferric salt of hydroferrocyanic acid ; on pouring 
potassium or sodium hydrate over it ferrocyanide of potassium 
and ferric hydrate are produced :— 
In copper solutions potassium ferrocyanide produces a 
reddish-brown precipitate of FeCy6Cu2. 
(FeCy6)3(Fe2)2 + 12K0H = 3K4FeCy6 + 2Fe2(OH)6. 
Oxidizing agents convert the ferro into potassium ferricya- 
nide—K6Fe2Cy12—red prussiate of potash. This conversion is 
most conveniently effected by conducting chlorine into the 
solution of the yellow prussiate :— 
2K4FeCy6 + Cl2 = K6Fe2Cy12 + 2KC1. 
By this the ferrocyanogen group FeCy6 is changed to the 
ferri, Fe2Cy12. 
The red prussiate crystallizes from water in red rhombic 
prisms. Concentrated hydrochloric acid precipitates the free 
hydro-ferricyanic acid H6Fe2Cy12, which decomposes very 
easily. 
Potassium ferri cyanide with ferrous solutions affords a 
dark blue precipitate, Fe2Cy12Fe3, very similar to Prussian 
Blue, and called Turnbull's Blue:— 
K6Fe2Cy12 3FeS04 —Fe2Cy12Fe3 -f- 3K2S04. COBALT. 
397 
This blue is the ferrous salt of hydroferricyanic acid. 
Alkalies change it to ferricyanide of potassium and ferrous 
hydrate:— 
Fe2Cy12Fes + 6K0H = Fe2Cy12K6 + 8Fe(0H)a * 
In ferric solutions, potassium ferricyanide does not cause pre- 
cipitation. Ferrocyanide gives Prussian blue, while in ferrous 
solutions it forms a bluish-white precipitate. 
By these reactions, ferric salts may readily be distinguished 
from the ferrous. Potassium sulphocyanide (CNSIv) pro- 
duces a dark red coloration in ferric solutions, while it leaves 
the ferrous unaltered. 
2. COBALT. 
Co — 58.6. 
Occurs in nature as smaltite (CoAsj) and cobaltite (CoAs2) 
CoS2). The metal is obtained by the ignition of cobaltous 
oxide with carbon or in an hydrogen stream. It has a red- 
dish-white color and strong lustre, is very tenacious and diffi- 
cultly fusible. Its specific gravity is 8.9. It is attracted by 
magnets, but to a less degree than iron. It is not altered by 
the air or water. Only slightly attacked by hydrochloric and 
sulphuric acids ; nitric acid dissolves it readily, forming cobalt 
nitrate. 
The predominating compounds have the form CoX2, and 
are called cobaltous. They are very stable and mostly isomor- 
phous with the ferrous salts. The hydrous cobaltous com- 
pounds have a reddish color, the anhydrous are blue. 
COBALTOUS COMPOUNDS. 
Cobaltous Chloride—CoCl2—is obtained by the solution of 
cobaltous oxide in hydrochloric acid, and crystallizes with 
6H20 in red monoclinic prisms. When heated, it loses water, 
and becomes blue in color, and anhydrous. On writing on 
paper with the solution, the pale reddish characters are almost 
invisible, but wrhen warmed they appear distinctly blue 
(Sympathetic ink). 
* According to recent investigations it appears that Turnbull’s blue 
VI 
and Prussian blue possess the same composition (FeCy6)2 j j^'2 The 
simpler relations are retained here. 398 
INORGANIC CHEMISTRY. 
Cobaltous Hydrate—Co(OH)2—a reddish precipitate pro- 
duced by the alkalies in hot, cobaltous solutions. Exposed 
to the air, it becomes brown by oxidation. From cold solutions 
basic salts are precipitated. Heated out of air contact the 
hydrate passes into green cobaltous oxide CoO. 
Cobaltous Sulphate—S04C0-f- 7H20—crystallizes in mono- 
clinic prisms; from a warm solution there separates CoS04 
-f- 0H20. It is isomorphous with ferrous sulphate, and forms 
double salts with alkaline sulphates. 
Cobaltous Nitrate—Co(N03)2 -f 6H20—forms red delique- 
scent prisms. 
Cobaltous Sulphide—CoS—is a black precipitate, produced 
in neutral cobalt solutions by alkaline sulphides, and insoluble 
in dilute acids. 
Cobalt Silicates. When glass is fused with a cobalt com- 
pound it is colored a dark blue and when pulverized becomes 
a blue powder which is used as a paint, under the name of 
smalt. 
Smalt is commercially prepared by fusing cobalt ores with potashes 
and quartz, whereby the cobalt forms a silicate (smalt) with the Si02 
and potassium, while the other metals accompanying Co in its ores, 
such as Bi, As, and especially nickel, are thrown out as a speiss. This 
is called speiss-cobalt and serves for the preparation of nickel. 
On igniting cobalt oxide with alumina a dark blue mass is 
produced—cobalt ultramarine or Thenard’s Blue. When zinc 
oxide and cobalt oxide are ignited a green color—green 
cinnabar—is obtained. 
COBALTIC COMPOUNDS. 
Cobaltic Oxide—Co203—is left as a black powder on the 
ignition of cobaltous nitrate. At a red heat it becomes 
eobaltous-cobaltic oxide, Co304, and at a white heat cobaltous 
oxide. The hydrate—Co2(OH)6—separates as a dark brown 
powder, if chlorine be passed through an alkaline solution 
having a cobaltous salt in it. 
By the action of sulphuric acid upon the oxide and its 
hydrate there is formed a cobaltous salt and hydrogen set free. COBALT 
399 
Chlorine is generated when it is heated with hydrochloric 
acid :— 
Co2Os + 6HCl=2CoCl2+8HaO + Cl2. 
In dilute, cold hydrochloric acid the cobaltic hydrate dis- 
solves almost without any liberation of chlorine ; in solution, 
is probably Co2Cl6, which in evaporation is decomposed into 
2CoCl2and Cl2. 
Cobaltous-Cobaltic Oxide—Co304 = Co203CoO—corres- 
ponding to magnetite, Fe304—is formed by the ignition of the 
oxygen cobalt derivatives, and is a black powder. 
Of the oxygen salts of cobalt only a few double salts are 
known, of which potassio-cobaltic nitrite is worthy of note. 
When potassium nitrite KN02 is added to a cobaltous solu- 
tion acidified with acetic acid, nitrogen is set free, and in 
course of time there separates the double salt 
Co2(N02)6. 6KN02 + nH20, 
as a yellow crystalline powder. This reaction is very char- 
acteristic, for cobalt, and serves to separate it from nickel. 
Ammonio-Oobalt Compounds. Cobalt is capable of form- 
ing with ammonia a series of peculiar compounds, in which it appears 
in its highest state of oxidation; the structure of them has not yet 
been explained. On adding ammonium hydrate to a cobaltous chlo- 
ride solution, the first formed precipitate dissolves in the excess of the 
reagent, and when this liquid is permitted to stand pxposed to the air, 
the color, at first brown, passes gradually into red. On adding con- 
centrated hydrochloric acid to this, a brick-red, crystalline powder, of 
the composition—Co2Cl6 10NH3 + ‘2H20—called Roseocobalt—is 
precipitated. If, however, the red solution be boiled with hydro- 
chloric acid, a red powder—purpureocobaltic chloride, Co2Cl6 -f- 
10NH3—separates out. If the ammoniacal red solution contain much 
ammonium chloride, hydrochloric acid will precipitate a yellowish- 
brown compound—luteo-cobaltic chloride Co2C16.12NH3. 
The other salts of cobalt, such as the sulphate and nitrate, yield simi- 
lar compounds, e. g., Co2(NO3)6.10NH3, roseocobaltic nitrate. 
Cyanogen Cobalt Compounds.—In solutions of cobaltous 
salts, potassium cyanide produces a bright brown precipitate of cobalto- 
cyanide. Co(CN)2 soluble in an excess of the reagent. In the air 
the solution absrt-bs oxygen and forms potassium cobalticyanide, 
K6Co2(CN)12 corresponding to potassium ferricyanide. On evapora- 
tion of the solution the cobalticyanide crystallizes in colorless rhombic 
prisms, very soluble in water. Sulphuric acid precipitates from the 
concentrated solution, hydrogen cobalticyanide, H6Co2(CN)12, 
crystallizing in needles. 400 
INORGANIC CHEMISTRY. 
3. NICKEL. 
In native condition nickel exists in meteorites; its most 
important ores are Niccolite—NiAs—and Gersdorffite, NiS2. 
NiAs2 (constituted like cobaltite.) In its ores nickel 
is always accompanied by cobalt, and vice versd, cobalt 
mostly by nickel. The isolation of the latter from its ores, as 
also from speiss-cobalt (p. 398), is very complicated. Gener- 
ally nickel appears in commerce in cubical forms; in 
addition to the chief ingredient, there always are present 
copper, bismuth and other metals. To get chemically pure 
nickel, its oxalate or carbonate should be ignited in an 
hydrogen stream. 
Nickel is almost silver white, very lustrous, and very tena- 
cious. Its specific gravity is 9.1 and that of the fused variety 
8.8. It fuses at a somewhat lower temperature than iron, and 
like this is attracted by the magnet. In the air it is not 
altered ; it dissolves with difficulty in hydrochloric and sul- 
phuric acids, but readily in nitric acid. 
It forms almost exclusively ows compounds of the form 
NiX2; nickelic oxide behaves like a peroxide, and does not 
afford corresponding salts. 
Ni = 58.6. 
Nickelous Hydrate—Ni(OH)2—is a bright green precipitate 
produced by alkalies in nickelous solutions. It dissolves in 
ammonium hydrate, with a blue color. When heated it 
passes into gray nickelous oxide, NiO. 
Nickelous Chloride—NiCl2 + 6H20—forms green, mono- 
clinic prisms. When heated they lose water and become yel- 
low. 
Nickelous Cyanide—Ni(CN)2—is precipitated, as a green- 
colored mass, from nickel solutions, by potassium cyanide. It 
is soluble in excess of the precipitant. From the solution 
there crystallizes the double cyanide, NiCy22KCy -f- H20. 
Acids decompose this readily. Cyanogen compounds, consti- 
tuted like those of iron and cobalt, are not with nickel. 
Nickelous Sulphate—NiS04 -f- 7H20—appears in green, 
rhombic prisms, isomorphous with the sulphates of the mag- 
nesium group, and forms analogous double salts. 
Nickelous Sulphide—NiS—is precipitated, black-colored, 
by alkaline sulphides from nickel solutions. PLATINUM. 
401 
Nickelic Oxide—Ni203—and Hydrate — Ni2(OH)6—are 
perfectly similar to the corresponding cobalt salts; when 
warmed with hydrochloric acid they give up chlorine. 
Nickel is used for certain alloys. Argentan consists, ordin- 
arily, of 50% copper, 25% nickel and 25% zinc. The more 
nickel there is present in the alloy the whiter and harder it is, 
and the more capable of receiving a high polish. The Ger- 
man nickel coins consist of 75% Cu and 25% Ni. At pres- 
ent, to protect cast iron ware from rusting and to impart to it 
a beautiful white surface, it is coated with a layer of nickel. 
This is accomplished in an electrolytic manner, or by boiling 
the iron ware in a solution of zinc chloride and nickel sulphate. 
GROUP OF THE PLATINUM METALS. 
Here belong palladium, rhodium, ruthenium, osmium and 
iridium—the constant companions of platinum in its ores. 
On page 385 we observed that these metals are divided into 
two groups ; the group of light platinum metals and the group 
of heavy ones, which have higher atomic weights and specific 
gravities:— 
Ru, 103.5 sp. gr. 12.26 Rh, 104 sp. gr. 12.1 Rd, 10.6 sp. gr. 11.8 
Os, 198* “ “ 22.4 Ir, 192.7 “ 22.38 Pt, 196.7 “ “ 21,4 
The relations of the metals of this group to each other, like 
those of the iron group, are entirely similar, and they show 
in their physical and chemical properties a great similarity 
to the corresponding members of the iron group. Like iron, 
osmium and ruthenium have a gray color, are very difficultly 
fusible and readily oxidized in the air. Palladium and 
platinum, however, like nickel, have an almost silver-white color, 
are more fusible, and are not oxidized by oxygen. In chemical 
relations osmium and ruthenium, like iron, show at the same 
time a metalloidal nature, inasmuch as their highest oxygen 
compounds form acids. Their derivatives exhibit great simi- 
larity to those of iron :— 
II 
OsO 
Osrnous 
oxide. 
IV 
Os203 
Osmic 
oxide. 
IV 
0s02 
Osmium 
dioxide. 
(Os03) 
Osmic 
trioxide. 
RuO 
Ruthenous 
oxide. 
2 0 3 
Ruthenic 
oxide. 
Ru02 
Ruthenium 
dioxide. 
(Ru03) 
Ruthenium 
trioxide. 
*See note (page 385). 402 
INORGANIC CHEMISTRY. 
The acid oxides 0s03 and Ru03 are unknown, but the cor- 
responding acids H20s04 (osmic acid) and H2Ru04 (ruthenic 
acid) and derived salts, are known. In addition, osmium and 
ruthenium exhibit another very high degree of oxidation— 
0s04, perosmic oxide, and Ru04, per-ruthenic oxide—which 
are not known with iron ; in these derivatives the metals 
appear to be octads, yet these oxides form no corresponding 
acids or salts. 
Rhodium and iridium, like cobalt, do not form acidic-like 
derivatives. Their salts correspond to the forms :— 
RhO 
Rhodous 
oxide. 
IV 
R1i203 
Rhodic 
oxide. 
IV 
Rh02 
Rhodium 
dioxide. 
The rhodic compounds are the most stable. 
Palladium and platinum, finally, are relatively of more 
basic nature, as their ous derivatives PdX, and PtX2 are 
proportionally more stable than the ic forms PdX4 and PtX4. 
Palladium also forms a lower oxide, palladium suboxide Pd20 
in which it approaches silver. 
The platinum metals are almost exclusively found in nature 
in the so-called platinum ore, which usually occurs in small 
metallic grains in accumulated sands of a few regions (in 
California, Australia, the Island of Sumatra and chiefly 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 contains usually 50-80 °Jo platinum, besides palladium 
(to 2 %), iridium (7 %), osmium (II c/0), and ruthenium 
(II %), and also different other metals, like gold, copper, iron. 
The separation of the platinum metals is generally executed 
in the following manner: The gold is first removed by dilute 
aqua regia. Then the ore is treated with concentrated aqua 
regia, whereby platinum, palladium, rhodium, ruthenium and 
a portion of iridium are dissolved. Metallic grains or leaflets, 
an alloy of osmium and iridium—platinum residues—remain. 
Then ammonium chloride is added to the solution, and plati- 
num and iridium are precipitated as double salts. When the 
precipitate is ignited a spongy mass of iridium-bearing plati- 
num (platinum sponge) is obtained, which is directly applied 
in the manufacture of platinum vessels. The filtered RUTHENIUM AND OSMIUM. 
403 
solution from the insoluble chlorides contains palladium, 
rhodium and ruthenium, which are thrown down as a metallic 
powder by iron; their further separation is then effected in 
various ways. 
For the manufacture of platinum objects spongy platinum 
was formerly solely employed ; this was pressed into moulds, 
then glowed and hammered out. Now the fusibility of Pt in 
the oxy-hydrogen flame is made use of, and the molten metal 
run into moulds. 
By means of the oxy-hydrogen blowpipe iridium and rho- 
dium-bearing platinum may be directly fused out of the 
platinum ore; osmium and ruthenium, for the most part, are 
consumed in this operation. The presence of iridium and 
rhodium makes platinum harder and less readily attacked by 
many reagents. 
RUTHENIUM and OSMIUM 
Ruthenium has a steel gray color; it is very hard, brittle, and diffi- 
cultly fusible. In pulverized condition it oxidizes, when ignited in the 
air, to RuO and Ru203. It is insoluble in acids, and only slowly dis- 
solved by aqua regia. Fused with potassium hydrate and nitrate, it 
forms potassium rutheniate K2Ru04. 
Ruthenium heated in chlorine gas yields ruthenium dichloride RuC12, 
a black powder, insoluble in acids. The sesquichloride Ru2Cl6 is 
obtained by the solution of Ru2(OH)6, and is a yellow, crystalline 
mass, which deliquesces in the air. With potassium and ammonium 
chlorides, it yields crystalline double chlorides, e. g., Ru2C16.4KC1. 
The tetrachloride RuC14 is only known in double salts. Ruthenious 
oxide RuO, the sesquioxide Ru203, and dioxide Ru02, are black 
powders, insoluble in acids, and are formed when ruthenium is roasted 
in the air. 
Ru = 103.5. 
Os —198.* 
The hydrates Ru2(OH)6 and RufOH)4 are produced by the action 
of the alkalies upon the corresponding chlorides, and are very readily 
soluble in acids. Ruthenic acid H2Ru04 is not known in free con- 
dition. Its potassium salt K2Ru04 is formed byfusing-the metal with 
potassium hydrate and nitre. It dissolves in water with an orange 
yellow color. When chlorine is conducted through the solution ru- 
thenium tetroxide Ru04 separates as a gold-yellow crystalline mass. 
It fuses at 40° and boils about 100°, and yields a yellow vapor, the odor 
of which is similar to nitrogen dioxide, N02. At 108° it decomposes 
with explosion Water breaks it up with formation of Ru2(OH)6. In 
concentrated potassium hydrate it dissolves to Ru04K2 By less con- 
* Compare note (p. 385). 404 
INORGANIC CHEMISTRY. 
tinuous introduction of Cl into the solution of Ru04K2 greenish 
black crystals separate out, which are isomorphous with potassium 
permanganate, and appear to be Ru04K. 
Osmium is very much like the preceding. It is not even fusible in 
the oxy-hydrogen flame ; it only sinters together somewhat. As a 
fine powder it burns when glowed in the air to 0s04. 4 Nitric acid and 
aqua regia oxidize it to the same. The compounds OsCl2, OsO, 
Os2Cl6 and 0s203,0s02 and 0sCl4 are very similar to the cori'es- 
ponding compounds of ruthenium. By fusion with potassium hydrate 
and nitre we get potassium osmate—K20s04—which crystallizes from 
aqueous solution with 2H20 in dark-violet octahedra. The most 
stable and very characteristic derivative of osmium is the tetroxide, 
0s04, which is produced by glowing the metal in the air or by the 
action of chlorine on osmium in presence of water. It crystallizes in 
large colorless prisms, which fuse below 100° and distill at somewhat 
higher temperature. It has a very sharp, piercing odor, similar to that 
of sulphur chloride. Reducing and organic substances precipitate 
from it pulverulent osmium. On this depends its application' in 
microscopy. 0s04 and Ru04 do not afford corresponding salts. 
RHODIUM and IRIDIUM. 
These have more of a white color and are more easily fusible than 
ruthenium and osmium. (Iridium fuses at 1950°.) In pure condition 
they are not attacked by acids or aqua regia ; alloyed with platinum 
they, however, dissolve in aqua regia. 
Rhodium forms three oxides: RhO, Rh203 and Rh02, of which 
the second forms salts with acids. When rhodium is heated with 
nitre RhO, results. 
Rh = 104. 
Ir — 192.7. 
Of the chlorides only Rh2Cl6 is known. It results when the metal 
is heated in chlorine gas. It is a brownish-red mass. With alkaline 
chlorides it forms readily crystallizing, red-colored double salts. 
Iridium has perfectly analogous derivatives : IrO, Ir2Os, Ir02 and 
IrCl2, Ir2Cl6, IrCl4. The sesquichloride, Ir2Cl6, formed by heating 
Ir in chlorine, is an olive-green, crystalline mass, insoluble in water 
and acids. It forms double salts with the alkaline chlorides, e. g., 
Ir2Cl6, 6KC1 + 6H20, which crystallize from water in green crystals. 
They are also produced by the action of S02 upon the double salts of 
IrCl4. 
Iridium Tetrachloride—IrCl4—is formed by 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). When alkaline 
chlorides are added to the solution double chlorides are precipitated, 
e. <7., IrCl4. 2NH4C1, isomorphous with the double chlorides of plati- 
num. When a solution of IrCl4 is boiled with KOH, Ir(0H)4 will 
be precipitated. PALLADIUM. 
405 
PALLADIUM. 
Pd - 106.2 
Palladium, in addition to occurring in platinum ores, is 
alloyed with gold (Brazil), and in some selenium ores(Hurtz); 
it has a silver white color, and is somewhat more fusible than 
platinum. In finely divided condition it dissolves, in boiling 
concentrated HC1, sulphuric and nitric acid. When ignited 
in the air, it, at first, through oxidation, becomes dull; at 
higher temperature, however, the surface again assumes a 
metallic appearance. 
Palladium possesses, like some other metals (as platinum 
and silver), but in much higher degree, the ability of absorb- 
ing hydrogen gas (occlusion). 
Freshly ignited palladium leaf absorbs at ordinary tempera- 
tures upwards of 370, at 90-100°C., about 650 vols. of hydro- 
gen. A greater absorption, at ordinary temperatures, may be 
easily accomplished in the following manner:— 
Water is decomposed by the galvanic current, palladium 
foil being used as negative electrode. The liberated hydrogen 
is then taken up by the palladium (to 960 vols); the metal 
expands (tV its volume), becomes specifically lighter, but 
entirely retains, however, its metallic appearance. According 
to the investigations of Debray, the compound Pd2H is pro- 
duced, in which the hydrogen is contained dissolved, and 
deports itself similar to an alloy (compare p. 39). Palladium 
charged with hydrogen usually remains unaltered in the air 
and in a vacuum; sometimes, however, it becomes heated in the 
air, as the hydrogen is oxidized to v'ater. The same occurs 
upon heating palladium hydride to 100°; in vacuo, all the 
hydrogen escapes as gas. Palladium hydride is a strong 
reducing agent, similar to nascent hydrogen. 
Ferric salts are reduced to the ferrous state ; chlorine and 
iodine in aqueous solution are converted into potassium chlo- 
ride and hydriodic acid. 
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. On heating palladium sponge in the air 
until the white metallic color becomes black, through the 
superficial formation of a layer of palladious oxide, it will at 
ordinary temperatures absorb hydrogen very energetically, 
with partial oxidation to water. 406 
INORGANIC CHEMISTRY. 
When palladium sheet or sponge is introduced into the 
flame of a spirit lamp, it is covered with smoke; this is due 
to the fact that the metal withdraws the hydrogen of the 
hydrocarbons, and carbon is set free. 
Palladium forms two series of compounds; the palladious 
PdX2 and palladic, PdX4. The first are well characterized 
and are distinguished by their stability. 
Palladious Chloride—PdCl2—remains as a brown, delique- 
scent mass, in evaporating the solution of palladium in aqua 
regia. With alkaline chlorides it yields easily soluble, crystal- 
line double salts, e. 2 2KC1. 
Palladious Iodide—Pdl2—is precipitated as a black mass, 
insoluble iu water, by potassium iodide, from palladium solu- 
tion. 
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 dissolves in sulphuric or nitric acids the 
corresponding salts are produced. 
The sulphate, PdS04 -j- 2H20,is composed of brown crystals, 
readily soluble in water. Much of the latter decomposes it. 
Palladio Chloride—PdCl4—is formed when the metal is 
dissolved in aqua regia. On evaporation it decomposes into 
PdCl2 and Cl2. On adding potassium or ammonium chloride 
to the solution, difficultly soluble red double chlorides, similar 
to the corresponding platinum salts, appear. 
PLATINUM. 
Pt = 196.7. 
Its separation from the ore was described on p. 402. The 
metal has a grayish-white color and a specific gravity of 
21.5. It is very tough and malleable, and may be drawn out 
into very thin wire and rolled into foil. At a strong heat it 
becomes soft without melting; in the oxy-hydrogen flame it 
fuses (about 1770°—Violle) and is somewhat volatile. On 
fusion it absorbs oxygen, which it again gives up on cooling 
(like silver). At ordinary temperatures it condenses oxygen 
too, upon its surface, especially when in finely divided state, as 
platinum black or sponge. The first is obtained if reducing PLATINUM. 
407 
substances, like zinc, be added to solutions of platinic chlo- 
ride or upon boiling with sugar and sodium carbonate; it 
absorbs as much as 250 volumes of oxygen. Platinum black 
remains as a spongy mass upon the ignition of PtCl4.2NH4Cl. 
Upon this property of platinum to condense oxygen, as we 
have seen, depends the introduction of various reactions : thus 
hydrogen inflames in the air, if it be conducted upon platinum 
sponge ; sulphur dioxide combines at 100° with O to form the 
trioxide. At a red heat platinum permits free passage to 
hydrogen, while it is not permeable by oxygen and other 
gases. 
Platinum is not attacked by acids; it is only soluble in 
liquids generating free chlorine, as aqua regia. In conse- 
quence of this opposition to acids, and also its unalterability 
upon ignition, this metal answers as an undecomposable 
material for the production of chemical crucibles, dishes, 
wire, etc. The usual presence of iridium in ordinary platinum 
increases its durability. 
The alkaline hydrates, sulphides and cyanides attack it 
strongly at a red heat. With phosphorus, arsenic, and many 
heavy metals, especially lead, it gives rise to easily fusible 
alloys; many heavy metals are also reduced from their salts 
by platinum. Therefore such substances must not be ignited 
in platinum crucibles, etc. 
Platinum, like palladium, forms platinous PtX2 and 
platinic PtX4 compounds; in the first it is more basic, in the 
latter more of an acid nature. 
Platinic Chloride—PtCl4—is obtained by the solution of 
platinum in aqua regia, and, on evaporation of the solution 
remains as a red-brown crystalline mass, very deliquescent in 
the air. It forms characteristic double chlorides, PtCl4.2KCl, 
with ammonium and potassium chloride. These are difficultly 
soluble in water ; hence, on mixing the solutions, they immedi- 
ately separate out in form of a crystalline yellow powder. 
Ignition completely decomposes the ammonium salt, leaving 
spongy platinum. 
Similar insoluble double chlorides are formed by platinum 
chloride with those of rubidium, csesium and thallium, while 
that with sodium is very soluble in water. At 200°, PtCl4 
breaks up into PtCl2 and Cl2. 
On adding NaOH to platinic chloride and then supersatu- 
rating with acetic acid, there separates a reddish-brown pre- 408 
INORGANIC CHEMISTRY. 
cipitate of platinic hydrate, Pt(OH)4. This dissolves readily 
in acids (excepting acetic), with formation of salts. The 
oxygen salts, as Pt(S04)2, are very unstable. The hydrate 
also has an acidic character (platinic acid), and dissolves in 
alkalies, forming salts with them. These, also, result in 
fusing platinum with potassium and sodium hydrate. The 
barium salt, + 3H20, is precipitated from platinic 
chloride, by barium hydrate, as a yellow, crystalline com- 
pound. As regards the acidic nature of its hydrate, platinum 
approaches gold. If hydrogen sulphide be conducted through 
platinic solutions, black platinum disulphide, PtS2, is precipi- 
tated, soluble in alkaline sulphides, with formation of sul- 
pho-salts. 
Platinous Chloride—PtCl2—is a green powder, insoluble in 
water, remaining after heating PtCl4 to 200°. It also affords 
double salts with alkaline chlorides, e. g., PtCl2. 2NaCl. 
Warmed with potassium hydrate it gives the hydrate Pt(OH)2. 
CYANOGEN COMPOUNDS. 
Like iron and cobalt, platinum yields double cyanides cor- 
responding to ferrocyanides. When platinous chloride is dis- 
solved in potassium cyanide there crystallizes on evaporation 
platinum-potassium cyanide, K2PtCy4 -f- 3H20, in large prisms 
exhibiting magnificent dichroism ; in transmitted light they are 
yellow ; in reflected light blue. This salt must be viewed as 
the potassium compound of hydro-platino-cyanic acid H2PtCy4. 
Separated from its salts it crystallizes in gold-yellow needles. 
Its salts with the heavy metals are obtained by double replace- 
ment, and all show a beautiful play of colors. 
PLATINUM-AMMONIUM COMPOUNDS. 
There is a whole series of these, which must be viewed as 
platinum bases and their salts. They are constituted accord- 
ing to the following empirical formulas :— 
Pt(NH3)2X2, Pt(NH3)2X4, Pt(NH3)4X2,Pt(NH3)4X4, 
in which X indicates various acid residues, or halogen atoms. 
They arise by the action of ammonium hydrate upon plat- 
inous chloride. The bases are obtained by substituting hy- SPECTRUM ANALYSIS. 
409 
droxyl groups for the acid residues, e. g., Pt(NH;i%(OH)2. 
In their chemical properties these resemble alkaline hydrates. 
The other platinum metals form similar amine derivatives. 
The nature and chemical constitution of these interesting 
compounds is, however, not yet entirely explained. 
SPECTRUM ANALYSIS.* 
We observed that various substances, if introduced into a 
non-luminous flame, imparted to it a characteristic coloration. 
Thus the sodium compounds color it yellow, the potassium, 
violet, thallium, green, etc., etc. Upon the decomposition of 
the thus formed rays of light, and, indeed, of every light, by 
means of the prism, and the consideration of the resulting 
spectrum, depends spectrum analysis, established in 1859 by 
Kirchhoff and Bunsen, which, by its important applications 
and universal use, constitutes one of the greatest of scientific 
achievements of all ages. 
As we well know, every substance, solid or liquid, heated to 
white heat emits rays (e. g., molten platinum; lime heated 
in the oxy-hydrogen flame ; the ordinary flame containing 
glowing particles of carbon,) of every refrangibility ; and 
hence, furnishes, if the light be conducted through a prism, 
a continuous spectrum, which brings to view all the colors of 
the rainbow, from red to violet. Glowing gases and vapors, 
on the contrary, whose molecules can execute unobstructed 
oscillations, emit light of definite refrangibility, and there- 
fore afford spectra, consisting of single, bright lines. Thus the 
spectrum of the yellow sodium flame is recognized as com- 
posed of one very bright yellow line, which by increased mag- 
nifying power is shown to consist of two lines lying very near 
each other. This reaction is so very delicate, that by means 
of it Tis-tshrtru of a milligram of sodium may be detected. 
The violet potassium light affi>rds a spectrum, consisting of a 
red and a blue line. The crimson strontium light shows 
in the spectrum several single red and a blue line. (See the 
spectrum plate.) Each of these lines corresponds to a very 
definite coefficient of refraction, therefore in the spectrum it 
occupies a very definite relative position. 
* A more exhaustive, concise and distinct presentation of the spec- 
trum phenomena may be found in Hermann W. Vogel’s “Practische 
Spectralanalyse iridischer Stoffe.” 1877. 410 
INORGANIC CHEMISTRY. 
If variously coloring substances be introduced into a flame, 
usually the most intense color obscures the others; in the 
spectrum, however, each individual substance shows its peculiar 
bright lines, which appear simultaneously or succeed each 
other, corresponding to the volatility of the various substances. 
The spectrum apparatus or spectroscope, figured in Fig. 89, 
serves to observe the spectra. 
In the middle of the apparatus is a flint-glass prism P. At 
the exterior end of the tube A is a movable vertical slit, in 
front of which is placed the source of 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) 
observed by the telescope B. To ascertain the relative posi- 
tion of the spectrum lines the tube C is employed. This at 
its outer end is provided with a transparent horizontal scale. 
Fm. 89. 
When a luminous flame is placed before the scale its divisions 
are reflected from the prism surface and reach into the tele- 
scope B. We see, then, in B, simultaneously, the spectrum to 
be investigated and the scale divisions, and can then readily 
determine the relative position of the spectrum lines. To 
study the spectra at the same time, and compare them, there is 
attached in front of one-half (the lower and upper) of the slit SPECTRUM ANALYSIS. 
411 
of the tube A a three-sided, right-angled glass prism, which 
directs the rays of a light source placed at the side (/, Fig. 89) 
through A upon the prism P. Then by means of B, two 
horizontal spectra will be observed, one above the other, and 
between the bright scale divisions. 
Adjustment of the Spectroscope.—To observe the spectra in the 
apparatus described, it is necessary to previously adjust the same cor- 
rectly. The tube A contains, besides the slit, also a lens (Collimator 
lens), which serves to make 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 tele- 
scope (B) is screwed out and adjusted to some distant object, that it 
may be adapted for the reception of parallel rays ; it is then replaced 
in the stand, pointed towards the slit, illuminated by a sodium chloride 
flame, and the slit then so far extended as to appear perfectly distinct 
in the telescope. To observe the spectrum lines as sharply as possible, 
the slit must be made quite narrow ; for feebly luminous lines it must, 
however, be widened. The horizontal black lines appearing in the 
spectrum arise from dust particles adhering to the slit. 
The proper position of the tube with the slit towards the prisms 
is usually fixed by the frame on which they rest. It must be so, 
that the refracted rays pass through the prism as symmetrically as 
possible, i. e-, in the minimum of their deviation, otherwise the spec- 
trum is less distinct and (owing to unequal refraction) will appear dis- 
torted. The symmetrical passage of the spectrum rays approximates 
most nearly if the medial green rays pass through the prism symmet- 
rically. Such a position, therefore, must be given the prism, with 
reference to the slit-bearing tube, that the middle green rays (line E of 
the sun spectrum) pass through in the minimum of their deviation. 
Then it is only necessary to so arrange the telescope that the green 
rays lie in the middle of the field of vision. 
The determination of the position of the lines of the spectra is usually 
effected by means of a scale (see above) and so arranging it (according 
to Bunsen) that the yellow sodium line coincides with the division 
line 50 on the scale; then the red potassium line (a) lies at 17, the 
violet (/?) at 152 (apparatus of Desaga). As, however, the refrac- 
tion and dispersion of the rays are influenced by the quality of the 
glass of the prism, the scale indications of different forms of appara- 
tus are not directly comparable. Hence they must be referred to an 
absolute measure. This is most conveniently attained by reduction to 
the sun spectrum, which, by means of the comparison prism, may be 
rendered at the same time visible in the telescope. The dark sun lines 
are now determined, with which the flame lines under investigation 
coincide. For accurate determinations, the spectrum lines are repre- 
sented in wave lengths, according to the millionth of a millimeter. 
The above described apparatus is that usually employed in chemical 
laboratories. In addition exist still others, accommodated to special 
purposes,—for investigation under the microscope, for the observation 412 
INORGANIC CHEMISTRY. 
of the sun and stars. For accurate observations, to obtain correct de- 
tailed, broad spectra, the light is permitted to pass through several 
(3-9) prisms, consisting of hollow glass filled with carbon disulphide, 
which refracts light very strongly. Instead of prism spectra, in many 
instances, diffraction spectra may be advantageously applied. 
The direct line (d vision directe) spectroscopes are very excellent 
for laboratory purposes, with which the spectra, without deflection, 
may be seen in the direction in which the luminous objects realty are. 
This is accomplished by a combination of several prisms of crown and 
flint glass, whereby dispersion, with simultaneous removal of deflection, 
is attained. 
To observe the spectra of metals in many instances, thus 
with the alkalies and alkaline earths, it is only necessary to 
introduce their volatile salts into a non-luminous alcohol or 
gas flame. A reduction of the metal usually takes place, and 
the spectra of the free metals themselves are obtained; thus, for 
example, sodium chloride is decomposed in the flame at first 
into HC1 and NaOH, which is then reduced by the carbon of 
the glowing gases to metallic sodium, which colors the flame 
yellow. The difficultly decomposable compounds (as the bari- 
um salts) frequently yield independent spectra, differing from 
those of the free metal; this is plainly recognizable in the 
copper compounds. 
However, most metals require for their conversion into gases 
much higher temperature than that of the gas flame. To 
vaporize them and observe their spectra, the electric spark 
is made to pass from electrodes constructed of them. In this 
manner all metals, even the most non-volatile, like gold, iron 
and platinum-, can be investigated. The spectra of these 
metals are generally quite complicated, and exhibit a great 
number of single lines. Thus, for iron, over 450 lines have 
been established. 
Instead of making the electrodes from the metals under study, we 
can, according to Bunsen (Poggend, Ann., 155), employ carbon points 
saturated with solutions of the metallic salts. To produce the electric 
spark, a Ruhmkorf induction apparatus with a dip battery of 4 
elements, and an attached Leyden jar, will suffice. Such spark spectra 
frequently distinguish themselves from flame spectra obtained in the 
gas flame. 
A peculiar, very interesting and practical procedure for the produc- 
tion of spark spectra, is due to Lecoq de Boisbaudran (Spectres 
lumineux, Paris, 1874). He allows the induction sparks to strike into 
the solution of the metal under consideration. This is placed in a 
small reagent tube, in the bottom of which is fused a platinum wire. 
Above the surface of the liquid is the second electrode, a platinum SPECTRUM ANALYSIS. 
413 
wire connected with the positive pole of the induction spiral. In this 
manner the spectra of all the metals may be easily obtained, and 
indeed, by it Lecoq de Boisbaudran discovered gallium. 
The spectra of the elementary gases may be determined by 
passing electric sparks through them, whereby these will 
be variously colored. Hydrogen illuminates with a red light, 
giving in the spectrum a bright red, a blue and a green line. 
Nitrogen shines with violet light and affords a spectrum of 
many lines, chief of which are the violet. The spectra of 
gases may be more conveniently observed by aid of Geissler’s 
tubes, which are filled with very dilute gases and the induc- 
tion stream then passed through them. 
By these methods we can very readily distinguish the in- 
dividual chemical elements, and even detect them in traces. 
By means of them, since the year 1860, different new 
elements: Caesium, rubidium, thallium, indium, scandium, 
gallium and several others which have not been accurately 
studied, have been discovered. 
In addition to the just described direct, bright spectra, 
there are yet dark absorption spectra. If a white light giving 
an uninterrupted spectrum be allowed to pass through 
different transparent bodies, the latter will absorb rays of 
definite refrangibility, allowing all others to pass. Therefore, 
in the spectroscope w'e observe the sun spectrum interrupted 
by dark lines or bands. Thus solutions of didymium and 
erbium absorb certain rays, and in the spectrum give corres- 
ponding dark lines. The gases deport themselves similarly. 
White light that has traversed a broad layer of air shows several 
dark lines in the spectrum peculiar to nitrogen, oxygen and 
steam. To a much higher degree is this power of absorption 
peculiar to all glowing gases or vapors. If a white light, 
like the Drummond calcium light, be conducted through the 
yellow sodium flame (through glowing sodium vapors), there 
will appear in the rainbow spectrum of the white light a dark 
line, which, as regards position, accurately corresponds to that 
of the yellow sodium line; the latter thus appears converted into 
a dark line. If white light be passed through the potassium 
flame, two dark lines will be visible in the spectrum, correspond- 
ing to the red and blue lines of the potassium spectrum. Such 
spectra are designated the inverted spectra of the correspond- 
ing metals. In this way the inverted spectra of all elements 
may be obtained, which accurately correspond to the direct 414 
INORGANIC CHEMISTRY. 
bright spectra. The cause of these phenomena lies in the prop- 
osition deduced by Kirchhoff from the undulatory theory of 
light, viz., the proportion between the power of emission and 
absorption for all bodies at equal temperatures is the same. 
According to this, glowing gases only absorb rays of just the 
same refrangibility which they emit. For example, bright 
white light is passed through the yellow sodium flame, the 
yellow rays of the same are absorbed and retained, while all 
others, almost entirely unaltered, pass on. Therefore, in the 
rainbow spectrum of white light the yellow rays of definite 
refrangibility fail; if, now, 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 phenomena have laid open to spectrum analysis, a 
new and broad domain, inasmuch as they open up avenues 
for the investigation of the chemical nature of the sun and 
other bodies. 
As is known, the bright rainbow sun spectrum is intersected 
by a quantity of dark lines v7hich have been called, from their 
discoverer, the Frauenhoff lines. The explanation of these 
becomes simple after what has been given in the preceding 
paragraphs, as shown by Kirchhoff, with the following hypo- 
thesis upon the nature of the sun. The latter is composed of 
a solid or liquid, luminous nucleus, surrounded by an atmos- 
phere of glowing gases and vapors. Therefore, the uninter- 
rupted spectrum of the glowing nucleus must be intersected 
by the dark lines of the inverted spectra of these gases and 
vapors, which are present in the sun’s atmosphere. An accu- 
rate comparison of the Frauenhoff lines with the spectrum 
lines of the various elements, has revealed the fact that in the 
sun’s atmosphere are chiefly found iron, sodium, magnesium, 
calcium, chromium, nickel, barium, copper, zinc and hydro- 
gen. Thus, for all the 450 lines of the iron spectrum have 
been discovered the corresponding dark lines in the sun’s 
spectrum. The conclusive inferences upon the chemical con- 
stitution of the sun, possess just such, indeed, a much higher 
degree of probability than is peculiar to many other deduc- 
tions. 
The investigation of the sun’s spectrum has cleared up many 
other changes occurring there, and on this was founded an en- 
tire sun-meteorology; all the fixed stars thus far investigated SPECTRUM ANALYSIS 
415 
possess a constitution like that of the sun. They give spectra 
intersected by dark lines ; therefore they consist of incan- 
descent nuclei surrounded by gaseous atmospheres. The 
spectra of nebulse only show bright lines; hence these con- 
sist of masses of glowing vapor yet uncondensed. 
Periodicity of the Spectrum Lines. As all other properties of 
the elements and their compounds proved themselves to be periodic 
functions of the atomic weights, the same is to be expected with 
reference to the spectrum phenomena. As yet but few such regularities 
have occurred. The most elements, chiefly the metalloids and the 
difficultly volatile metals, afford very complex spectra, which at the 
same time frequently vary with the temperature, so that for some of 
them spectra of 1st, 2d, etc., order are distinguished. It, however, 
appears that not all the spectra lines are of the same importance, as in 
several instances it has been possible to refer the various lines of a 
spectrum to particular fundamental lines whose relations to each other 
are comparable to that of the harmonies to primary tones. Thus the 
four lines of the hydrogen spectrum may be looked upon as the harmo- 
nies of a single wave. Therefore, in the comparison of spectra only in- 
dividual lines are to be regarded. This is clearly observed with the 
easily volatile metals belonging to the homologous groups K Rb Cs, Ca 
Sr Ba, Ga In Tl, whose lines lying in the violet portion of the spectrum 
proceed more towards the red end as the atomic weights increase ; 
with the latter, or that of the atomic volumes, the wave lengths (in 
millionths of a millimeter) become successively greater 
K 39 Wave length, 404 
Rb 83 “ “ 420.421 
Cs 132 “ “ 456.459 
Ca 40 Wave length, 422 
Sr 87 “ “ 461 
Ba 137 “ “ 525.550 
Ga 69 Wave length, 403.417 
In 113 “ “ 410.450 
T1 209 “ “ 535. 
By the arrangement of the spectra in the spectrum plate, these rela- 
tions are made apparent. 
A similar shifting of the spectral lines is observed with heterologous 
elements belonging to the same periods : K Ca, Rb Sr, Ba Cs, so 
that a conclusion upon the spectra of the succeeding elements appears 
possible. Indeed, the element scandium (45) succeeding calcium, shows 
intense violet lines of the wave length 425-440. 
P. 137, 4th line. Ins|e$'6f of P. 233, 22d line. 
Instead of Bn, read ‘jKX.'\ P..i">l, nth iine, the of pa£<ftomitted in (compare 
p. ) is 69. P. page omitted under Sails is Wy Atomic weight 
of strontium lie 87* not 872. 
ERRATA. HEAT OF FORMATION OF THE MOST IMPORTANT COMPOUNDS OF THE METALS. 
ACCORDING TO J. THOMSEN. 
In usual state of aggregation (columns a), and in dilute aqueous solution (columns b) 
a 
b 
a 
b 
a 
b 
a 
b 
a 
b 
a 
b 
a 
b 
- KC1 - 
105,6 
101,2 
KBr 
95,3 
90,2 
KI 
80,1 
75,0 
KOH 
103,9 
116,4 
KSH 
65,1 
k2so4 
344,5 
337,2 
kno3 
119,4 
111,4 
-NaCl- 
97,6 
96,5 
NaBr 
85,7 
85,6 
Nal 
69,1 
70,3 
NaOH 
102,0 
111,8 
NaSH 
— 
60,4 
Na*S04 
328,5 
329,0 
NaN03 
111,2 
106,2 
- LiCl 
93,8 
102,2 
LiBr 
— 
91,3 
Lil 
— 
76,1 
LiOH 
— 
117,4 
LiSH 
— 
66,0 
Li2S04 
334,1 
340,1 
LiN03 
111,6 
112,2 
= BaCl2 
194,2 
196,3 
BaBr2 
169,4 
174,4 
Bal2 
— 
144,0 
Ba02H2 
216,3 
226,6 
BaS2H2 
— 
107,1 
BaS04 
337,5 
— 
BaN206 
225,7 
216,3 
SrCl2 
184,5 
195,3 
SrBr2 
157,7 
173,8 
Srl2 
— 
143,3 
Sr02H2 
216,4 
226,1 
SrH2S2 
— 
106,6 
SrS04 
320,8 
— 
SrN206 
219,8 
215,2 
=: CaCl2 
170,2 
187,6 
CaBr2 
141,2 
165,7 
Cal2 
— 
135,3 
Ca02H2 
214,7 
217,8 
CaS2H2 
— 
98,3 
CaS04 
319,9 
— 
CaN206 
203,2 
207,1 
MgCl2 
151,0 
186,9 
MgBr2 
— 
165,0 
Mgl2 
— 
134,6 
Mg02H2 
217,2 
MgS2H2 
— 
97,8 
MgS04 
302,2 
323,0 
MgN206 
— 
206,3 
MnCl2 
111,9 
128,0 
MnBr2 
— 
105,1 
Mnl2 
— 
75,7 
A103H3 
296,9 
MnS, n(H20) 
46,3 
— 
ai2s3o12 
— 
878,9 
MnN2Ofi 
— 
147,5 
AlCla 
160,9 
237,7 
AlBr2 
129,6 
204,9 
aii3 
70,3 
159,3 
MnO, H20 
94,7 
ZnS, nH20 
41,5 
— 
MxiS04 
249,8 
263,6 
ZnN206 
— 
132,3 
- ZnCl2 
97,2 
112,8 
ZnBr2 
75,9 
90,9 
Znl2 
49,2 
60,5 
ZnO, H20 
82,6 
CdS, nH20 
33,9 
— 
ZnS04 
230,0 
248,4 
CdN206 
— 
115,8 
CdCl2 
93,2 
96,2 
CdBr2 
73,9 
74,3 
Cdl2 
44,9 
43,9 
H O (liquid) 
68,3 
FeS, nH20 
23,7 
— 
CdS04 
221,1 
231,8 
FeN206 
— 
119,5 
FeCl2 
82,0 
99,9 
FeBr2 
— 
78,0 
Fel2 
— 
47,6 
FeO, H20 
68,2 
CoS, nH20 
21,7 
— 
FeS04 
— 
235,6 
CoNA 
— 
114,3 
CoCl2 
76,4 
94,8 
CoBr2 
— 
72,9 
CoI2 
— 
42,5 
Fe203, 3H20 
191,1 
NiS,nH20 
19,3 
— 
CoS04 
— 
230,4 
NiN206 
— 
113,2 
NiCl2 
74,0 
93,7 
NiBr2 
— 
71,8 
Nil, 
— 
41,4 
CdO, H O 
65,6 
T12S 
21,6 
— 
NiS04 
— 
229,3 
HNOs 
49,1 
FeCl3 
96,0 
127,7 
FeBr3 
— 
— 
— 
— 
— 
CoO, H,0 
63,4 
PbS 
20,4 
— 
T12S04 
221,0 
212,7 
PbN206 
105,5 
97,9 
SnCl2 
80,8 
81,1 
SnBr2 
— 
— 
Sul, 
— 
— 
NiO, H20 
60,8 
CuS 
10,0 
— 
nm 
210,7 
T1N03 
58,1 
48,1 
HOI 
— 
39,3 
TlBr 
41,4 
— 
Til 
30,1 
— 
SnO, H20 
58,0 
Cu2S 
20,2 
— 
PbS04 
216,2 
— 
HN0, 
41,5 
— 
SnCl4 
127,2 
157,0 
HBr 
— 
28,3 
Pbl, 
39,6 
— 
Sn02, H20 
133,5 
HgS 
16,8 
— 
nm 
192,9 
CuN206 
— 
82,2 
' T1C1' 
48,5 
38,4 
PbBr2 
64,4 
54,4 
Hgl, 
34,3 
— 
HiO (vapor) 
58,0 
Ag2S 
53,1 
— 
CuS04 
182,5 
198,3 
HgN03 
— 
38,9 
' PbCl2 
82,7 
75,9 
HgBr 
34,1 
— 
Hgl 
24,2 
.— 
PbO 
50,3 
ES 
4,5 
9,2 
Ag2S04 
167,2 
162,7 
HgN206 
— 
67,1 
HgCl 
41,2 
— 
HgBr2 
50,5 
— 
Cul 
16,2 
— 
As203 
154,6 
147,0 
AgNOs 
28,7 
23,3 
HgCl2 
63,1 
59,8 
CuBr 
24,9 
— 
Cul, 
— 
— 
ti2o 
42,2 
39,2 
CuCl 
32,8 
— 
CuBr2 
32,5 
20,4 
Agl 
13,8 
— 
HgO 
30,6 
CuCl2 
51,6 
62,7 
SbBr3 
— 
— 
h! 
— 
13,1 
Hg20 
42,2 
SbCl3 
86,3 
— 
AgBr 
27,7 
— 
T1IS 
— 
10,5 
CuO 
37,1 
AgCl 
29,3 
— 
TlBrs 
— 
56,1 
AsI3 
12,6 
— 
Cu20 
40,8 
AsC13 
74,6 
—- 
AsBr3 
47,1 
— 
Aul 
—5,5 
— 
T1A,3H20 
86,0 
TlCls 
— 
89,0 
HBr 
8,4 
— 
HI 
-6,2 
— 
PdO, H20 
22,7 
HOI 
22,0 
— 
AuBr 
-0,1 
— 
Ptl2 
— 
Au A, 3H20 
—13,2 
AuCl 
5,8 
— 
AuBr3 
8,8 
5,0 
AuC13 
22,8 
27,2 
PtBr2 
— 
— 
PtCl2 
— 
— SSPtMX* 
PAGE 
Absorptiometer Ill 
Acetylene 145 
Acids..... 50, 75, 168 
Acid radicals 169 
Active oxygen 76 
“ chlorine 296 
Affinity, chemical 15 
Agate 227 
Alabandite 381 
Algar red, powder 137 
Alkaline metals 265 
“ earths 291 
Allotropy 79 
Alloys 252 
Alum 347, 382 
burned 348 
“ cubic 348 
“ Roman 348 
“ shales 348 
Alumina 346 
Aluminates 345 
Aluminite 347 
Aluminium 342 
“ bronze 343 
“ chloride 343 
“ oxide 344 
“ sulphate 347 
Alunite 348 
Amalgams 253 
Ammonia 115 
“ heat of formation of 139 
Ammonium 120 
“• anglesite 364 
“ chloride 288 
** compounds 287 
“ salts 348 
Anatase 229 
Anhydrides of acids 169 
Anhydride 296 
Anhydro-acids 186, 206, 227 
A 
PAGE 
Animal charcoal 141 
Anthracite 141 
Antimony 135 
“ butter 136 
“ chloride 136 
“ hydride 136 
“ oxide 215 
“ sulphides 216, 217 
Antimonic acid 216 
Antimonyl 215 
Apatite 297 
Apparatus of Carrfe 118 
“ Dcebereiner.... 38 
“ “ Marsh 133 
Aqua regia 195, 337 
Argentan 331, 401 
Arseniates 213 
Arsenic 131 
“ acid 212 
“ sulphides 213, 214 
Arsenious chloride 135 
“ oxide 211 
Arsenites «• 212 
Arsine 132 
Asbestos 308 
Atmosphere 108 
Atomic compounds 165 
“ heat 247 
“ volume 245 
“ weights..., 17, 158 
“ “ thermal 250 
Atoms 60, 68, 157 
Aurates 340 
Auric acid 339 
Auripigment 213 
Aurous compounds 338 
Auric “ 339 
Avogadro, law of. 68 
Azote 109 
Azurite 330 418 
INDEX 
B PAGE 
Barium 301 
“ peroxide 302 
“ sulphate 302 
“ chromates 373 
Baryta water 302 
Bases '50, 75 
Basicity of acids 168, 208 
Bauxite 345 
Bell metal 331 
Blue, Prussian 396 
Beryllium 309 
Bessemer steel 389 
Bismuth 137, 365 
Bismuthic acid 366 
Bismuthinite 365 
Bitter salt 307 
“ water 82 
Bleaching 42, 179 
Boiler incrustation 298 
Boracite 231 
Boron 230 
“ fluoride 232 
Brass 330 
Braunnite 380 
Bromine 43 
“ hydrate 44 
“ hydride 52 
Bromic acid 173 
Bronze 331 
Brown coal 141 
Burned alum 348 
C 
Cadmium 312 
Caesium 276 
Calamine 312 
Calcium 292 
“ carbonate 298 
“ manganite 383 
“ oxide 292 
“ silicate 298 
“ sulphides 299 
‘ ‘ sulphate 296 
Calcite 392 
Calomel 318 
Calorie 57, 80 
Caput mortuum 393 
Carbon 140 
“ chlorides 151 
“ dioxides 219 
B 
PAGE 
Carbon disulphide 223 
“ group 140 
“ monoxide 222 
“ oxychloride 223 
Carbonates 222 
Carbonic acid 222 
Carry’s 118 
Cast iron 388 
“ steel 389 
Cassiterite—Tin-stone 357 
Caustic potash 274 
“ soda 278 
Cement 293 
Cementation steel 389 
Cerium 351 
Cerite 350 
Cerussite 361, 364 
Chalk 298 
Chamber acid 182, 184 
Charcoal, animal 141 
Chemical affinity , 15 
. “ energy 15 
“ structure 160 
Chloranhydrides 187 
Chloric acid 172 
Chlorides of sulphur.... 101, 102 
Chlorine 39 
“ gas 383 
“ hydrates 41 
“ iodide 58 
“ oxides 169 
“ water 41 
Chloroform 151 
Chlorous acid 171 
Chromates 371 
Chromic acid 372 
Chromite 373 
Chromium group 367 
“ 368 
“ alum 370 
“ oxychloride 374 
“ yellow 374 
Chrysoberyl 346 
Cinnabar 321 
“ green 398 
Cobalt 397 
“ cyanides 399 
Cobalt rose 399 
“ ultramarine 398 
Cobaltite 397 INDEX 
419 
PAGE 
Cobaltous compounds 397 
Cobaltic “ 398 
Coke 141 
Colcothar 393 
Colloids 227 
Compounds atomic 165 
“ definite 83 
“ endothermic 57 
“ exothermic 57 
“ molecular 165 
Copper 325, 330 
“ hydride 328 
“ vitriol 329 
Corundum 242 
Critical temperature 219 
Crocoisite 361, 368 
Cryolite 283, 342 
Crystallization, water of 83 
Crystallography 20 
Crystalloids 227 
Cupellation 332 
Cuprammoniuin compounds, 329 
Cupric compounds. 328 
Cuprite 327 
Cuprous compounds 327 
Cyanogen 224, 394 
“ compounds 408 
Corrosive sublimate 319 
Cubical alum 348 
D 
Davy’s lamp 150 
Decipium 341 
Dialysis.. 227 
Diamond . 140 
Dicyanogen 225 
Didymium 351 
Diffusion of gases 113 
Dissociation 85, 96, 130, 184, 198 
Disodium phosphate 284 
Disulphuric acid 186 
Dithionic acid 190 
Doebereiner’s lamp 38 
Drummond’s light 75 
E 
Earth metals 291 
Earths alkaline 291 
“ rare 350 
Epsom salt 373 
PAGE 
Ekaboron 350 
Eka aluminium 352 
Electrolysis of salts 260 
Elements 12 
“ classification of.... 30 
Emery 344 
Equivalence ! 162, 164 
“ of forces 14 
Equivalent weights 162 
Erbium *.... 351 
Ethyl 144 
Ethylene 114 
Eudiometer Ill 
Euxenite 350 
P 
Feldspar 340 
Ferric acid 394 
“ compounds 392 
Ferricyariogen 395 
Ferrous compounds 391 
Ferrocyanogen 395 
Flame 145 
“ illuminating power ofi 147 
“ non-luminosity of. 150 
Flores zinci 311 
Flint 299 
Fluorine 46 
Fluorite 292 
Fulminating silver 333 
a 
Gadolinite 350 
Gahuite 346 
Galenite 332, 361 
Gallium 352 
“ group 351 
Gas densities 62 
“ laughing 204 
“ volume 64 
Gases, condensation of 35 
“ diffusion of. 113 
“ kinetic theory of 64 
“ measuring of. 112 
Gersdorflite 400 
Glass 299 
Glover’s tower 182 
Glucinum 309 
Gold 337 
“ purple 340 420 
INDEX 
PAGE 
Gold fulminating 340 
“ sulphide 340 
Granite 349 
Graphite 140, 141 
Greenockite 313 
Guignet’s green 330, 370 
Gypsum 296 
H 
Halogen compounds 253 
' hydrides 47 
Halogens 39 
Haloid acids 51 
“ salts 51 
Hausmannite 380, 382 
Heat, latent 81 
“ solution of. 83 
“ specific 274 
“ unit of. 80 
“ atomic 274 
Hematite 392 
Hornsilver 331, 334 
Hydrargillite 345 
Hydrates 75, 174, 254 
“ of acids 185 
Hydrobromic acid 52 
“ chloric “ 47 
“ cyanic “ 225 
“ carbons 142 
“ fluoric acid 55 
“ sulphurous acid 180 
Hydrogen 31 
“ chloride 47 
“ cyanide 225 
“ fluoride 55 
“ iodide 53 
“ peroxide 90 
“ persulphide 100 
“ sulphide 97 
Hydroxides 75, 254 
Hydroxylamine 121 
Hypochlorites... 170 
Hypochlorous acid 170 
Hypophosphites 206 
Hypophospliorous acid 206 
Hyponitrous acid 204 
Hyposulphites 189 
Hyposulphurous acid 189 
I 
Ice machine of Carrfe 118 
PAGE 
Indium 853 
Iodic acid 173 
Iodine 44 
Iridium 402, 404 
Iron 387 
“ alum 394 
“ cast 388 
“ fused 392 
“ group . 386 
“ pyrite 391 
“ pyrophoric 390 
“ vitriol 391 
“ wrought 388 
Isomerism 79 
Isomorphism 251 
K 
Kaolin 342, 349 
Kelp 45 
Kermes minerale 217 
Kieserite 306 
L 
Lanthanum 341, 350 
Lapis lazuli 349 
Laughing gas 204 
Lead 356, 361 
“ chamber crystals 184 
“ oxide 362 
“ peroxide 363 
“ sulphide 363 
“ tree 362 
“ vitriol (sulphate) 364 
“ white 363 
“ alloys 362 
Lime 293 
“ stone 292 
“ water 293 
“ chloride of. 294 
Litharge 362 
Lithium 288 
Lunar caustic 336 
Luteocobalt chloride 399 
M 
Magnesia 306 
Magnesite 308 
Magnesium 305 
“ group 3C3 
Magnetite 393 
Malachite 330 INDEX 
421 
PAGE 
Manganese 379 
Manganic acid 880 
salts 381 
Manganite 380 
Marble 292 
M arsh ’ s apparatu s 133 
Marsh gas..... 142 
Massicot 363 
Maximum valence 165 
Mercuric compounds 319 
Mercurous “ 318 
Mercury 314, 316 
Meta-acids 192 
Meta-antimonic acid 216 
Metahydrates 255 
Metals 244 
Metalloids 12, 30 
Metaphosphoric acid 208 
Metastannic acid 358 
Methyl 142 
Mineral waters 82 
Molecular theory 69 
“ compounds 165 
Molecule 68, 69, 157 
Molecules of elements 69, 79, 96, 
125 
Molybdenum 375 
Molybdenite 375, 376 
Molybdic acid 376 
Mortar 293 
Mosaic gold 361 
Multiple proportions 62 
Magisterium bismuthi 367 
N 
Nickel 400 
“ plating 401 
Niobium 137 
Nitrates 193 
Nitric acid 193 
“ “ anhydride 196 
“ “ fuming 195 
Nitrogen 106 
Nitrogen chloride 122 
“ iodide 123 
“ oxide 201 
“ pentoxide 196 
“ tetroxide 197 
Nitrosulphuric acid 198 
Nitrosylchloride 201 
PAGE 
Nitrosylsulphuric acid 198 
Nitrous acid 197 
Nordhausen sulphuric acid... 186 
O 
Orthoclase 349 
Olefiant gas 145 
Organic chemistry 142 
Ores 310 
Orthite 350 
Ortho-acids 192 
“ nitric acid 194 
“ phosphoric acid 208 
Osmium 401, 403 
Oxidation 75 
Oxides 254 
“ basic 75 
“ acid-forming 75 
Oxygen 71 
Ozone 76 
P 
Palladium 401, 405 
“ hydride 39, 405 
Passive iron 390 
Pattison’s method 332 
Pentathionic acid 190 
Periodates 175 
Perchloric acid 172 
Periodic system 233 
Periodicity of heat of forma- 
tion 242 
“ “ spectral lines.. 415 
“ “ valence 240 
Periodic acid 174 
Permanent white 365 
Permanganic acid 379, 383 
Peroxides 255 
Phillipium 341 
Phlogiston 109 
Phosgene 223 
Phosphates 206 
Phosphine 126 
Phosphites 206 
Phosphonium 128 
“ iodide !. 128 
Phosphorus 123 
“ bromides 131 
‘‘ chlorides 129 
“ iodides 131 422 
INDEX, 
PAGE 
Phosphorus oxychloride 131 
“ pentoxide.....205, 209 
“ sulphides 211 
Phosphorous acid 207 
Phosphoric u 208 
Phosphorite 297 
Photography 335 
Platinum 406 
“ bases 408 
“ metals 401 
“ ores 402 
“ sponge 38, 402, 406 
Plumbic acid 363 
Polyacids 186, 227 
“ silicic acid 227 
“ thermic “ 188 
Porcelain 349 
Potashes 273 
Potassium 267 
“ amide 276 
“ aurate 340 
“ bromide 268 
“ chloride 269 
“ cyanide 270 
“ fluoride 270 
“ hydrate 268 
“ iodide. 270 
“ manganate 352 
“ oxide 268 
“ plum bate 363 
“ saltpetre 272 
Prussiate of potash, yellow... 395 
“ “ red 396 
Purpureocobalt chloride 399 
Pyrolusite 380, 382 
Pyrites 181, 394 
Pyroacids s. polyacids.. 186, 227 
Pyrophosphoric acid 208 
Pyrosulphuric “ 186 
Q 
Quantivalence .. 160 
Quartz 226 
Quadrant oxides 256 
Quicksilver 320 
R 
Red chromate of potash 373 
Red lead 363 
Rare metals 350 
PAGE 
Radicals s. residues 169 
Realgar 214 
Reduction 76 
Residues 169 
Respiration 221 
Rhodium 404 
Rhodochroisite 380, 381 
Rock salt 278 
Roseo-cobalt 399 
Rose’s metal 367 
Rubidium 276 
Ruby 342 
Ruthenium 401, 403 
Rutile 229 
Roman alum 348 
S 
Safety lamp 150 
Salts 54, 75, 168, 256 
“ basic 257 
Saltpetre 272 
Salt producers 39 
“ springs 278 
Sapphire 342 
Sassolite 232 
Scandium 350 
Scheele’s green, 330 
Scheelite 377 
Schlippe’s salt 217 
Schweizer’s reagent 328 
Sea water 82 
Selenic acid 191 
Selenium 102 
Silicates 227 
Silicic acid 226 
Silicon 152 
“ chloride 153 
“ chloroform 154 
“ dioxide.. 226 
“ disulphide 228 
“ fluoride 155 
“ hydride 152 
Silver 331 
“ cyanide 835 
“ nitrate 335 
u oxides 333 
Silvering 336 
Slags 381 
Smaltite 397 
Smalts 398 423 
PAGE 
Smithsonite 392 
Soda 282 
“ caustic 278 
Sodium 277 
“ carbonate 282 
“ chloride 278 
“ hydrate 277 
“ hyposulphite 281 
“ iodate 279 
“ saltpetre 284 
“ sulphate 277 
“ oxide 277 
“ phosphates 284 
“ silicate 286 
Soft solder 362 
Specific gravity of gases...36, 63 
“ “ “ metals 244 
“ heat 247 
“ volume 245 
Spectrum analysis 409 
Spiegeleisen 388 
Stalactite 298 
Stannates 356 
Stannic compounds 360 
Stanniol 357 
Stannous compounds 359 
Stassfurt salts 267 
Status nascendi 43, 69 
Steel 388 
“ cast 389 
“ Bessemer 389 
Stibnine 136 
Stibnite 135 
Stone coal 140 
Stolzite 377 
Strass 299 
Strontianite 300 
Strontium 300 
Structure, chemical 160 
Substitution 151 
Sulphur auratum 217 
Sulphates 177 
Sulphur 94 
chlorides 101 
“ dioxide 177 
“ heptoxide 180 
“ sesquioxide 180 
“ trioxide 181 
Sulpho-stannate 359 
Sulphuretted hydrogen 97 
INDEX 
PAG B 
Sulphuric acid 181 
“ anhydride 181 
“ “ English.. 184 
“ “ fuming... 186 
Sulphides 99 
Sulphites 179 
Sulphurous acid 179 
Sulphuryl 188 
“ chloride 188 
Sylvanite 337 
Sylvite 269 
Sympathetic ink 397 
T 
Talmi gold 331 
Tantalum 137 
Telluric acid 191 
Telluride, hydrogen 104 
Tellurium 103 
Tension of vapors 81 
Terbium 341 
Tetanic acid 191 
Tetranoxide 328 
Tetrathionic acid 190 
Thallic acid 356 
Thallium 354 
“ glass 356 
Thenard’s blue 398 
Thionyl chloride 179 
Thiosulphuric acid 189 
Thorium 228, 230 
Thulium 18, 341 
Tin 356 
“ stone 357 
Tinkal 285 
Titanium 229 
Tombac 331 
Trithionic acid 190 
Turnbull’s blue 396 
Turpetum mineral 321 
Type metal 362 
Tungsten 367, 377 
“ steel 377 
Tungstic acid 377 
U 
Ultramarine 349 
Uranates 378 
Uranium 378, 367 
“ yellow 379 424 
INDEX 
PAGE 
Uranium glass 379 
Uranyl 378 
V 
Valence 160 
Vanadium 137 
Vapor density 45 
Varec 45 
Vivianite 392 
Vitriol, oil of, see Sulphuric 
acid 
Vitriol, iron 391 
W 
Water 80 
Water glass 226 
“ of constitution 307 
White iron 387 
PAGE 
White lead 364 
Wolfram 377 
Wollastonite 299 
Wrought iron 388 
Wolfenite 361 
Y 
Ytterbium 351 
Yttrium 351 
Yellow prussiate of potash.... 395 
Z 
Zinc 310 
“ blende 312 
“ dust 310 
“ vitriol 311 
“ white 311 
Zircon 228, 230 INORGANIC CHEMISTRY, 
A TEXT-BOOK FOR STUDENTS. 
By PROF. VICTOR V. RICHTER, 
University of Breslau. 
Authorized Translation from the Third German Edition, 
By EDGAR F. SMITH, M.A., Ph.D., 
Prof, of Chemistry in Wittenberg College, Springfield, Ohio ; formerly in the Laboratories of 
the University of Pennsylvania and Muhlenberg College ; Member of the Chemical Societies 
of Berlin and Paris, of the Academy of Natural Sciences of Philadelphia, etc., etc. 
S9 Wood-cuts and Colored Lithographic Plate of Spectra. 
12mo, Cloth, 421 Pages. Price $2 50. 
In most of the chemical text-books of the present day, one of the 
striking features and difficulties with which teachers have to contend 
is the separate presentation of the theories and facts of the science. 
These are usually taught apart, as if entirely independent of each 
other, and those experienced in teaching the subject know only too 
well the trouble encountered in attempting to get the student properly 
interested in the science and in bringing him to a clear comprehension 
of the same. In this work of Prof, von Richter, which has been 
received abroad with such hearty welcome, and for which there has 
been a great demand, two editions having been rapidly disposed of, 
theory and fact are brought close together, and their intimate relation 
clearly shown. From careful observation of experiments and their 
results, the student is led to a correct understanding of the interesting 
principles of chemistry. The descriptions of the various inorganic 
substances are full, and embody the results of the latesf discoveries. 
The periodic system of Mendelejeff and Lothar Meyer constitutes 
an important feature of the book. The thermo-chemical phenomena 
of the various groups of elements also receive proper consideration, 
both in their relation to chemical affinity and the law of periodicity. 
The matter is so arranged as to adapt the work to the use of the be- 
ginner as well as for the more advanced student of chemical science. 
The arrangement of types, size of the book, etc., are such as to 
facilitate its use. 
Favorable terms will be made for the adoption of this ivork in 
schools and colleges. Professors of chemistry are invited to corres- 
pond with the publishers in reference to its introduction, and for 
further information. 
P. BLAKISTON, SON & CO., Publishers and Booksellers, 
MEDICAL AND SCIENTIFIC BOOKS, 
JVo. 1012 Walnut Street, Philadelphia. From JOHN MARSHALL, M.D., Nat. Sc.D. (Tubingen), Demon- 
strator of Chemistry in the University of Pennsylvania. 
Knowing the value of the excellent work of Prof, von Richter in 
the original, I am pleased to hear of a forthcoming English transla- 
tion, by such an able chemist and admirable translator as Prof. Smith. 
The work is of undoubted value. The theory of chemistrjq which 
is generally the bugbear of students, is in this book very clearly 
explained, and the explanations are so well distributed through the 
book that students are brought easily from the simplest to the most 
difficult problems. 
That part descriptive of the elements and their compounds is full, 
and all that could be desired in a text-book, while the cuts, with which 
the work is profusely illustrated, are an excellent aid to the student. 
Altogether, it is one of our best modern works on chemistry. 
Muhlenberg College, Chemical Laboratory, 
Allentown, Pa. 
This volume is especially adapted for use as a text-book; and, if for 
no other reason than for its superior arrangement, should find its way 
into our best schools and colleges. The author has apparently accom- 
plished the difficult task of furnishing to the student a book that 
combines, with the utmost clearness, the two departments of general 
chemistry, in such a way as to remove the trouble which has for a 
long time been experienced in applying the theoretical knowledge 
obtained to the explanation of the inorganic portion of the work. In 
this volume, however, the two departments are treated at one and the 
same time. The name of the translator should insure it a cordial 
reception by all American chemists. 
N. WILEY THOMAS. 
I am well acquainted with von Richter’s Inorganic Chemistry in its 
original German form. Its success abroad has been exceptional, 
having rapidly run through several large editions in Germany, and 
having been translated into at least two other languages. 
With the large number of text books on Inorganic Chemistry at 
present in the field, this success could only have been achieved be- 
cause of distinctive features which merited commendation. 
These we find to be the clear presentation of simple underlying 
theory and succinct statements of illustrative facts; but above all, in 
the well chosen order in which present views of chemistry are de- 
veloped in successive steps. 
Dr. Smith’s previous successes in the role of translator and editor 
of German works on chemistry insure an accurate reproduction of 
the original. 
We have no doubt that the work will be as well received by Ameri- 
can chemists as it has been abroad. 
Prof, of General and Organic Chemistry, University of Penn’a. 
SAM’L P. SADTLER,