A TEXT BOOK 
OF 
INORGANIC CHEMISTRY. 
RICHTER. NOW READY. 
The Chemistry of the Carbon 
Compounds; 
OR, 
ORGANIC CHEMISTRY. 
PROF. VICTOR von RICHTER, 
UNIVERSITY OF BRESLAU. 
Authorized Translation from the Fourth German Edition, 
EDGAR F. SMITH, M.A., Ph.D., 
Translator of Richter's Inorganic Chemistry; 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. 
Illustrated. i2mo. 725 Pages. Cloth, $3.00. 
In this volume Prof, von Richter gives a full description of the Carbon De- 
rivatives. The treatment of the various classes of bodies is comprehensive, 
the most recent and interesting results of chemical research being given with 
special fulness. 
Excellent methods of preparing the most typical bodies will be found dis- 
tributed through the entire volume. These will enable the student to acquaint 
himself practically with those portions of organic work which are now receiv- 
ing so much attention. The reactions of the groups of compounds are com- 
plete, the methods of analysis described are the most modern and reliable, so 
that the volume may be depended on as a trustworthy laboratory guide. 
Numerous references to the literature of the subject will be found upon almost 
every page. 
The arrangement of types allows of the book being used by both beginners 
and advanced students, the more general facts being printed in large type, and 
may be employed for common class use. The smaller type covers the matter 
usually given in more advanced lectures. 
ff-iy Examination copies at a reduced price. Correspondence invited from 
teachers and professors. 
P. BLAKISTON, SON & CQ., Publishers, Philadelphia. A 
TEXT-BOOK 
OF 
INORGANIC CHEMISTRY. 
■ BY 
PROF. VICTOR von RICHTER, 
UNIVERSITY OF BRESLAU. 
AUTHORIZED TRANSLATION, 
BY 
EDGAR F. SMITH, 
PROFESSOR OF CHEMISTRY IN WITTENBERG COLLEGE, SPRINGFIELD, OHIO. 
SECOND AMERICAN FROM THE FOURTH 
GERMAN EDITION. 
WITH EIGHTY-NINE ILLUSTRATIONS ON WOOD 
AND 
COLORED LITHOGRAPHIC PLATE OF S H A," 
PHILADELPHIA: 
P. BLAKISTON, SON & CO., 
No. xoi2 Walnut Street. 
1 885. Copyrighted, 1885, by P. Blakiston, Son & Co. PREFACE 
TO THE 
SECOND AMERICAN EDITION. 
The present edition is a translation of the fourth German 
edition. Many parts of the work have been rewritten, and 
new matter incorporated. The features which recommended 
the previous edition have been preserved, and as now pre- 
sented it is hoped that the work may continue to be of 
service to the student of chemical science. 
The translator would here express his obligations to Mr. 
Allen J. Smith, who has read the entire proof, revised the 
index and table of contents, as well as rendered other valuable 
assistance. PREFACE 
TO THE 
FIRST AMERICAN EDITION. 
The success of Prof, von Richter’s work abroad would 
indicate its possession of more than ordinary merit. This we 
believe true, inasmuch as, in presenting his subject to the 
student, the author has made it a point to bring out promi- 
nently the relations existing between fact and theory. These, 
as well known, are, in most text-books upon inorganic chem- 
istry, considered apart, as if having little in common. The 
results attained by the latter method are generally unsatisfac- 
tory. 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 
connection between their results and the theories based upon 
them, is obviously preferable to their separate study, especially 
when they are treated in widely removed sections or chapters 
of the same book. Judging from the great demand for von 
Richter’s work, occasioning the rapid appearance of three 
editions, the common verdict would seem to be unanimously 
in favor of its inductive methods. 
In the third edition, of which the present is a translation, 
the Periodic System of the Elements, as announced by Men- 
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- IO 
PREFACE TO THE FIRST AMERICAN EDITION. 
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 
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 OF CONTENTS. 
INTRODUCTION. 
Province of Chemistry, 17. Definition of Chemistry, 19. Chemical Ele- 
ments, 20. Principle of Indestructibility of Matter, 20. Principle 
of Conservation of Energy, 21. Chemical Energy, 23. Conditions 
of Chemical Action, 24. Chemical Symbols and Formulas, 25. 
Crystallography, 28. 
SPECIAL PART. 
Classification of the Elements, 37. 
Hydrogen, 38. Purifying and Drying of Gases, 40. Apparatus 
for the Generation and Collection of Gases, 40. Condensation of 
Gases, 42. 
Group of Halogens, 46. 
Chlorine, 46. Bromine, 50. Iodine, 51. Fluorine, 53. General 
Characteristics of the Halogens, 53. Compounds of the Halogens 
with Hydrogen, 54. Hydrogen Chloride, 54. Acids, Bases, Salts, 
57. Hydrogen Bromide, 59. Hydrogen Iodide, 60. Hydrogen 
Fluoride, 62. General Characteristics of the Hydrogen-Halogen 
Compounds, 64. Thermo-Chemistry of the Hydrogen-Halogen Com- 
pounds, 64. Compounds of the Halogens with Each Other, 66. 
Weight Proportions in the Union of the Elements, Law of Constant Pro- 
portions, Atomic Hypothesis, 66. Density of Bodies in State of Gas, 
Volume Proportions in the Union of Gases, Atomic Molecular The- 
ory, 70. Avogadro’s Law, 75. Status Nascens, 76. 
Oxygen Group, 78. 
Oxygen, 78. Oxidation and Reduction, 81. Ozone, 83. Isomerism 
and Allotropy, 86. Compounds of Oxygen with Hydrogen, 87. 
Water, 87. Critical Pressure, 88. Solutions, 90. Thermo-Chemistry 
of Solutions, 91. Dissociation, 92. Quantitative Composition of 12 
CONTENTS. 
Water, 94. Molecular Formula of Water, Atomic Weight of Oxy- 
gen, 95. Hydrogen Peroxide, 97. Sulphur, 102. Molecules of the 
Elements, 104. Compounds of Sulphur with Hydrogen, 105. Hy- 
drogen-Sulphide, 105. Molecular Formula of Hydrogen-Sulphide, 
Atomic Weight of Sulphur, 107. Hydrogen Persulphide, 108. Com- 
pounds of Sulphur with the Halogens, 109. Selenium, no. Tellu- 
rium, in. Summary of the Elements of the Oxygen Group, 112. 
Thermo-Chemistry of their Hydrogen Compounds, 113. 
Nitrogen Group, 114. 
Nitrogen, 114. Atmospheric Air, 116. Measuring Gases, 120. 
Diffusion of Gases, 121. Compounds of Nitrogen with Hydrogen, 
123. Ammonia, 123. Ammonium Salts, 127. Quantitative Composi- 
tion of Ammonia, Atomic Weight of Nitrogen, 127. Hydroxylamine, 
128. Compounds of Nitrogen with the Halogens, 129. Phosphorus, 
130. Compounds of Phosphorus with Hydrogen, 133. Compounds 
of Phosphorus with the Halogens, 136. Arsenic, 139. Arsine, 139. 
Compounds of Arsenic with the Plalogens, 142. Antimony, 143. 
Stibine, 143. Compounds of Antimony with the Halogens, 143. 
Characteristics of the Elements of the Nitrogen Group, 145. Their 
Thermo-Chemical Deportment, 145. 
Carbon Group, 146. 
Carbon, 146. Carbon Compounds of Hydrogen, 148. Methane, 
148. Atomic Weight of Carbon, 149. Ethane, 150. Ethylene, 150. 
Acetylene, 151. Nature of Flame, 151. Compounds of Carbon with 
the Halogens, 157. Silicon, 158. Hydrogen Silicide, 159. Com- 
pounds of Silicon with the Halogens, 160. Hydrogen Silico-Fluo- 
ride, 162. 
Atom and Molecule, 163. Determination of Molecular Value from Chem- 
ical Reactions, 165. Valence of the Elements, Chemical Structure, 166. 
Equivalence, 168. Maximum Valence, 170. 
Oxygen Compounds of the Metalloids, 173. 
Oxygen Compounds of the Halogens, 174. 
Oxides of Chlorine, 175. Hypochlorous Oxide, 175. Hypochlorous 
Acid, 176. Chlorine Trioxide, Chlorous Acid, Chlorine Tetroxide, 
176. Chloric Acid, 177. Perchloric Acid, 178. Oxides of Bro- 
mine, 178. Oxides of Iodine, 179. Hydrates of the Acids, 180. 
Thermo-Chemistry of the Oxygen Compounds of the Halogens, 
181. 
Oxygen Compounds of the Elements of the Sulphur Group, 182. 
Oxygen Compounds of Sulphur, 183. Sulphur Dioxide, 183. Sul- 
phurous Acid, Hyposulphurous Acid, 185. Sulphur Trioxide, 186. Sul- CONTENTS. 
13 
phuric Acid, 187 Pyrosulphuric Acid, 191. Sulphuric Acid Chlor- 
anhydrides, 192. Polythionic Acids, 194. Thiosulphuric Acid, 195. 
Oxygen Derivatives of Selenium and Tellurium, 196. Thermo- 
Chemistry of the Oxides and Acids of the Sulphur Group, 197. 
Oxygen Compounds of the Elements of the Nitrogen Group, 198. 
Oxygen Derivatives of Nitrogen, 199. Nitric Acid, 199. Nitrogen 
Pentoxide, Nitrogen Trioxide, 202. Nitrous Acid, 203. Nitrogen 
Tetroxide, 204. Nitrosyl-sulphuric Acid, 205. Nitric Oxide, 207. 
Nitrous Oxide, 209. Hyponitrous Acid, 210. Thermo-Chemistry of 
the Oxides and Acids of Nitrogen, 211. Oxides of Phosphorus, 211. 
Hypophosphorous Acid, 212. Phosphorous Acid, 213. Phosphorus 
Trioxide, Pyrophosphoric Acid, 214. Metaphosphoric Acid, 215. 
Phosphorus Pentoxide, 216. Chlor-anhydrides of the Acids of Phos- 
phorus, 217. Phosphorus Compounds with Sulphur, 217. Oxides of 
Arsenic, 218. Arsenic Trioxide, 218. Arsenious Acid, Arsenic 
Acid, 219. Compounds of Arsenic with Sulphur, 220. Sulpho-Salts, 
221. Oxygen Derivatives of Antimony, 221. Antimony Oxide, 221. 
Antimonic Acid, 222. Antimony Sulphides, 223. Thermo-Chemis- 
try of the Acids of the Nitrogen Group, 224. -Vanadium, Niobium, 
Tantalum, 224. 
Oxygen Compounds of the Elements of the Carbon Group, 225. 
Oxides of Carbon, 226. Carbon Dioxide, 226. Carbon Monoxide, 
229. Compounds of Carbon with Sulphur, 231. Cyanogen Com- 
pounds, 232. Thermo-Chemistry of the Carbon Compounds, 233. 
Oxygen Compounds of Silicon, 233. Dialysis, 234. Crystalloids 
and Colloids, 235. Silicates, 235. Titanium, Zirconium, Thorium, 
236. Boron, 237. Boron Hydride, 238. Boron Trichloride, 239. 
Boron Fluoride, 239. Boric Acid, 239. 
The Periodic System of the Elements, 240. 
Periodicity of Chemical Valence, 247. Periodicity of Thermo-Chem- 
ical Phenomena, 249. , 
THE METALS. 
Physical Properties of the Metals, 251. Atomic Volumes, 252. Specific 
Heat, Atomic Heat, 254. Thermal Atomic Weights, 257. Isomor- 
phism, 258. Chemical Properties of the Metals, 259. Alloys, 259. 
Halogen Compounds, 260. Oxides and Hydroxides, 261. Peroxides, 
262. Salts, 263. Action of Metals on Salts and Acids, 265. Elec- 
trolysis of Salts, 267. Transposition of Salts, i70. Principle of the 
Greatest Heat-Disengagement, 271. 14 
CONTENTS. 
Group of the Alkali Metals, 273. 
Thermo-Chemistry of the'Alkali Metals, 274. Potassium, 275. Potas- 
sium Oxide, Potassium Hydroxide, 276. Potassium Chloride, Potas- 
sium Bromide, Potassium Iodide, 277. Potassium Cyanide, 278. Potas- 
sium Chlorate, 278. Potassium Hypochlorite, 279. Potassium Sulphate, 
279. Potassium Nitrate, 280. Gunpowder, 281. Potassium Carbonate, 
281. Potassium Silicate, 282. Potassium Sulphides, 283. Recognition 
of the Potassium Compounds, 284. Rubidium, Caesium, 284. Sodium, 
284. Sodium Hydroxide, 285. Sodium Chloride, 286. Sodium 
Sulphate, 287. Supersaturated Solutions, 288. Sodium Hyposul- 
phite, 289. Sodium Carbonate, 290. Sodium Nitrate, Sodium 
Phosphates, 292. Borax, 293. Sodium Silicate, 294. Recognition 
of Sodium Compounds, 294. Lithium, 294. Ammonium Com- 
pounds, 295. Ammonium Chloride, Ammonium Carbonate, 296. 
Ammonium Phosphates, Ammonium Sulphide, 297. Recognition of 
Ammonium Compounds, 298. 
METALS OF GROUP II, 298. 
Group of the Alkaline Earths, 299. 
Calcium, 300. Calcium Oxide, Calcium Hydroxide, 300. Cement, 
301. Calcium Chloride, 301. Calcium Fluoride, Chloride of Lime, 
302. Calcium Sulphate, 303. Calcium Phosphates, 304. Calcium 
Carbonate, 305. Glass, 306. Calcium Sulphides, 307. Strontium, 
307. Barium, 308. Barium Oxide, 308. Barium Peroxide, Barium 
Sulphate, 309. Recognition of the Compounds of the Alkaline 
Earths, 310. 
Magnesium Group, 310. 
Magnesium, 312. Magnesia, 312. Magnesium Chloride, 313. 
Magnesium Sulphate, Magnesium Phosphates, 314. Magnesium 
Carbonate, 315.- Recognition of Magnesium Compounds, 315. 
Beryllium, 316. Zinc, 317. Zinc Oxide, 317. Zinc Chloride, Zinc 
Sulphate, Zinc Sulphide, 318. Cadmium, 319. Thermo-Chemistry 
of the Metals of Group II, 320. Mercury, 321. Amalgams, 324. 
Mercurous Compounds, 324. Mercuric Compounds, 326. 
Copper, Silver, Gold, 328. General Characteristics, 328. Forms of 
. Combination, 330. Copper, 332. Metallurgy of Copper, 332. Cu- 
prous Compounds, 333. Cupric Compounds, 335. Copper Sulphate, 
335. Alloys of Copper, 337. Silver, 337. Metallurgy, 338. Silver 
Oxide, 339. Silver Chloride, 341. Photography, 341. Nitrate of 
Silver, 342. Silvering, 343. Gold, 343. Aurous Compounds, 345. 
Auric Compounds, 345. CONTENTS. 
15 
METALS OF GROUP III, 346. 
Group of Earth Metals, 348. 
Aluminium, 348. Aluminium Chloride, 349. Aluminium Oxide, 
351. Aluminates, 352. Alum, 354. Aluminium Silicates, 355. 
Porcelain, 355. Ultramarine, 355. Rare Earth Metals, 356. Scan- 
dium, 356. Yttrium, Lanthanum, Cerium, Didymium, Samarium, 
Ytterbium, Erbium, Terbium, 357. 
Gallium Group, 358. 
Gallium, 358. Indium, 359, Thallium, 360. Thallous Compounds, 
361. Thallic Compounds, 362. 
Tin and Lead, 362. 
Tin, 364. Stannous Compounds, 365. Stannic Compounds, 366. 
Stannates, 367. Lead, 367. Lead Oxide, 369. Plumbic Acid, 370. 
Galenite, 371. 
Bismuth, 372. Bismuthic Acid, 373. 
Chromium Group, 373. 
Chromium, 375. Chromous Compounds, Chromic Compounds, 376. 
Chromates, 378. Chromium Oxy-chloride, 381. Molybdenum, 382. 
Tungsten, 384. Uranium, 385. 
Manganese, 386. 
Forms of Combination, 387. Manganous Compounds, 388. Man- 
ganic Compounds, 389. The Acids of Manganese, 391. 
METALS OF GROUP VIII, 392. 
Iron Group, 394. 
Iron, 395. Metallurgy of Iron, 396. Ferrous Compounds, 398. Ferric 
Compounds, 400. Ferric Acid Compounds, 401. Cyanogen Com- 
pounds, 402. Cobalt, 404. Cobalt-amine Compounds, 406. Cobalt- 
cyanogen Compounds, 406. Nickel, 407. 
Platinum Metals, 408. 
Ruthenium and Osmium, 410. Rhodium and Iridium, 411. Palla- 
dium, 411. Platinum, 413. 
Spectrum Analysis, 416. Periodicity of the Spectrum Lines, 421. A TEXT-BOOK 
OF 
INORGANIC CHEMISTRY. 
INTRODUCTION. 
The study of Nature reveals an endless multitude of objects 
or bodies. That which forms the basis of the latter, strongly 
characterized by extent and weight, we designate substance 
or matter. The investigation of the internal and external 
structure of bodies, their classification according to conform- 
able or distinguishing characteristics, constitute the task of 
the descriptive sciences; of mineralogy, of geology, of de- 
scriptive botany and zoology, of anatomy, etc. 
A closer scrutiny of natural objects discloses the fact that 
they in time succumb to many more or less serious altera- 
tions or changes. We observe that minerals form, crystal- 
lize, or disintegrate and crumble to pieces; that plants and 
animals spring up, grow, and then fall into decay and de- 
composition. Such changes in the condition of bodies occur- 
ring with time are entitled phenomena. The investigation of 
these during their progressiont 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. 
Like every other classification or definition, the division of the natural 
sciences into speculative and descriptive is not strictly correct. It does 
not completely cover the nature of the phenomena. We approximate the 
actual facts more closely by designating the natural sciences as general and 
special. The general sciences, mechanics, physics, and chemistry, occupy 
themselves with the study of the general properties and transformations 
of bodies, regardless of the external form, and deal chiefly with their sub- 
stance only. Special branches—like botany and zoology— consider distinct 18 
INORGANIC CHEMISTRY. 
classes of bodies, first, in reference to their form (morphology, etc.), and 
afterwards in relation to their transformations and alterations. Physi- 
ology of animals and plants, and geology, investigate the physical and 
chemical phenomena of particular classes of bodies, and are, therefore, 
speculative sciences. On the other hand, chemistry is also a descriptive 
science, inasmuch as it considers the external properties of chemical sub- 
stances. 
Although no abrupt boundaries are presented in nature, but 
gradual transitions and intermediate steps throughout, two 
tolerably distinct classes of phenomena may be observed. 
Some changes in the condition of bodies are only superficial 
(external), and are not accompanied by material alteration in 
substance. Thus heat converts water into steam, which upon 
subsequent cooling is again condensed to water, and at lower 
temperatures becomes ice. In these three conditions the solid, 
liquid, and gaseous, the substance or the matter of water or 
ice is unchanged ; only the separation and the motion of the 
smallest particles—their states of aggregation—are different. 
If we rub a glass rod with a piece of cloth, the glass acquires 
the property of attracting light objects, e.g., particles of paper. 
It becomes electrified. An iron rod allowed to remain sus- 
pended vertically for some time slowly acquires the power of 
attracting small pieces of iron. Through the earth’s magnet- 
ism it has become magnetic. In both instances the glass and 
iron receive new properties; in all other respects, in their 
external and internal form or condition, they have suffered 
no perceptible alteration ; the glass is glass, and the iron re- 
mains iron. All such changes in the condition of bodies, 
unaccompanied by any real alteration in substance, are known 
as physical phenomena. 
Let us turn our attention now to the consideration of another 
class of phenomena. It is well known that ordinary iron 
undergoes a change, which we term rusting; i.e., it is trans-i 
formed into a brown substance which is entirely different from 
iron. On mixing finely divided copper filings with flowers; 
of sulphur (pulverulent sulphur) there results an apparently! 
uniform, grayish-green powder. If this be examined, how- 
ever, under a magnifying glass, we can very plainly distinguish 
the red metallic copper particles in it from the yellow of 
sulphur ; by treating with water, the specifically lighter sul- 
phur particles can easily be separated from those of the copper. 
Hence this powder represents nothing more than a mechanical 
mixture. If, however, this mixture be heated, e.g., in a glass 
test-tube, it will commence to glow, and on cooling, a black. INTRODUCTION' 
fused mass remains, which differs in all respects from copper 
and sulphur, and even under the strongest microscope does 
not reveal the slightest trace of the latter, and elutriation 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 been pro- 
duced, and is named copper sulphide. Mixtures of sulphur 
with iron or with other metals act in a similar manner; and 
the resulting bodies are known as sulphides. 
Such mutual action of different bodies occurs not only 
under the influence of heat, but frequently at ordinary tem- 
peratures. If, e.g., mercury and sulphur are rubbed con- 
tinuously in a mortar, there is produced a uniform, black 
compound called mercury sulphide. The action of gaseous 
chlorine upon various metals is quite energetic. When finely 
divided antimony is shaken into a flask filled with yellow 
chlorine gas, flame is produced ; each antimony particle burns 
in the chlorine with a bright white light. The product of 
this action of solid metallic antimony and gaseous yellow 
chlorine is a colorless, oily liquid, known as antimony chlo- 
ride. Such occurrences, therefore, in which a complete and 
entire alteration takes place in the bodmsentering the reaction, 
are termed chemical phenomena. Chemistry, then, is that de- 
partment of natural science which occupies itself with the study 
of those phenomena in nhich 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. Tne opposite may occur, con- 
sisting 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 de- 
posited upon the upper, cooler portions of the tube, globules 
of mercury. From this we observe that on heating solid red 
mercuric oxide two different bodies arise: gaseous oxygen 
and liquid mercury. We conclude, then, that mercuric oxide 
holds in itself, or consists of, two constituents—oxygen and 
mercury. This conclusion, arrived at by decomposition, or 
analysis, may be readily verified by combination or synthesis. 
It is only necessary to heat mercury for some time, at a some- 
what lower temperature than in the preceding experiment, in 
an atmosphere of oxygen, to have it absorb the latter and 
yield the compound we first used—red mercuric oxide. The 20 
INORGANIC CHEMISTRY. 
direct decomposition of a compound body into its constituents 
by mere heat does not often happen. Generally, the coopera- 
tion of another substance is required, which will combine with 
one of the constituents and set the other free. In this man- 
ner we can, for example, effect the decomposition of the pre- 
viously synthesized mercury sulphide, viz. : by heating it with 
iron-filings; the iron unites with the sulphur of the mercury 
sulphide, to form iron sulphide, while the mercury is set 
free. 
If, in a similar manner, natural objects be decomposed, 
bodies or substances are finally reached which have, withstood 
all attempts to bring about their division into further constitu- 
ents, and which cannot be formed by the union of others. 
Such substances are chemical elements ; they cannot be con- 
verted into each other, but constitute, as it were, the limit of 
chemical change. Their number, at present, is about 66 ; 
some have been only recently discovered. To them belong 
all the metals, of which iron, copper, lead, silver, and gold 
are examples. Other elements do not possess a metallic ap- 
pearance, and are known as metalloids ; i.e., they are substances 
resembling (as regards their further indivisib'lity) 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. 
When the elements unite with each other in smaller or larger 
numbers they produce the compound bodies known to us. 
Water is a compound of two gaseous elements—hydrogen and 
oxygen ; common salt consists of the metal sodium and the 
gas chlorine. The elements make up not only our own earth, 
but the heavenly bodies are composed of them, at least as far 
as has been proved by spectrum analysis. 
If the quantities by weight of substances entering into a 
chemical change be determined, we notice that in all trans- 
positions, in the decomposition of a compound into its con- 
stituents, and in the union of the elements to form compound 
bodies, loss in weight never occurs. The weight of the re- 
sulting 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 
THE PRINCIPLE OF THE INDESTRUCTIBILITY OF MATTER. INTRODUCTION. 
conclusively that the growth of plants occurs only in conse- 
quence of the absorption of substance from the earth and at- 
mosphere. 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 of com- 
bustible 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 dis- 
appear, and is not produced anew. This fundamental truth is 
called the principle of the indestructibility of substance {matter). 
Lavoisier, in the eighteenth century, first established it by 
convincing experiments. Combined 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. These names, however, only represent kinds 
of phenomena, without explaining their true nature. Of the 
nature of some of these forces we know, positively, that they 
consist of various modes of motion of portions of matter. 
In the case of mechanical force it is obvious that it depends 
solely upon the motion of masses; but other forces are nothing 
more than modes of motion. The phenomena of light are 
explained by the very rapid movements of the smallest parti- 
cles, and these act upon the eye through the aid of a gaseous 
medium—ether. The phenomena of heat are due to the less 22 
INORGANIC CHEMISTRY. 
rapid motion of weighable portions of matter which affect our 
sense of feeling. Accurate physical investigations have estab- 
lished that the different forces or modes of motion can never 
be destroyed, but only transferred from one body to others, 
and changed from one kind to another. The movements or 
vibrations of one variety pass into those of another. For ex- 
ample, a discharged bullet is heated by coming in contact with 
any obstruction in its course; the visible motion of the entire 
mass in this instance is transformed into the invisible motions 
of the smallest particles, and appears as heat. The heat mo- 
tions can, on the other hand, be again changed into mechani- 
cal motion (molecular motion), or into light, magnetism, or 
electricity. 
In all these transformations of the different modes of motion 
into one another, we observe a perfect equivalence of their 
quantity. If a mass motion, whose quantity is designated as 
mechanical work, can produce a certain degree of heat, so vice 
versa, the latter can perform the same mechanical work (the 
mechanical equivalent of heat, light, electricity). Upon this 
equivalence of transformation rests the principle of the conserva- 
tion of force or energy, according to which the various forces or 
motionsof matter a an neither beannihilated nor produced anew. 
This principle, forming one of the most important corner-stones 
of natural science, was first sharply defined by the speculative 
observations of Dr. J. R. Mayer, of Heilbronn, in 1842, and 
since then has been repeatedly confirmed by experiment.' 
The most recent advance in physics has led to the negation of the ob- 
jective existence of all abstract physical forces. Not considering the 
phenomena of electricity and those of chemical affinity—the reduc- 
tion of which to forms of motion is clearly foreseen, and not to be 
doubted—the only remaining, enigmatical force is that of attraction 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 “ absurdumit 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 trans- 
ference of the gravitation into material bodies, further contradicts the prin- 
ciple of conservation of energy, as gravity is neither transferred nor ex- 
hausted—whether it be through the approach of bodies, whereby the force 
always increases—or by planet movement, in which the centrifugal com- 
ponent is constantly 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 canno.t proceed. 
If we desire to make a preliminary presentation upon these relations, 
the following would be the simplest and most probable: Space is filled INTRODUCTION. 
23 
by the smallest possible material particles, but as they are all alike, they 
do not possess gravity, and are found in constant transferable motion— 
ether substance. By 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-ag- 
gregate, a second appear in space, an effort to approach each other pro- 
duced 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 energy of place are removed. A much 
clearer and more distinct presentation and confirmation of these repre- 
sentations, especially as regards the nature of forces, may be 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 mode 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 kind 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 ab- 
sorbed, disappears as such, and is transformed into chemical 
energy. Thus, for instance, in the union of i kilogram of 
hydrogen with 8 kilograms of oxygen a quantity of heat is 
liberated which can perform a mechanical work equal to 
34.462x423.5 = 14,629,000 kilogrammeters. In the decom- 
position, on the other hand, of 9 kilos of water into hydro- 
gen and oxygen, the same force or quantity of heat is neces- 
sary. 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 con- 
trary, with difficulty, or not at all. The reason for this de- 
portment is to us entirely unknown, but it is in all probability 
due to the different form and mode of motion of the smallest 
particles of matter. We designate it by 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, gen- 
erally, the bodies least alike chemically unite with each other 
most readily. 24 
INORGANIC CHEMISTRY. 
The main condition in the chemical action of bodies con- 
sists in their immediate contact. While, in physical chang-s, 
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 indispensably intimate 
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, accom- 
panied with effervescence. As shown in this case, the requi- 
site 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 fluida. 
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 con- 
ducted through it, chemical action at once takes place accom- 
panied by an explosion. 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 comprehended that with the increase of temperature 
the chemical attraction of the various substances is in condi 
tion to overcome the cohesive force. Light and electricity 
behave like heat. On the other 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 it so that decomposition or alteration of 
the compound body will result at a certain temperature. 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 decomposi- 
tion. Light and electricity do the same. 
In addition to these circumstances we should remember that 
both the compound bodies and elements consist of several 
CONDITIONS OF CHEMICAL ACTION. INTRODUCTION. 
25 
atoms combined with each other, and that this union must 
first be broken to render possible the action of other particles. 
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 ref- 
erence will be made to individual generalizations. The quan- 
titative relations, on the other hand, in the action of bodies 
on each other, in the combination of the elements, are thor- 
oughly investigated, and lead to very important laws, specu- 
lations, and theories, which constitute the main portion of the 
scientific chemistry of to-day. 
As briefly represented in the preceding, two different factors are to be 
considered in every chemical phenomenon. First, there 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 unite chemically to produce water, there occurs, in 
addition, a considerable development of heat, by which the chemical 
energy and affinity become evident (p. 64). Although the physical cause 
of the latter, formerly identified with electrical differences, is as yet unex- 
plained, 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 is 
their affinity ; so that the former may serve as an approximate measure of 
the latter. From this follows the general proposition that the qualitative 
alterations occurring in the reaction of bodies invariably pursue the di- 
rection in which the most heat is disengaged. This fundamental propo- 
sition of the greatest development of heat (Bertholet) may be considered a 
special case of Clausius’s proposition of entropy or energy, according to 
which every form of energy tends to pass into heat-motion. Since, how- 
ever, chemical substances do not consist of free atoms, but of atomic ag- 
gregations (molecules) which require tearing asunder in order that their 
components may act upon one another, and because the stability of com- 
pounds, their decomposition (disgregation) by heat and solvents, do not 
always accord with their heat of formation, it becomes necessary to formu- 
late the principle mentioned above to read as follows : In chemical changes 
the tendency is to produce those bodies in whose formation the greatest de- 
velopment of heat (degradation of energy) occurs. This proposition is quite 
important for the illustration and explanation of chemical processes. It 
will be more fully developed, in its most interesting applications, under 
the different groups. 
CHEMICAL SYMBOLS AND FORMULAS. 
For simplicity and convenience the elements are represented 
by the first letters of their names, derived either from the 
Latin or Greek. Hydrogen is represented by the letter H, 
from the word hydrogenium; nitrogen by N, from nitrogen- 
ium. When several elements happen to have the same letter 26 
INORGANIC CHEMISTRY. 
there is added to the capital a second, small letter; thus, Na 
represents natrium; Ni, nickel; Hg, mercury (hydrargyrum), 
etc. The subjoined table comprises all the elements known 
at present with certainty (66), together with their chemical 
symbols and atomic weights. The latter have been deter- 
mined with more or less accuracy. The number of decimal 
places corresponds approximately to the accuracy of the atomic 
numbers; those without decimals are, therefore, less accu- 
rately determined: 
Elements. 
Symbol. 
Atomic 
Weight. 
Elements. 
'o 
•O 
s 
in 
Atomic 
Weight. 
• 
Aluminium 
A1 
27.O 
Mercury 
Hg 
199.8 
Antimony (Stibium). 
Sb 
II9.6 
Molybdenum 
Mo 
95-8 
Arsenic 
As 
74-Q 
Nitrogen 
N 
I4.OI 
Barium 
Ba 
136.8 
Nickel 
Ni 
58.6 
Berv Ilium 
Be 
9.I 
Niobium 
Nb 
94 
Bismuth 
Bi 
207 
Osmium 
Os 
195 
Boron 
B 
IO.9 
Oxygen 
O 
15.96 
Bromine 
Br 
79.76 
Palladium 
Pd 
106.2 
Cadmium 
Cd 
111.9 
Phosphorus 
P 
W.q6 
Caesium 
Cs 
I "t2-7 
Platinum 
Pt 
194.3 
Calcium 
Ca 
3Q.Q 
Rhodium 
Rh 
104 
Carbon 
c 
1r-99 
Rubidium 
Rb 
85 
Cerium 
Ce 
141 
Scandium 
Sc 
44 
Cr 
52.4 
Sulphur 
S 
31.08 
Cu 
63.2 
Selenium 
Se 
78.9 
Chlorine 
Cl 
35-37 
Silver (Argentum)... 
Ag 
107.66 
Cobalt 
Co 
58.6 
Silicon 
Si 
28 
Didymium 
Di 
142 
Sodium (Natrium)... 
Na 
22.Q9 
Erbium 
Er 
166 
Strontium 
Sr 
87-3 
Fluorine 
El 
19 
Tantalum 
Ta 
182 
Gallium 
Ga 
69.8 
Tellurium 
Te 
126 
Au 
196.2 
Thallium 
T1 
203.6 
H 
1 
Thorium 
Th 
232 
In 
113.4 
Tin (Stannum) 
Sn 
117.5 
Iodine 
I 
126.54 
Titanium 
Ti 
48 " 
Iridium 
Ir 
192.5 
Tungsten (Wolfram) 
W 
183.6 
Fe 
3 3.9 
Uranium 
Ur 
239 
Kalium (Potassium). 
K 
J J y 
39-03 
Vanadium 
Vd 
.51.2 
La 
138.2 
Ytterbium 
Yb 
173 
Li 
7 
Yttrium 
Y . 
89 
Lead (Plumbum) — 
Pb 
206.4 
Zinc 
Zn 
64.9 
Mg 
239 
Zirconium 
Zr 
90 
Manganese 
Mn 
54-8 
The existence of some other elements, as Decipium, Holmium, Terbium, 
Wasium, Vesbium, Norwegium, and Thulium, is yet doubtful. INTRODUCTION. 
27 
Compounds produced by the union of the elements are rep- 
resented by placing their corresponding symbols together and 
designating these chemical formulas. Common salt, a com- 
pound of sodium and chlorine, is represented by the formula 
NaCl; mercuric oxide, a compound of mercury and oxygen, 
by HgO ; iron sulphide by FeS ; hypochlorous acid, a com- 
pound of chlorine, hydrogen and oxygen, by CIOH. 
Chemical formulas not only express the nature of the ele- 
ments, but also the relative proportions by weight, according 
to which they unite, compared with hydrogen as unity. Thus 
H represents i part by weight of hydrogen ; Cl, 35.37 parts 
by weight of chlorine ; Na, 22.99 Parts by weight of sodium. 
(See table, p. 26). 
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 matter in general, two possibilities ap- 
pear 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 state of physical 
and chemical investigation, so that the atomic constitution of 
matter only has value at present. The inductive derivation 
and establishment of the atomic theory will be given subse- 
quently (see page 66); here we will only state the following 
propositions: Each distinct element consists of similar atoms, 
of like size and similar weight, while atoms of different ele- 
ments possess a different weight. The absolute atomic weights 
are, at present, not determined with sufficient accuracy ; the 
relative weights are referred to the hydrogen atom, which has 
the smallest weight, hence is made equal to 1 (H = 1). The 
chemical union of the atoms produces 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 indicates that in it 22.9 
parts, by weight, of sodium are combined with 35.37 parts of 
chlorine. If several atoms of an element are present in a 
compound, this is denoted by numbers which are attached to 
the symbol of the atom : 
HC1 HaO IIsN CH4 
Hydrochloric acid. Water. Ammonia. Methane. 28 
INORGANIC CHEMISTRY. 
The formula of water (H20) means that its molecule con- 
sists of 2 atoms of hydrogen (2 parts by weight) and 1 atom 
of oxygen (O = 15.96 parts by weight). The formula of sul- 
phuric acid (H2S04) indicates it to be a compound consisting 
of 1 atom of sulphur (31.98 parts), 4 atoms of oxygen 
(4 X 15.96 = 63.84 parts), and 2 atoms of hydrogen (2 X 1 
—2 parts), from which the composition of the acid may be at 
once calculated into per cent., or into any desirable quantity 
by weight. 
Atomic Composition. In per cent. 
Sulphur, S = 31.98, ...... 32.69 
Oxygen, O4 = 63.84, ...... 65.26 
Hydrogen, H2 = 2, . . . . . . 2.05 
H2SO4 = 97.82 . . . . . . 100.00 
The chemical union of bodies is*sho.wn by the sign 4., and 
the resulting products are placed to the right, following the 
= sign : 
HgS -f- Fe = FeS -f- Hg. 
By this* equation of chemical transposition is meant that by 
the union of mercury sulphide (HgS) and iron (Fe), iron sul- 
phide (FeS) and free mercury (Hg) are formed. At the same 
time such equations show the proportions by weight of the 
substances entering into and resulting from the reaction ; the 
weight of the acting substances is equal to that of those re- 
sulting. Therefore every chemical equation is at once an 
expression of the principle of the indestructibility of matter 
(substance). (See p. 21.) 
CRYSTALLOGRAPHY. 
Chemistry occupies itself chiefly with the investigation of 
the chemical alterations of bodies. Its subject is not the latter 
in themselves, in their external properties, but only with ref- 
erence to their material composition, and their genetic rela- 
tions 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 
occurring bodies, but includes also the innumerable substances 
which have been prepared artificially, it becomes the province 
of Inorganography. Pure chemistry considers the physical INTRODUCTION. 
29 
properties only as far as they serve for the external character- 
ization and eventual recognition of the given substances and 
for the deduction of chemical irregularities. The most im- 
portant physical properties,— the state of aggregation, the 
temperature of fusion and boiling, the sptcific gravity, capa- 
city for heat, etc.—are partly treated in Physics, and in part 
will be considered later, in special cases. 
Here, therefore, the morphological characters of the solid 
bodies will receive only a brief consideration. 
The homogeneous solids exhibit either similar properties in 
all their parts, are amorphous, or show differences in certain 
definite directions, giving rise to a crystalline appearance. 
The cause of this deportment lies in the arrangement of the 
smallest particles of substance 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 den- 
sity 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 for- 
mation, 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 numerous, since several 
thousands are known. It is, however, possible to reduce the 
numberless varieties to a few classes or systems, by comparing 
their modes of formation, and by referring their principal 
elements—the planes—to definite axes, /.<?., to directions or 
lines, which are imagined to be so placed through the mid- 
dle point of the crystals that their planes lie symmetrically 
with reference to them. In this manner, by distinguishing the 
various axis-intersections, crystallography arrives at the follow- 
ing 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 therefore the po- 
sition of a crystal face is also determined by its points of in- 
tersection with the three axes, or by the distances from the 
centre of the axis at which the plane (by suitable expansion) 
cuts or intersects the three axes. These distances are termed 30 
INORGANIC CHEMISTRY 
the parameters 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. i). If a plane 
intersect all three axes at equal distances (as in the octahe- 
dron) the parameter ratio is a: a : a; if at unequal distances, 
the ratio is a : a: ma, or a : ma : na; in which case the parameter 
of the first axis is always made = i. If a plane lie parallel 
to an axis intersecting it at infinity, its parameter with refer- 
ence to this axis = infinity (a : a : oo a); if it be parallel to 
two axes, two parameters are = infinity (a: eo a: co a). Hence 
these parameter ratios designate the position of a plane; and 
as all the planes in simple crystal forms are similar, these 
symbols represent the entire simple form, the whole 
of all planes, which are deduced from the same parameter 
ratio. 
In the regular system where all three axes are of equal 
lengths, the parameters m, n (the parameter of the first axis 
equalling i) are simple rational numbers (i, £, 2, 3, 4, etc.). 
In other systems of ciystallization either two or three axes 
are unequal, and do not stand in any rational ratio to their 
lengths (a : b : c), and, therefore, to the parameters of the 
planes (a : mb : nc); but the ratio of the parameter coefficients 
1. m. n., however, in them is, as learned from experience, a 
simple, rational one, like that of the regular system. 
According to the number of axes, five of the crystallographic systems 
are triaxial or trimetric ; for practical reasons \ye assume the presence of 
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 regu- 
lar, the quadratic, and the rhombic; or they include, also, oblique angles 
—in the two clinometric systems, the monoclinic and triclinic. In the 
regular system the three directions of development are at right angles 
and perfectly similar, for which reason the bodies belonging to it refract 
light simply—isometric system. Amorphous substances do the same. The 
remaining five systems with unlike axes are, on the contrary, doubly re- 
fracting, and exhibit simple refraction in but one direction (optically uni- 
axial systems—the quadratic and hexagonal), or, in two directions (opti- 
cally, diaxial systems—the rhombic, mono-, and triclinic). Bravais and 
Sohncke have recently suggested a geometrical derivation of the different 
systems of crystallization; the basis of the same is the molecular consti- 
tution of matter. The authors proceed from the possible groupings of 
points or molecules in space, i.e , from such groupings in which about each 
point the arrangement of the remaining points is the same. In this way it 
is possible to embrace the point positions which can occur in space into seven 
groups, corresponding to the six crystallographic systems. The rhombo- 
hedral forms, which crystallography views as the hemihedral of the hex- 
agonal, then appear as an independent system. INTRODUCTION. 
31 
i. Regular System.—The crystals of this system are 
similarly developed in three directions, and therefore have 
three similar axes, a, of the same length, which intersect each 
other at right angles (Fig. i). If we imagine the faces of the 
crystal arranged in such manner that they intersect the three 
Fig. 1. 
Fig. 2. 
Fig. 3. 
axes at similar or dissimilar distances, we have the following 
seven simple forms for the various possible parameter ratios 
page 30): 
1. The octahedron (Fig. 2), with the parameter ratio a : a : a; 
for which the abbreviated symbol O maybe written. 2. The 
cube, or hexahedron (Fig. 3),with the symbol a : 00 a: co a or 
briefly 00 O 00. 3. The rhombic dodecahedron, twelve sides, 
a: a: 00 a or 00 O (Fig. 4). 4. The trigonal tris-octahedron 
Fig. 4. 
Fig. 5. 
Fig. 6. 
(Fig. 5), a: a: ma or mO ; there are several of these,as m can 
equal f, 2 or 3. 5. The trapezohedron (Fig. 6), a: ma: ma 
or mOm, in which m = 2 (in garnet) or m = 3 (in ammo- 
nium chloride). 6. The tetrahexahedron (Fig. 7), a: ma : 00 a 
or mO 00, in which m= 2, or 3. 7. The hexoctahedron 
(Fig. 8), a: ma: na or mOn, where, e.g., in fluorite, m — 2 
and n — 3. 
These seven simple forms may appear combined on one 
crystal, and we thus obtain crystal combinations. For example, 
Fig. 9 represents a combination of octahedron and cube (on 32 
INORGANIC CHEMISTRY. 
alum); Fig. 10, a combination of octahedron, cube, and do- 
decahedron (galenite). 
Fig. 7. 
Fig. 8. 
Fig. g. 
In addition to these simple forms, appearing with their full 
number of planes, hence termed holohedral, others are found 
having only the half or fourth of the possible forms, hemihe- 
Fig. 10. 
Fig. ii. 
Fig. 12. 
dral or tetartohedraliorms. We may suppose them produced 
by the enlargement of the symmetrically distributed half 
number of faces of the holohedral form, and therefore they 
are marked with the corresponding holo- 
hedral symbol divided by 2. Such hemi- 
hedral forms are: The tetrahedron 
(Fig. 11), resulting from the octahedron, 
and the pentagonal dodecahedron — 
(Fig. 12) derived from the tetrahexa- 
hedron. 
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 equal secondary axes, a, and an unequal, 
longer or shorter, principal axis, c (Fig. 13). The axis 
ratio, therefore, is a : a : c; the ratio of a : c is definite but not 
rational for every body of the system (see page 30), e.g., for 
FlG- !3- INTRODUCTION. 
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 c 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 pyra- 
mids of the second order, in which the two secondary axes 
pass through the middle of the edges. The planes of the latter 
Fig. 14. 
Fig. 15. 
Fig. i6. 
are parallel to one secondary axis, therefore its symbol is 
a: 00a : c, or P 00. Fig. 15 represents a combination of a 
pyramid of the first and second orders. 
In addition to the above pyramids, others may occur upon the same 
crystal, which intersect the principal axis at the distance me. In this case 
m is also a simple rational number, e.g.t 3, 2, etc. The symbol of 
such secondary pyramids of first order, is a : a : me = mP : that of the 
second order, a: do a: me = mPoo . The coefficient of the secondary 
axis is always written before P. 
If the pyramid'planes intersect the principal axis at infin- 
ity, the quadratic prism results, and, indeed, either a prism of 
the first order, with the symbol a : a : ooc = Poo ; or a prism 
of second ofder, a : oo a : oo c = oo Pco . Fig. 16 represents 
a combination of the quadratic prism with the quadratic pyra- 
mid, as seen on zircon and potassium phosphate. With the 
parameter ratio of the planes, a : ma : nc, we get the ditetrag- 
onal pyramid nPm. Its corresponding ditetragonal prism has 
the symbol a : ma : op c = oo Pm. 
Different hemihedral "forms are possible in this system, of 
' . p 
which may be mentioned the tetragonal sphenoid t-, corre- 
sponding to the tetrahedron. Chalcopyrite, tin, potassium 
phosphate, etc., crystallize in this system." 
* The principal (ground) form is that to which the remaining forms of 
the same mineral may be most readily referred. 34 
INORGANIC CHEMISTRY. 
3. Hexagonal System.—The forms of this system, like 
those of the preceding, exhibit one peculiarly striking direc- 
tion 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 equal secondary axes (Fig. 17), 
intersecting each other at 6o°. The'principal axis is at right 
angles to these. - The axis ratio 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. 
The fundamental (ground) form of the system is the hexag- 
onal pyramid a : a.: (00 a) : c = P, from which is derived the 
hexagonal prism a : a : (00 a) : 00 c '= 00 P ; and, indeed, as in 
the quadratic system, there are pyramids and prisms of first 
Fig. 17. 
Fig 18. 
Fig. 19. 
and second order—the latter with the symbol a: 2a: (2a) : c = 
P2, and a : 2a : (2a) : 00 c = °o P2. Further, other pyramids can 
occur, iatersecting.the principal axis, at the distance me; theii 
symbol is mP and mP2. Fig. 18 represents the combination 
of pyramids and prisms found on apatite. With the common 
parameter ratio a : na : (fa) : me (in which the parameter o) 
the third secondary axis, ra, is always given as in all-othei 
hexagonal forms, by the parameters of the first two secondary 
axes), result the dihexagonal pyramid, mPn, and the dihex- 
agonal prism, 00 Pn. 
mP , 
The rhombohedra iy=±mR (Fig. 19) are the hemi- 
hedral forms of the pyramids mP, produced by the growth ol 
alternate faces. Another important hemihedral form is tht 
scalenohedron ±—y- derived from the dihexagonal pyramid 
It is a remarkable fact that the hemihedral forms of the hex- 
agonal system occur much more frequently in nature in num- 
berless combinations (especially in calcite), and they are, there 
fore, sometimes treated as a separate system. INTRODUCTION. 
35 
4 Rhombic System.—Three axes of unequal length, a. 
b. c., at right angles to each other. Any one, as c, is chosen 
as principal axis, and of the secondary axes, the shorter a is 
Fig 20. 
Fig. 21. 
Fig. 22. 
designated as the brachydiagonal, the longer b as macrodiag- 
onal (Fig. 20). The axis ratio a : b : c is definite for every 
substance, eg., for sulphur, 0.811 : 1 : .0899. 
The principal forms of the system are the rhombic pyramid 
a : b : c = P (Fig. 21), the rhombic prism a : b : co c — 00 P, 
and the domes—brachydiagonal 00 a : b : c = P 00 and the 
macrodiagonal a : 00 b : c = p 00 , consisting of two pairs of 
planes. Fig. 22 shows a combination of two pyramids with the 
brachydiagonal dome. Sulphur (native and that crystallized 
from carbon disulphide), potassium nitrate, aragonite and 
barite belong to 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 is at right angle to b and to c, but 
c is oblique to b at the angle ft. The 
crystals, of this system are mostly devel- 
oped according to an axis oblique to 
another (c to b), and hence one of these 
is chosen as principal axis c. Of the two 
secondaries, one, b is called the clino- 
diagonal, and a the otthodiagonal. The 
axis ratio is always definite for every 
substance; eg., for ferrous sulphate 
1.1704: 1 : 1.5312 with the angle ft = 
76° 33', The monoclinic pyramid is the principal form 
a b : c = P, consisting of two hemipyramids, + P and — P, 
(Fig. 24). It corresponds to the monoclinic prism 00 P. 
Fig 23. 36 
INORGANIC CHEMISTRY. 
Sulphur (fused), soda, borax, di-sodium phosphate (HNa2PO,± 
+ 12 H20), Glauber’s salt, and orthoclase, crystallize in this 
system. 
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 
Fig. 24. 
Fig. 25. 
Fig. 26. 
very Complicated, since the triclinic pyramids P must be con- 
sidered as consisting of four quarter pyramids, and the triclinic 
prism oo P (Fig. 25), of two hemiprisms. Potassium dichro- 
mate, albite, boracic acid, and copper sulphate crystallize in 
this system. Fig. 26 represents one of. the common forms of 
the latter. 
Crystals found in nature have grown and therefore rarely 
occur so regularly developed as represented in the preceding 
drawings. They are usually developed more or less in the 
direction of single axes, and on that account 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 remains 
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,” we cal- 
culate the axis ratio from the angles. 
Substances crystallizing in two or three different systems are 
said to be dimorphous, or trimorphous.(see sulphur). Different 
substances’ crystallizing 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. 20). 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; their oxygen deriva- 
tives 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 com- 
pounds 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-form- 
ing nature. A sharp line of difference between metals and 
metalloids does not exist. There are elements which in their 
external appearance resemble metals, while in a chemical res- 
pe.ct they deport themselves throughout as metalloid's, and vice 
versa. Thus hydrogen, a gaseous element, is like the metals 
in its entire chemical character, while metallic antimony 
arranges itself with the metalloids. 
It- is therefore advisable to divide the elements into separate 
natural groups, based upon their chemical-analogies. 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 shall treat of 
the periodic system more at length ; it forms the basis of this 
text-book, and in accordance with this doctrine we consider 
the elements in single natural groups of similar chemical 
deportment. The first of these groups, comprising almost 
all the so-called non-metals' and some metals, are the fol- 
lowing : INORGANIC CHEMISTRY. 
Fluorine 
Chlorine 
Bromine 
Iodine 
Oxygen 
Sulphur 
Selenium 
Tellurium 
Nitrogen 
Phosphorus 
Arsenic 
Antimony 
Bismuth 
Carbon 
Silicon 
Tin 
Boron 
Hydrogen does not belong to either of these groups ; unit- 
ing the metallic and non-metallic characters in itself, it repre- 
sents, as it were, the type of all elements, and therefore it will 
receive first attention. Boron occupies an isolated position. 
It has been classed with the metalloids, but differs somewhat 
from them in chemical deportment. It forms the transition 
to the metallic elements, beryllium and aluminium. 
HYDROGEN. 
H = 1. H2 = 2. 
Hydrogen (Hydrogenium), a gaseous body, is rarely found 
in a free condition upon the earth’s surface, as it combines 
readily with the oxygen of the air. It is present in consider- 
able quantity in 
the photosphere of 
the sun and fixed 
stars. In combi- 
nation, we fipd 
it principally as 
water and in sub- 
stances of vegeta- 
bleand animal ori- 
gin. Paracelsus first 
discovered this ele- 
ment in the six- 
teenth century, 
* and called it in- 
. flammable air. In 
1781' Wafts and 
Cavendish showed 
that water resulted 
from the combus- 
tion of hydrogen in the air. In 1783 Lavoisier proved that 
hydrogen was a constituent of water—a chemical compound 
of the elements hydrogen and oxygen. ‘ • 
Preparation.—It may be readily obtained from water, a 
compound of hydrogen and oxygen. The-decomposition of 
Fig. 27. HYDROGEN. 
39 
the same by the removal of oxygen can be effected by some 
metals, like Na and K, at the ordinary temperature. Both 
metals act very energetically upon it, liberating gaseous hy- 
drogen. 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 dis- 
engaged, displace the water and collect in the upper part of the 
cylinder. The reaction occurring between the sodium and 
water is expressed by the following chemical equation: 
H20 + Na = NaOH -f H. 
Water. 
Sodium. 
Hydrogen. 
The compound NaOH, known as sodium hydroxide, 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, exposed to a red 
Fig. 28. 
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 -f H2S04 . = ZnS04 + 2H. 
Sulphuric acid. 
Zinc sulphate. 40 
INORGANIC CHEMISTRY/ 
Place granulated zinc (obtained by dropping molten zinc 
into water) in a double-necked flask (Fig. 28), and intro- 
duce sulphuric acid (diluted with about 3 vols. of HaO) 
through the funnel tube, b. The liberation of gas begins 
immediately, and the hydrogen, escaping through the exit 
tube,/, is collected as previously described. 
The hydrogen thus formed has a faint odor due to a 
slight admixture of foreign substances. Pure hydrogen may be 
obtained by heating potassium’formate with potassium hy- 
drate : 
CH02K + . KOH ' = • K2COs -f 2li, 
or from palladium hydride. 
Purifying and Drying of Gases.—To free gases of the substances me- 
chanically carried along during their disengagement, it is best to conduct 
them through variously constructed wash-bottles, filled with water or 
‘ - * liquids that will absorb the im- 
purities. Ordinarily the so-called 
Woulff’s bottles are employed 
(compare Figs. 36 and 42). The 
open tube,1 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 ate always moist, as they 
contain aqueous vapor. To remove 
this conduct them through vessels 
or tubes filled with hygroscopic sub- 
stances (see Fig. 34). - Calcium 
chloride, burnt lime, sulphuric acid, 
etc., are used for this purpose. 
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 con- 
trol the generation of the gas, we 
have recourse to different forms of 
apparatus. One of the most practi- 
cable of these is that of Kipp. It 
consists of two glass spheres, d and b, Fig. 29, in the upper opening of 
which there is a third sphere, c, fitting air-tight and provided with an elon- 
gated tube.. 'It serves as a funnel. Granulated zinc is placed in the mid- 
dle sphere through the tubuluie e, and dilute sulphuric acid is poured into 
the spherical funnel, which first fills d, then ascends to b, where it comes in 
contact with the metal; at once the evolution of hydrogen commences and 
the gas escapes through e. Upon closing the stop-cock of the tube fixed 
Fig. 29. HYDROGEN. 
41 
in e, the hydrogen that is set free presses the sulphuric acid out of b, 
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 com- 
mences anew. The vessel a contains water to wash the escaping hy- 
drogen. 
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, are 
Fig. 30. 
connected by a rubber tube. The bottle A is filled -with granulated zinc, 
and B with dilute sulphuric acid. The cock JR closes A. When this is 
opened the sulphuric acid flows from B\.o A, to the zinc, and the evolution 
of gas commences. On closing the stop-cock the hydrogen 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 preserv e gases. 
Tn 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 tube, f allows the operator to 
observe the height of the water-level. When the lower cylinder is filled 
with water close a and e (the last traces of air can be removed by opening 
b\. To fill the gasometer with gas, remove the cover of the side tubulure 
</, and introduce the tube from which the gas is escaping. The latter 
rushes up into the cylinder* while the water flows out the tubulure. 42 
INORGANIC CHEMISTRY. 
When the water is displaced by the gas, close d, after filling the upper 
cylinder, and then, if desired, open a, and the gas can be set free, either 
by e or b. 
In addition to the gasometer described, various other forms are em- 
ployed ; gas-bags are very well adapted for preserving gases. 
Physical Properties.—Hydrogen is a colorless, odorless, and 
tasteless gas. Its ability to refract light and conduct heat is 
greater than that of all 
other gases. This is in 
accord with its metal- 
lic nature. By cooling 
(— 140° C.) and power- 
ful pressure (600 atmos- 
pheres), it is condensed 
toasteel-blue.non trans- 
parent liquid, which 
upon evaporation even 
becomes solid ; conse- 
quently liquid hydro- 
gen resembles a molten 
metal, or liquid mer- 
cury at ordinary tem- 
peratures. The density 
of hydrogen equals 
about 0.03 (water = 1). 
Condensation of Gases.— 
Until recently hydrogen and 
several other gases (oxygen, 
nitrogen, carbon monoxide, 
nitric oxide, methane) were 
considered non-coercible 
gases (permanent gases), 
as they had not been con- 
densed by the strongest pressure lover 2000 atmospheres). This supposed 
non-condensability was due to the fact that a general property of gases,called 
by Andrews their critical te?nperature, had been overlooked. As first 
observed with carbon dioxide (see this) all gases show, indeed, a definite 
temperature, beyond which they cannot be condensed by any pressure. 
Conversely, all liquids, if raised above this temperature, are even under 
the greatest pressure, transformed into gases (absolute boiling tempera- 
ture of Mendelejejf). Hence, in the condensation of gases, the corres- 
ponding lower temperature must be applied along with pressure. In 
fact of late years all the so-called permanent gases have been condensed 
by this means (Pictet and Cailletet). The lower temperatures are attained 
by the rapid evaporation of liquid carbon dioxide (— 130°), or hypo- 
nitrous oxide (—140°) with aid of air-pumps; or the strongly com- 
Fig. 31 HYDROGEN, 
43 
pressed gas is permitted to expand rapidly, whereby so much heat is 
absorbed through the evaporation that the remainder of the gas con- 
denses. 
Like all difficultly coercible gases, hydrogen is but slightly 
soluble in water, 100 volumes dissolving 1.9 volumes H. It 
is the lightest of all gases, being 14.43 times lighter than air. 
Its specific gravity, compared with air as unity, is = 
0.06930. In many respects it is more convenient to com- 
pare the specific gravity of gases with H as unity, espe- 
cially as the composition of air, and therefore its weight by 
volume, is not constant. If the specific gravity of gases com- 
pared with H = 1 be represented by A, and the specific 
gravity compared with air= 1 by D, then A-D* 14.43and 
A 
D — A cubic decimetre (= 1 litre) of hydrogen, 
according to Regnault’s determinations, weighs 0.08958 
grams at o° and an atmospheric pressure of 760 millimeters 
(in the 490 of latitude), a litre of air, which is 14.43 times 
heavier, weighs, therefore, 1.2926 grams. 
That hydrogen is lighter than air is shown by a balloon of 
collodion or gum filled with the former rising in the latter; 
this can also be seen in soap 
bubbles filled with hydrogen. 
In consequence of its lev- 
ity, hydrogen may be col- 
lected in inverted vessels 
(opening turned down) by 
replacing the air, and can also 
be poured from one cylinder 
into another,as is represented 
in Fig. 32. The hydrogen 
flows from the inclined cyl- 
inder into the one held vertically and filled with air, which it 
expels. Such a separation of gases, based on their varying 
specific gravity, is only temporary, as they soon mingle with 
each other by diffusion. By virtue of its levity and mo- 
bility, hydrogen penetrates porous bodies with ease, and dif- 
fuses through both animal and vegetable membranes, as well 
as through gutta-percha. Metals, e.g., iron, platinum, palla- 
dium, permit a free passage to hydrogen, when they are raised 
to a red heat. They are impenetrable to other gases. This 
behavior is in part probably dependent upon the chemical 
attraction of these metals for hydrogen. 
Fig. 32. 44 
INORGANIC CHEMISTRY. 
Chemical Properties.—Hydrogen is characterized by its 
ability to burn in the air, when it combines with the oxygen 
of the latter and forms water; hence 
its name hydrogenium (from vdwp, 
water,. and pvaw, I produce). Its 
flame is faint blue, and almost non- 
luminous, but possesses a very high 
temperature. When a mixture of 
hydrogen and air is ignited a violent 
explosion ensues; therefore, before 
bringing a light in the vicinity of hy- 
drogen 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 inverted cylinder con- 
taining an atmosphere of the gas (Fig. 33), the latter will 
ignite at the mouth of the vessel, but the candle will be extin- 
guished. 
Water is the product of the combustion of hydrogen in the 
air. It is a chemical compound containing hydrogen and oxy- 
Fig. 33. 
Fig. 34. 
gen. To render the formation of it visible, by the combus- 
tion of hydrogen, the flame of the latter is made to burn under HYDROGEN. 
45 
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 
sulphuric acid or a tube filled with calcium chloride, to absorb 
all moisture. 
The union of hydrogen with oxygen to form water occurs 
only at high temperatures, at a red heat, in contact with 
a flame, or by the passage of an electric spark through the 
mixture. The combination can be effected at ordinary tem- 
peratures with the aid of platinum sponge; the latter consists 
of finely divided metal, obtained by the ignition of 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 power of the metal 
to condense hydrogen and oxygen upon its surface, and 
thereby increase their ability to unite. 
Upon this behavior depends the action of the so-called Dcebereiner 
Lamp (Fig. 35). This is a continuous hydrogen generator. The outer glass 
cylinder, c, is filled with dilute sul- 
phuric acid, into which projects 
the pear-shaped vessel b. This is 
open below and above, and is pro- 
vided with a stop-cock, e, affording 
communication with the air; in it a 
piece of* zinc is suspended by a wire. 
On opening the stop-cock the sul- 
phuric acid presses from the outer 
cylinder a into b and meets the zinc— 
when the liberation of hydrogen be- 
gins. The stop-cock directs the gas 
upon the support, f, in which is fixed 
some platinum sponge that effects the 
ignition. 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. 
The absorption of hydrogen 
by the metal palladium is 
very characteristic. As already 
known, water is so decomposed 
by the electric current that hydrogen separates at the 
electro-negative pole and oxygen at the electro-positive. 
Now, if a piece of palladium, in sheet or wire form, be 
attached to the electro-negative pole, the disengagement of 
hydrogen does not occur, because it is absorbed by the palla- 
dium, in a quantity over nine hundred times the volume of the 
Fig. 35. 46 
INORGANIC CHEMISTRY. 
latter. Palladium also absorbs hydrogen when it is heated 
to ioo°. 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 hydrogen, Pd2H, therefore, 
conducts itself like an alloy of two metals. From the specific 
gravity of the compound (according to Graham), the specific 
gravity of the condensed hydrogen is readily found to be o 62 
(water =1), and is, therefore, somewhat heavier than the 
metal lithium. The metals potassium and sodium absorb hy- 
drogen when heated from 200° to 400°, forming the alloys 
(Na2H and K2H) in which the density of hydrogen is again 
equal to 0.62. From these facts we could conclude that hy- 
drogen possesses a metallic character, and if condensed it 
would form a metallic,mercury-like liquid—aconclusion which 
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, it is able to expel 
different metals from their salt solutions, the same as other 
metals do. When palladium hydride is heated it suffers de- 
composition, hydrogen escaping, just as the mercury is ex- 
pelled from mercury alloys. 
HALOGEN GROUP. 
To this group belong chlorine, bromine, iodine, and fluo- 
rine. 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. 
i. CHLORINE. 
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, which is composed 
of chlorine and sodium. The Swedish chemist, Scheele, dis- 
covered chlorine in 1774. Its elementary character was first 
established by Gay-Lussac and Thenard in France (1809), and 
by Davy in England (1810). 
Preparation.—To obtain free Ghlorine heat a mixture of 
black oxide of manganese (MnO>2) and hydrochloric acid in a 
Cl = 35-37- Cl2 = 70.7. HALOGENS. 
47 
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 that is carried along mechani- 
cally by passing it through water in a three-necked Woulff’s 
Fig. 36. 
bottle, and then collecting it over water. The reaction which 
occurs above is indicated in the following equation : 
MnOa + 4HCI = MnC]2 + Cl2 + 2H2Q. 
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 6 of water) is employed: 
Mn02 + 2NaCl + 2H2S04 = MnS04 + Na2S04 + Cl2 + 2H20. 
Manganese 
dioxide. 
Sodium 
chloride. 
Sulphuric 
acid. 
This reaction comprises two phases: First, the sodium chloride (NaCl) 
is decomposed by the sulphuric acid, yielding sodium sulphate and hydro- 
chloric acid: * 
2NaCl + H4S04 = Na2SQ4 + 2HCI. 
The latter acid then acts, together with a new portion of sulphuric acid, 
upon the manganese dioxide ; 
Mn02 +• H,S04 + 2HC1 = MnS04 + 2H20 -f Cl2. 48 
INORGANIC CHEMISTRY. 
The second method is more advantageous for laboratory purposes ; the 
first, however, is preferred in practice, as it is cheaper. 
The resulting manganous chloride (MnCL) is converted by the Process 
of Weldon into manganese peroxide (see this). Technically, chlorine is 
also obtained by the Process of Deaccrn, by conducting HC1 mixed with air 
over strongly ignited 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. It cannot be collected over mercury, as it readily com- 
bines with the latter. When perfectly dry chlorine is sought, conduct the 
liberated gas through Woulff’s bottles containing sulphuric acid, to absorb 
the moisture, then collect in an empty upright flask (compare Fig. 44, p. 
61). As chlorine is so much heavier than air it will displace the latter. 
Physical Properties. — Chlorine is a yellowish-green gas 
(from ylwpds), with a penetrating, suffocating odor.' Its spe- 
cific gravity compared with hydrogen (i) is 35.37 ; with air 
-2C. -27 
( — 1) it is = 2.45. At 150 C., and a pressure of 4 
atmospheres (at —40° C., under the ordinary pressure) it con- 
denses to a yellow liquid, boiling at — 33.50. To effect the 
condensation of chlorine take a bent glass tube (Fig. 37), 
introduce into the leg closed at 
one end some crystals of chlo- 
rine hydrate (Cl2 + ioH20, see 
below), then seal the open end. 
The limb containing the com- 
pound is placed in a water-bath ; 
the other is cooled in snow. 
Upon heating the water a 
little above 30° the chlorine 
hydrate is decomposed 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 chlorine distils back and is reabsorbed by the 
water. Charcoal saturated with chlorine may be substituted 
for the chlorine hydrate. This substance takes up 200 vol- 
umes of chlorine, which are disengaged again 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 prop- 
erties of the free gas ; it is therefore frequently employed for 
laboratory uses as a substitute for chlorine. The yellow, scale- 
Fig. 37. HALOGENS. 
49 
like crystals of chlorine hydrate (Cl2 + ioH20) separate 
when water saturated with the gas is cooled below o°. This 
compound is regarded as one of chlorine with water. At 
ordinary temperatures it decomposes into its constituent mole- 
cules. 
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 
(false 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 just as energetically with hydrogen. 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 hydro- 
gen compounds; most of these are so decomposed by the 
chlorine that it removes the hydrogen from them, and forms 
hydrochloric acid. Thus water is decomposed by chlorine 
into hydrochloric acid and oxygen : 
H20 + Cl2=2HCl + 0. 
If a glass cylinder be filled with, and inverted over chlorine 
water and exposed to the sunlight, a gas will be evolved, and 
will collect in the upper portion of the vessel; 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 
separated in a free condition. A piece of tissue paper satu- 
rated with newlv distilled turpentine oil, and 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 decolorized 
bv moist chlorine gas. The same occurs with the dark blue 
solutions of indigo and litmus ; colored flowers are rapidly 
bleached by it. On this principle depends the application of 
chlorine in bleaching goods, and in destroying decaying 
matter and miasmata in chlorine disinfection. (See Bleaching 
Lime). 
The bleaching action of chlorine is mostly influenced by the presence 
of water. It probably depends on the oxidizing action of the oxygen lib- 50 
INORGANIC CHEMISTRY. 
erated 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 ele- 
ments, at the moment of their birth (in statu nascendi), act more ener- 
getically than when free; the cause for this will be explained hereafter. 
Br = 79.76. Br2= 159.5. 
2. BROMINE. 
Bromine, the perfect analogue of chlorine, was discovered 
by Balard, in 1826. It occurs in sea water as sodium bromide, 
accompanied by sodium chloride, but in much smaller quan- 
tity than the latter (especially in the water of the Dead Sea), 
and in many salt springs, as at Kreutznach and in Hall. When 
sea water or other salt water is evaporated, sodium chloride 
first separates; in the mother-liquor, among other soluble 
salts, are found sodium and magnesium bromides. Bromine 
is found in greatest abundance in the upper layers of the rock- 
salt deposits of Stassfurth, near Magdeburg, where it exists in 
the form of bromides together with other salts. At present, 
large quantities of bromine are obtained in America. The 
method of its preparation is similar to that employed under 
chlorine. A mixture of manganese dioxide and sodium 
bromide is warmed with sulphuric acid : 
Mn02 -f- 2NaBr -)- 2H2S04 = MnS04 -f- Na2S04 -)- 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, containingthe mixture, isheated in awater-bath ; 
the tube B serves to cool the vapors which are condensed by 
cold water or ice in the receiver C. When free chlorine is 
conducted into an aqueous solution of sodium bromide, bro- 
mine separates out. 
Bromine is a heavy, reddish brown liquid, of an exceed- 
ingly penetrating, chlorine-like odor (hence the name Bro- 
mine, from ftpw/joq, stench). At —7.30 it crystallizes to a yellow- 
green, scaly mass, having a metallic lustre, and resembling 
iodine. Liquid bromine at o° has the specific gravity 3.18 
(water == 1) ; it is very volatile, yielding dark brown vapors 
at 63° C., changing, at the same time, into a yellowish-brown 
vapor. Its density equals 79.7 (hydrogen = 1), or 5.53 
(air = 1). 
Bromine is more soluble in water than chlorine. Cooled 
below 40 C., the hydrate (Br2 -f- ioH20) crystallizes out: 
his is analogous to the chlorine hydrate. It is decomposed at HALOGENS. 
51 
moderate temperatures. Bromine dissolves with ease in alco- 
hol, and especially in ether, chloroform and carbon disulphide. 
In a chemical point of view, bromine is extremely like 
chlorine, combining directly with most metals to form bro- 
Fig. 38. 
mides ; but it possesses a weaker affinity than chlorine, and is 
liberated by the latter from its compounds: 
With hydrogen it only combines on warming, not in sun- 
light. Upon hydrocarbons it acts like chlorine, withdrawing 
hydrogen from them. Bromine water gives starch an orange 
color. 
KBr + Cl = KC1 + Br. 
1 = 126.54. I2= 253.08. 
3. IODINE. 
Iodine, as Well as bromine, occurs in combination with 
sodium, in sea water and some mineral springs, especially 
at Hall, in Austria, and the Adelheit Spring in Bavaria. In 
these springs the iodine can easily be detected directly ; in 
sea water it is, however, only present in such minute quan- 
tity-that its separation, practically, is disadvantageous. Sea 
algae absorb it from the water, and these are then thrown by 
the tide on various coasts, where they are burned, yielding an 
ash (known as kelp in Scotland, as varec in Normandy) which 
is the 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 Davy and Gay-Lussac, 52 
INORGANIC CHEMISTRY. 
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 saltpetre. It is set free from its cofnpounds in 
the same manner as chlorine and bromine—by distillation 
with manganese dioxide and sulphuric acid. It is more con- 
venient, however, to pass chlorine (or, better, nitrous acid) 
through a solution of the salt, when all the iodine will sepa- 
rate : 
KI -f Cl = KC1 + I. 
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 strongly so 
as bromine. Its specific gravity is 4.95. It fuses at 1130 to 
a dark brown liquid, and boils near 200°, passing at the same 
time into a dark violet vapor (hence the name Iodine, from 
\ atdrj 7, violet- blue'). 
The vapor density of iodine equals 8.7 up to 6oo° C. ( air = 1) or 126.5 
(H = 1), corresponding to the molecular weight I2 = 253. Above 700° 
the vapor density gradually diminishes and about 1500° it is only "half the 
original. This is explained by the gradual decomposition (see Dissoci- 
ation of Water) of the normal diatomic molecule I2 into the free atoms 
I —j— I. In like manner the bromine molecules Br2, at high temperatures, 
suffer a separation into the free atoms, while the gas density of chlorine 
manifests only a slight change ; its molecules are, therefore, dissociated 
with far greater difficulty. 
Iodine is very slightly soluble in water, more readily in 
alcohol (Tinctura Iodi), very easily in ether, chloroform and 
carbon disulphide, the last two assuming a deep red-violet 
color in consequence. It affords a particularly beautiful crys- 
tallization, consisting of forms of the rhombic system, when 
it separates from a solution of glacial acetic acid. 
In chemical deportment iodine closely resembles bromine 
and chlorine ; it possesses, however, weaker affinities, and for 
this reason is liberated from its compounds by those elements. 
With the metals it usually combines only when warmed ; with 
hydrogen it does not combine directly, and it does not re- 
move it from its carbon compounds. 
The deep blue color it imparts to starch is characteristic of 
iodine. On adding starch-paste to the solution of an iodide, HALOGENS. 
53 
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. 
FI = 19. (Fla = 38.) 
4. FLUORINE. 
Fluorine possesses such a 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), from which its other deriva- 
tives are prepared ; the latter resemble those of chlorine. 
When silver fluoride is heated in a stream of chlorine, or cal- 
cium fluoride in a current of oxygen, there escapes a colorless, 
very strongly smelling gas, which attacks glass and all the 
metals. Upon conducting the same into water, oxygen is 
disengaged from the latter: 
H20 -j- Fl2 = 2HF1 -f o. 
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). 
These four similar elements, fluorine, chlorine, bromine, 
and iodine, exhibit gradual differences in their properties ; 
and, what is remarkable, this gradation stands in direct rela- 
tion to the specific gravity of the elements in the state of gas 
or vapor. 
FI Cl Br I 
Specific gravity 19 35*37 79-76 126.5. 
With the increase of specific gravity occurs a simultaneous 
condensation of matter, which expresses itself in the dimin- 
ished volatility. Fluorine is a gas; chlorine can readily be 
condensed to a liquid ; bromine is a liquid at ordinary tem- 
peratures, and iodine is a solid. Other physical properties, 
as seen in the following table, are also in accord with the 
preceding. 54 
INORGANIC CHEMISTRY. 
Fluorine. 
Chlorine. 
Bromine. 
Iodine. 
Fusing point 
Boiling point 
Specific gravity in liquid or 
solid condition 
Color 
Colorless 
—33° 
*33 
Yellow 
-7° 
+63 
3*8 
Brown 
+ "3° 
—(-200° 
4-97 
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 
metallic and hydrogen compounds by the lower. We shall 
discover, later, that in the affinity-energy of the halogens for 
oxygen and some other metalloids, the reverse is true. 
COMPOUNDS OF THE HALOGENS WITH 
HYDROGEN. 
With hydrogen the halogens form compounds of an acid 
nature, readily soluble in water. 
1. HYDROGEN CHLORIDE. 
The direct union of chlorine with hydrogen takes place 
through the agency of heat, and by the action of direct sun- 
light or other chemically active rays; in diffused light the 
action is only gradual, and does not occur at all in the dark. 
On introducing 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 atmosphere of hydrogen, may be shown easily by the fol- 
lowing 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 con- 
duct dry chlorine into the cylinder. The burning hydrogen 
will inflame the chlorine, which will continue to burn in the 
former. From these experiments, we perceive that combusti- 
bility and combustion are only relative phenomena; if hy- 
drogen is combustible in chlorine (or air), so, inversely, is 
chlorine (or air) combustible in hydrogen. By the term 
HC1 == 36.37. Density = 18.18. HYDROGEN CHLORIDE. 
55 
combustion, in chemistry, is understood every chemical union 
of a body with a gas, which is accompanied by the phenom- 
enon of light. 
A mixture of equal volumes of chlorine and hydrogen ex- 
plodes under the conditions given above for the union of the 
Fig. 40. 
Fig. 39. 
gases, with very great violence. The product is gaseous hy- 
drogen chloride. 
The formation of the latter compound succeeds best by 
allowing sulphuric acid to act upon sodium chloride, when 
solid sodium sulphate and hydrogen chloride gas will result. 
Pour over 5 parts sodium chloride, 9 parts sulphuric acid, somewhat di- 
luted with water (2 parts), and warm the mixture gently in a flask, A 
(Fig. 41). The escaping hydrogen chloride is conducted 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 aTl moisture, and 
afterwards collected over mercury. 
2NaCl -f H2SOt = Na2S04 -f 2HCI. 
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 io° 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 — no° and boils 
at — 80.3°. 
The specific gravity (density) of the gas is 18.18 (H = 1), 
or 1.26 (air =1). 
Hydrogen chloride possesses an acid taste, and colors blue 56 
INORGANIC CHEMISTRY. 
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 o° C. 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 cyl- 
inder with the gas and then just dipping its open end beneath 
water; the latter rushes up into the vessel rapidly (as into a 
Fig. 41. 
vacuum), as it quickly absorbs the gas. The aqueous solution 
of hydrogen chloride, in ordinary language is known as muri- 
atic or hydrochloric acid (acidum hydrochloratum). For its 
preparation t-he gas is passed through a series of Woulff bot- 
tles (Fig. 42) containing water. The small bottle B'm which 
there is but little water, serves to wash the gas—free it of any 
mechanically admixed sulphuric acid. The same apparatus 
may be employed in the manufacture of chlorine water, and is 
generally used in the saturation of liquids with gases. 
A solution saturated at 150 C. contains about 40 per cent, 
hydrogen chloride, has a specific gravity of 1.2, and fumes in 
the air. On the application of heat, the gas again escapes, 
and the temperature of the liquid rises to iio° C., when a 
liquid distils over, containing about 20 per cent, of hydrogen 
chloride, having a specific gravity of 1.104 and almost cor- HYDROGEN CHLORIDE. 
5 7 
responds to the formula HC1 -f- 8H20. The composition of 
th*e distillate varies somewhat with the pressure. On con- 
ducting hydrogen chloride into hydrochloric acid cooled 
Fig. 42. 
to — 220, crystals of the formula HC1 -f- 2H20 separate ; these 
fuse at — 180 and then decompose. 
Hydrochloric acid finds an extensive industrial application, 
and is obtained in large quantities, as a by-product, in the 
soda manufacture. 
Chemical Properties. —Acids— Bases — Salts. — Hydrogen 
chloride, as well as its solution, possesses all the properties 
of acids, apd can well figure as a prototype of these ; it 
tastes intensely acid, reddens blue litmus paper, and saturates 
the bases (oxides and hydroxides) ; i.e., such bodies as impart 
a blue color to red litmus paper. If we add hydrochloric acid 
to a solution of a base, e.g., sodium hydroxide, until the reac- 
tion is neutral, we will obtain (besides .water), a neutral, solid 
compound—sodium chloride. 
NaOH- + HC1 == NaCl + HaO. 
Sodium 
hydroxide. 
Sodium 
chloride. 
HBr, HI, and HF1 deport themselves similarly to HC1. 
These halogen compounds of hydrogen are termed haloid 58 
INORGANIC CHEMISTRY. 
acids, to distinguish them from those which, in addition to 
hydrogen, contain oxygen, hence called oxygen acids. The 
latter conduct themselves like the former, and saturate bases, 
forming salts and water : 
KOH + UNO, = KNOj + H20. 
Potassium 
hydroxide. 
Nitric 
acid. 
Potassium 
nitrate. 
Water, 
In the same manner the acids act upon the basic oxides, to 
form salts and water: 
ZnO + 2IIC1 = ZnCl2 + H20. 
Zinc 
oxide. 
Zinc 
chloride. 
Zinc 
oxide. 
ZnO + 2HNO* = Zn(N03)2 + H20. 
Zinc' 
nitrate. 
Usually when acids act upon metals, 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 hydrogen result: 
From the examples cited it is manifest that acids are hydro- 
gen compounds which yield salts, by the replacement of their 
hydrogen by metals, by the action of metallic oxides, hydrox- 
ides, or free metals. The metallic oxides and hydroxides capa- 
ble of forming salts by the saturation of acids are called l?ases. 
Finally, by the term salts, we understand such compounds as 
are analogous'to sodium chloride, and are formed by the mutual 
action of bases and acids. Salts are distinguished as haloid 
salts dxi&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 established analytically 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 peroxide, on the other hand, be heated in it, 
chlorine will be disengaged : 
Na + HC1 = NaCl + H. 
If the electric current be permitted to act upon an aqueous 
solution of hydrochloric acid, the latter will be so decom- 
MnOj + 4HCI =' MnCl2 + 2H20 -f Cl2. HYDROGEN BROMIDE. 
59 
posed that chlorine separates at the electro-positive and hy- 
drogen at the electro-negative pole. (See p. 72.) 
HBr = 80.76. Density = 40.38, 
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. 45.) Like hy- 
drogen chloride, hydrogen bromide can be obtained by the 
action of some acids, e.g., phosphoric acid, upon bromides; 
sulphuric acid would not answer, as the resulting HBr is again 
.partly decomposed by it. Ordinarily it is prepared by the 
action of phosphorus tri-bromide (see Phosphorus) upon water: 
PBr3 + 3H20 = H3PO3 + 3HBr. 
Phosphorus 
tri-bromide. 
Phosphorous 
acid. 
Place some water (1 part) in a flask (Fig. 43), gradually ad- 
mit through the funnel, supplied with a cock, the liquid, PBr3 
(3 parts), and warm gently. The escaping HBr gas is collected 
Fig. 43. 
over mercury or conducted into water. To free it perfectly 
from accompanying PBr3 vapors it is passed through water 
(the U-shaped tube in Fig. 43 contains pieces of pumice stone, 
which are moistened with water). 
Instead of employing prepared brom-phosphorus, we may let bromine 
vapors act upon (red) phosphorus. This may be done by pouring water 60 
INORGANIC CHEMISTRY. 
or dilute hydrobromic acid (2 parts) over the phosphorus placed in a flask. 
Bromine (10parts) is added gradually while cooling and heat then applied. 
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. 
Hydrogen bromide is a colorless gas, fuming strongly in the 
air. Under great pressure it is condensed to a liquid, boiling 
at — 73-3° and solidifying at — 120°. Its density is 40.38 (H 
= 1) or 2.79 (air = x). 
In water the gas is very readily soluble, its saturated solution 
having a specific gravity of 1.78, and containing 82 per cent. 
HBr; at 150 it contains 49.8 per cent., and has the specific 
gravity of 1.515. At 1250 C. a solution distils over contain- 
ing 46.8 per cent. HBr, and closely approximates the formula 
HBr + 5H20 ; its specific gravity is 1.47 at 140 C. 
On conducting HBr into a solution of the same cooled to 
— 200, crystals of the formula HBr + 2H2O separate and melt 
at — ii°. Chemically, HBr is the perfect analogue of 
HC1; it is, however, less stable, and suffers a partial decom- 
position at 8oo° C. 
HI = 127.54. Density = 63.7. 
3. HYDROGEN IODIDE. 
The attraction of iodine for hydrogen is very slight. Their 
partial union 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 : 
PI3 + 3H20 = P03H3 + 3HI; 
A more convenient procedure is to warm a mixture of 
amorphous phosphorus (i part), iodine (15 parts), and water 
(14 parts) ; when an analogous reaction will ensue. Or add a 
solution of 2 parts iodine in r part hydriodic acid, of specific 
gravity 1.67 (obtained by distillation, see below), drop by 
drop, to red phosphorus, and aid the reaction 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. HYDROGEN IODIDE. 
61 
44), where it will displace the air in consequence of its five- 
fold greater density. 
Fig. 44. 
' To get an aqueous solution of HI, take more water, warm 
the solution, filter, and then distil. 
Another method for obtaining HI consists in passing hy- 
drogen sulphide into water having iodine dissolved in it until 
there is no further decomposition : 
H2s + h = 2HI + s. 
Filter off the separated sulphur and distil the liquid. 
Hydrogen iodide is a colorless gas; it fumes strongly in the 
air ; its density is 63.7 (H = 1) or 4.41 (air =1). Under a 
pressure of 4 atmospheres (at o°) it is condensed to a liquid 
which solidifies at — 550. It is easily soluble in water; the 
solution saturated at o° C., has a specific gravity 1.99, and 
fumes strongly in the air. At 1270 a solution of 1.6.7 specific 
gravity, and containing 57.7 per cent. HI, distils over, corre- 
sponding closely to the formula HI + 5H20. 
Hydrogen iodide is a rather unstable compound, decom- 
posing at 1800 into hydrogen and iodine. At high tempera- 
tures oxygen decomposes ifinto water and iodine : 
On bringing a flame near a vessel containing a mixture of 
HI and oxygen, violet iodine vapors will at once fill it. The 
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 
2HI + o = h2o + It. 62 
INORGANIC CHEMISTRY. ' 
their oxygen and reduces them. The oxygen of the air, even 
at the ordinary temperature, and especially in sunlight, gradu- 
ally 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 : 
Ill + Ag = Agl + H. 
Chlprine and bromine liberate iodine from HI. 
This compound is employed as a powerful reducing agent 
in laboratory work. 
4. HYDROGEN FLUORIDE. 
It is obtained, like hydrogen chloride, by decomposing 
fluorides with sulphuric acid. Finely pulverized fluoride is 
mixed with H2S04 and heat applied gently: 
HF1 = 20. Density = io (at ioo°) 
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. 
receiver containing some water. To get perfectly anhydrous 
hydrogen fluoride, heat hydrogen potassium fluoride, which 
then decomposes according to the following equation : 
KFljH = KFl-f HFl.^ HYDROGEN FLUORIDE. 
63 
Anhydrous hydrogen fluoride is-a colorless, very mobile 
liquid, fuming strongly in the'air, and attracting moisture 
with avidity; it boils at rf- 190 C., and has a specific gravity of 
0.98 at 120. To recondense the gas it must be cooled to 
-— 20°. 
The gas density of hydrogen fluoride equals 10 (hydrogen = 1) at ioo°, 
corresponding to the molecular formula HF1 = 20. At 30°, however, it 
is twice as large, equalling 20. It follows, therefore, that the molecules 
of the gas at the latter temperature correspond to the formula H2FI2, and 
consist of two chemical molecules of HF1. (Compare arsenic trioxide.) 
The concentrated aqueous solution fumes in the air; when 
heated HF1 escapes; the boiling temperature increases regu- 
larly until- constant at 120° C., a solution distils over, the 
specific gravity of which is 1.15, and its percentage of HF1 
35.3. The vapors as well as the solution are poisonous, ex- 
tremely 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 sil- 
icate, is also acted upon ; hence the use of the acid for etching 
this substance. (Compare silicon fluoride.) To do this, coat 
the glass with a 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 HF1. The exposed portions appear etched ; 
gaseous HF1 furnishes a dim, and liquid HF1 a smooth, trans- 
parent etching. 
Vessels of lead, platinum, or caoutchouc are employed for 
the preservation of hydrofluoric acid, as they are not affected 
by it. 
These 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 they are condensed by the aqueous vapor. Readily 
soluble in water, they are only partially expelled from their 
solution by boiling ; solutions of constant quantity distil over ; 
these may be regarded as chemical combinations of the hal- 
ogen hydrides with water. As acids they neutralize the bases 
and form haloid salts, which also can result by the direct union 
of the halogens with metals. 64 
INORGANIC CHEMISTRY. 
The densities of the halogen hydrides exhibit a gradation 
similar to that of the densities of the halogens (page 53) : 
HF1 HC1 -HBr HI 
Density, 10. 18.18 40.38 63.7. 
The difference in chemical deportment corresponds to this 
gradation. Hydrogen fluoride is the most' stable, and acts 
most energetically; chlorine combines with hydrogen in sun- 
light, 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 (r8o°), into its constituents; 
the more stable hydrogen bromide at 8oo°, while hydrogen 
chloride remains unaltered up to 1500° C. Corresponding to 
this we have the very energetic action of fluorine, and the tol- 
erably ready action of chlorine upon water, oxygen separating 
at the same time : 
H20 -f- Clj = 2HCI + O. 
Iodine does not act upon water. The opposite reaction 
occurs: oxygen decomposes hydrogen iodide into water and 
iodine : 
2HI + O = H2Q -(- I2. 
Bromine occupies an intermediate position between chlorine 
and iodine ; in aqueous solution it decomposes water into 
HBr -j- O, while a concentrated solution of hydrogen bro- 
mide, 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 follow chlorine and bro- 
mine, and finally, as the least energetic element, we have 
iodine. (See p. 54.) 
Thermo-chemistry : A measure of the chemical affinity of the halo- 
gens for hydrogen, also of the chemical elements and bodies for each other, 
is afforded by the quantity of heat (heat modulus) disengaged or absorbed 
in chemical union (compare pp. 23 and 25). This is determined in heat 
units or calories. That amount of heat is selected as unit which will raise 
1 gram of water from o° to i° C. (small calorie). 
The quantities of the combining elements are expressed in grams, in num- 
bers corresponding to their atomic weights. Hence in the union of 35.37 
grams chlorine (Cl) with 1 gram hydrogen (H), 22,000 calories, and in the 
union of 79.7 grams bromine with 1 gram hydrogen, 8400 calories are dis- 
engaged while in the union of 126.5 grams iodine with 1 gram hydrogen, 
6040 calories are absorbed. This may be more simply expressed accord- 
ing to the method of J. Thomsen, as follows: 
(H,C1) — -f- 22,000. (H,Br) — + 8400. (H,I) = — 6040. HYDROGEN FLUORIDE. 
65 
The first two reactions, in which heat is liberated, are exothermic, while 
the heat absorbing combination of iodine with hydrogen represents an 
endothermic reaction. 
We. must remind the reader that the observed heo.t modulus is not a 
direct measure of chemical affinity. The elements do not exist as free 
atoms, but as molecules that require a definite quantity of heat to decom- 
pose them into atoms. The union of chlorine with hydrogen proceeds 
according to the molecular equation : 
The heat here disengaged indicates that the affinity of 2H for 2CI is 
2 X 22000 calories greater than the affinity of H for H -(- Cl for Cl. Simi- 
larly, the heat absorbed in the formation of hydrogen iodide, shows that 
the affinity of 2I for 2H is 2 X 6040 calories less than that of the atoms 
H and I in their molecules. The heat modulus, therefore, is only a rela- 
tive measure of chemical affinity. 
The greater the heat developed in a reaction, the more energetically 
and the more readily will it occur, and in general, the resulting compounds 
will be the more stable. In accordance with this, as we have seen, chlo- 
rine and hydrogen unite readily with each other, while in the union of 
hydrogen and iodine, where heat is absorbed, the combination occurs with 
difficulty, and can only be effected by the addition of energy (heat). Con- 
versely, hydrogen iodide is easily decomposed into its elements, while 
hydrogen chloride is very stable. In its entire chemical deportment, hy- 
drogen bromide corresponds to its heat of formation, and occupies a posi- 
tion intermediate between HQ and HI. In similar manner is explained 
how, according to the law of greatest heat development, iodine is eliminated 
from its compounds with hydrogen and the metals, by bromine and chlo- 
rine, and also bromine by chlorine. Remembering the thermal relations 
in the formation of water, we can explain in the same manner, the varying 
decomposition of the halogen hydrides by oxygen, and the reverse—that 
of water by the halogens. To illustrate these interesting relations, let us 
consider the formation of hydrogen iodide, by action of hydrogen sulphide 
on iodine (p. 61), corresponding to the reaction: 
HH + C1C1 = 2HC1. 
HaS -f-1* = 2III + S. 
Since in the production of hydrogen iodide, in the above equation, 12080 
(= 2 X 6040) calories are absorbed—this reaction cannot occur of itself 
without added energy. It does take place in the presence of water as the heat 
afforded by the solution of HI in water is sufficient. In the solution of 
halogen hydrides, as.well as 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 2HI in water, exceeds 
the heat of formation of 2-HI (— 12,080 C.) and the decomposition heat of 
H2S (—4500 C.), so that the reaction can occur even with disengagement 
of heat. The transposition is incomplete if the quantity of water employed 
be small. * * 66 
INORGANIC CHEMISTRY. 
COMPOUNDS OF THE HALOGENS 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, mono-chlor-iodine results, and when the chlo- 
rine is in excess, there is formed trichlor-iodine. 
Iodine Chloride—IC1—is a red crystalline mass, fusing 
at 24.70 C., and distilling a little above ioo° C. Water de- 
composes 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 I205. It crystallizes in long, yellow needles, and, 
when heated, suffers decomposition into IC1 and chlorine 
(at ordinary pressure, the dissociation commences at 250 C.); 
dissolves in a little water without alteration ; large quantities 
cause partial decomposition, with formation of iodic acid. 
Iodine Bromide—IBr—obtained by thedirect union of the 
elements, consists of iodine-like crystals, fusing at about 30°. 
Iodine Pentafluoride—IF1.—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 halogen derivatives considered, as well as in all 
other chemical compounds, we determine the quantity of the 
elements (according to methods described in analytical chem- 
istry) we will discover that they are always combined with 
each other in the same proportions by weight. In every chem- 
ical compound the proportions by weight of the constituents con- 
tained in it are invariably the same. Thus chemical analysis 
shows the following percentage composition for the halogen 
derivatives of hydrogen : 
H S.o 
H 
= 2.7 
H = 1.2 
K 
p 
co 
FI = 95.0 
Cl 
= 97-3 
Br = 98.8 
I =99.2 
HF1 =100.0 
HC1 
=100.0 
HBr =100.0 
HI = 100.0 
Experience has shown that hydrogen, of all the elements, 
enters compounds in the least quantity, therefore its quantity is WEIGHT PROPORTIONS. 
67 
chosen as unity, and we calculate those weights of the elements 
which combine with one part by weight of H. In this man- 
ner we find the following proportions for the halogens: 
H — i H = i H == i H = i 
FI = 19 Cl = 35.37 Br — 79.7 I =126.5 
HF1 = 20 HC1 = 36.37 HBr = 80.7 HI = 127.5 
Experiments have also established the remarkable fact that 
the same proportions of the halogens by weight are also ob- 
tained by the union of the same with other elements. Thus 
19 parts of FI by weight combine with the following weights 
of the metals: 23 parts Na, 39 parts K, 32.4 parts Zn, 31.6 
parts Cu, 99.9 parts Hg, and 35.37 parts Cl. 79.7 parts Br, 
and 126.5 Parts I combine with exactly the same quantities of 
these metals by weight. Let us take another example. On 
bringing copper into the solution of a mercuric salt the former 
dissolves, while Hg separates out; indeed, 31.6 parts Cu dis- 
place 99.9 parts Hg. If zinc be brought into the copper so- 
lution thus obtained, it will dissolve, while copper separates— 
and 32.4 parts of Zn separate 31.6 parts Cu. Furthermore, 
zinc displaces the hydrogen in acids ; from all of them 32.4 
parts Zn separate 1 part H. In all these reactions we observe 
the elements appearing in the same quantities by weight. 
These remarkable facts are fully verified by experiments. 
Such facts may be formulated in a rule, and when a rule com- 
prises 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 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 in- 
cludes a long series of facts, if it is repeatedly substantiated by 
other phenomena and has acquired a high degree of proba- 
bility, 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 hy- 
pothesis may be necessary. Conversely, something which long passed as 
a fact or a theory may be shown to be erroneous, if not any longer con- 
sistent with new observations. Hypothesis and that which we designate 
a fact, are distinguished really by the different degree of probability only. 
If for example, we make a sight observation we assume the hypothesis 
that the same has been caused by an external process, of the reality ot 68 
INORGANIC CHEMISTRY. 
which (in distinction from subjective perceptions) we can only assure our- 
selves 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. The same may be true with regard to the 
supposition of atoms—whether we comprehend them as 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 phi- 
losophers even conjectured that matter consisted of indivisible 
and very small particles—atoms (from a privative, and rotioq, 
division). This a priori supposition was subsequently re- 
peatedly 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 these representations, the law of constant proportions 
becomes very simple ; we can comprehend that the quantities 
of the constituents of a compound should be constant, and 
that the relative quantities of the elements by weight, must be 
the same in all their compounds, as they express the relative 
weights of the atoms. 
As yet only the relative atomic weights of the elements have been de- 
termined by chemical researches; in these the hydrogen atoms, as they 
possess the least weight, have been taken as unity. Until now the knowl- 
edge of the absolute atomic weight, for chemical considerations, has been 
unessential. At the present time different physical phenomena permit fixing 
the absolute size of the atom with considerable approximate accuracy. 
Very different considerations lead to the same conclusion, that the atoms 
cannot be smaller than the fifty millionth part of a millimeter (Thomson). 
We can determine the diameter of the molecules more accurately with the 
gases. With hydrogen H2 = 4, with N2 = 3, with O2 = 7, it has been 
found to be the xo millionth part of a millimetre. 
If we grant that in the preceding halogen-hydrogen com- 
pounds one atom of hydrogen is .combined with every halogen 
atom, the conclusion follows, that the ratio found expresses the 
relative atomic weights of the halogens. This supposition, 
however, appears questionable, in view of the mor'e compli- 
cated proportions which occur in the union of some elements. 
Observation shows, to wit, that very frequently two elements ATOMS, 
69 
unite with each other in- not only one, but, indeed, several 
proportions. For example, 35.37 parts of chlorine combine 
not only with 31.6 parts copper and 99.9 parts mercury, but 
also with 63.2 parts copper and 199.8 parts mercury. One 
part, by weight, of hydrogen, combines with 8 parts of oxygen 
(more accurately 7.98) to form water, and with 16 parts oxy- 
gen (to form the so-called hydrogen peroxide ) ; further, with 
16 and 32 parts sulphur. Oxygen forms five different com- 
pounds with nitrogen according to the following proportions, 
by weight: 
Nitrogen. 
Oxygen. 
Nitrous oxide, . . 
. 14 
parts. 
8 
parts = 
= I 
X 
8. 
Nitric oxide, . . . 
• 14 
parts. 
16 
parts = 
= 2 
X 
8. 
Nitrous anhydride, . 
. 14 
parts. 
24 
parts = 
= 3 
X 
8. 
Nitrogen dioxide, . 
. 14 
parts. 
32 
parts = 
= 4 
X 
8. 
Nitric anhydride, 
. 14 
parts. 
40 
parts .= 
= 5 
X 
8. 
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 ele nts 
(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 P, .portions, 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 that 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 likely 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 = 16, in the latter = 4. • 
Analytical results afford nothing positive for the solution of 
this difficulty. This was the condition in which the.question 
relating to the magnitude of the atomic weights existed thirty 
years ago. To establish these correctly, various views were 
allowed to prevail, none, however, with positive foundation. 
The question can only be solved upon a new and accurate 
basis: the specific gravities of the chemical compounds in a 
gaseous or vapor form answer well for this purpose. 70 
INORGANIC CHEMISTRY. 
DENSITY OF BODIES IN STATE OF GAS. VOLUME RATIO IN 
THE UNION OF GASES. ATOMIC MOLECULAR THEORY. 
The halogens, fluorine, chlorine, bromine, and iodine, unite 
with hydrogen in only one proportion. The supposition, 
therefore, that in the halogen-hydrogen compounds, 1 atom 
of H is combined with 1 atom of the halogen, is the simplest 
and most probable. Then their weight proportions, derived 
from analysis directly express their relative atomic weights. By 
comparing these atomic numbers (referring to H = 1) with 
those expressing the density in a state of gas (also referred to 
H = 1) the astonishing result is seen that the two series are 
identical. 
Elements. 
Density 
Air = i 
Density. 
Hydrogen = 1 
Atomic Weights. 
Hydrogen, 
0.0692 
I 
I 
Fluorine, 
• (1-31) 
(19) 
19 
Chlorine, • 
2-45 
35-37 
35-37 
Bromine, 
5-52 
79-7 
79-7 
Iodine, 
8-75 
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 an equal 
number of atoms. Indeed, if in one volume of hydrogen, for 
example, there are contained iooo atoms of hydrogen, which 
equal iooo weight units, and in a like volume of chlorine 
there are also present iooo atoms of chlorine, which equal 
iooo X 35-37 weight units, then it is evident, that the rela- 
tion between the atomic weights, and that between the densi- 
ties (the weights of like gas volumes) must be the same. 
1000 X 1 
iooo X 35.4 
1 vol. Hydrogen. 
1 vol. Chlorine. 
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. Yielding 
to a too hasty generalization, this was incorrectly followed for 
all elements. 
We arrive at a perfectly similar, but much more general 
conclusion, by the consideration of the physical properties of GAS VOLUME. 
71 
gases or vapors. The similar deportment of the same under 
pressure (law of Mariotte and Boyle), their similar expansi- 
bility 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, 
which are separated by equal distances, very great in pro- 
portion to the particles (the distances of the centres are equal 
and suffer equal alterations.) It immediately follows from this, 
that equal numbers of particles are contained in equal volumes 
of all gases (under like temperature and pressure). The kinetic 
gas theory, based on the same supposition, explains the similar 
deportment of gases by the equal kinetic energy of the smallest 
gaseous particles. 
From the proposition, that in equal volumes an equal num- 
ber of particles is present, it follows directly that their relative 
weights are proportional to the volume weights or gas densities, 
and that by the determination of the latter, the first are also 
given. In what ratio these smallest particles (called mole- 
cules) stand to the chemically smallest particles (atoms), re- 
mains, first of all, undetermined, and can only be obtained by 
a comparison of the volume ratios according to which the 
bodies combine (p. 70). It is, however, even now seen that, 
at least in the case of compound bodies, the smallest gas par- 
ticles must be sums of atoms, as the same consist of combina- 
tions of atoms. 
It follows, from the equality of the atomic weights and the 
densities, that the halogens must combine with hydrogen in 
equal volumes, since 1 part of H by weight combines with 
35.37 parts of chlorine by weight, etc., and the weights of 
equal gas volumes stand in the same ratio. Further: 1 part 
H and 35.37 parts chlorine yield 36.37 parts HC1; one volume 
of the latter weighs, however, 18.1 (H = 1, p. 77); conse- 
quently, 36.3 parts HC1 occupy 2 volumes. Therefore, equal 
volumes of H and Cl yield a double volume of HC1, or, as 
ordinarily expressed, 1 volume H and 1 volume Cl yield 2 
volumes HC1. In a similar manner it may be deduced that 1 
volume H and 1 volume Br vapor yield 2 volumes HBr ; that 
1 volume H and 1 volume I vapor yield 2 volumes HI. 
These conclusions are confirmed by the following experiments: 
i. The concentrated aqueous solution of hydrochloric acid is decom- 
posed by the action of the galvanic current, and the chlorine and hydrogen 
collected; these gases separate at opposite poles. The electrolysis may 
be made in an ordinary voltameter (Fig. 46). Hofmann’s apparatus is better 72 
INORGANIC CHEMISTRY. 
adapted to this purpose (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 till 
the apparatus with liquid; and also by further 
additions, to press out the 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 introduced by means of caoutchouc corks. 
When the separating gases (in this case the chlo- 
rine) attack the platinum, carbon electrodes are 
substituted for the latter. 
To electrolyze hydrogen chloride, fill the ap- 
paratus 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 proved 
to be chlorine; the other combustible gas is hy- 
drogeti. 
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—is shown in the next ex- 
periment : 
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 GAS VOLUME, 
73 
conveniently done by conducting the gaseous mixture obtained by the elec- 
trolysis 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 mix- 
ture, 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 the hydrogen 
chloride that was produced dissolves; all hydrogen and all chlorine have 
disappeared. 
3. A modification of this experiment teaches us another important fact 
whjch has reference to the ratio of the volume of the hydrogen chloride to 
the volumes of its constituents. If the tube filled with equal volumes of 
Cl and H be opened, under Hg, 
after the explosion, no diminu- 
tion in volume will be detected, 
although the mixture of Cl and 
H has been changed to hydro- 
gen 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 vol- 
FiG.48. 
Fig. 49. 
time of Cl and one volume of H yield two volumes of hydrogen chlo- 
ride. 
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 FIC1 is decompos-ed, the Cl combines with the Na to form 
sodium chloride, while hydrogen is set free. Upon measuring the residual 
hydrogen it will be found that its volume is exactly the half of the 74 
INORGANIC CHEMISTRY. 
volume of HC1 originally introduced. In the same manner may be 
shown the fact that in two volumes of HBr and HI there is con- 
tained in each one volume 
of H. It follows further 
from the densities of bro- 
mine and iodine vapors, 
that the quantities of these 
elements in gas form com- 
bining with one volume of 
hydrogen also occupy one 
volume. Hence, one vol- 
ume 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. 
111G~ 5°~ 
The volume ratios in the chemical union of gases were first investigated 
by Humboldt and Gay-Lussac (1805-1808). The latter derived the two 
following empirical laws by experiment: (1) Gases unite according to 
simple volume ratios; (2) The volume of the resulting body bears a simple 
ratio to the volumes of the constituents. 
Comparing this fact announced by Gay-Lussac, that in the chemical 
union of gases simple volume ratios do occur, with that discovered by Dal- 
ton (p. 67), that the quantities by weight of the combining elements also 
bear a simple ratio, and granting the atomic constitution of matter, it fol- 
lows that the number of smallest gas particles (molecules) contained in 
equal 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 Ampere in 1814. 
As deduced on page 71, and confirmed by the described 
experiments, the quantities of the halogen-hydrogen com- 
pounds by weight, expressed by the chemical formulas, HC1, 
HBr, HI, occupy a volume twice as large as one part by 
weight of H, or 35.37 parts Cl, 79.7 parts bromine, 126.5 parts 
iodine. While the gas densities of the elements are equal to 
their atomic weights (p. 70) those of the compound bodies, 
consequently amount to half that expressed by their formulas. 
From this it would follow that in equal volumes of compound 
bodies only half as many atoms or particles are present as in 
an equal volume of an elementary form of matter. In fact, 
one volume of H, containing n atoms of H, combines with 
one volume of chlorine, which, too, contains n atoms of 
Cl. n parts HC1 result, which fill two volumes; there- 
fore, there are only \ parts of HC1 contained in one volume 
of HC1: ATOMS AND MOLECULES. 
75 
nH + nCl = nHCl. 
I vol. I vol. 2 vols. 
This conclusion contradicts the general postulate (p. 71), 
derived from the physical properties, viz., that all gases, both 
simple and compound, contain the same number of gaseous 
particles in equal volumes. This contradiction, which for a 
long time prevented the adoption of the atomic volume theory 
in chemical science, is now easily solved 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 mole- 
cules of compounds consist of atoms, is obvious, since, 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. 70), that in equal gas volumes of the. halogen elements 
there is contained an equal number of atoms must be formu- 
lated somewhat as follows: In equal volumes of all gases is 
found an equal number of molecules (law of Avogadro). 
The process of the combination of hydrogen with chlorine 
(and the other halogens) must be conceived therefore 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 HC1: 
Hi, 4- Cl> = 2HC1. 
We can now understand that hydrogen chloride contains 
just as many molecules in an equal volume as H and Cl. This 
is apparent from the following representation : 
nll2 
nCl2 
nllCl 
nllCl 
In a similar manner 2 volumes H (containing 2n molecules) 
give with 1 volume oxygen (containing n molecules), 2 vol- 
umes aqueous vapor; consequently, 2n molecules of water. 
In 2n molecules of the latter (H20) there are contained 2n 
atoms of O ; therefore in n molecules of oxygen, 2n atoms of 
oxygen—or one oxygen molecule consists 0/2 atoms. 
i volume. 
i volume. 
2 volumes. 76 
INORGANIC CHEMISTRY. 
nH2 
n02 
Yield 
nH20 
nH20 
nH2 
i vol. 
2 Vols. 
2 vols. 
In the same way it may be shown that the nitrogen mole- 
cule consists of 2 atoms of nitrogen (N2), the phosphorus 
molecule, of 4 atoms of phosphorus (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), and by the remarkable action 
of the elements in the moment of their liberation. 
Upon p. 50 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 explained. 
The free elements (their molecules) are compounds of similar 
atoms whose chemical affinity has always been partially satis- 
fied. In the moment of their separation from compounds free 
atoms appear, which, before they combine 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. ’ The same number of molecules is contained in equal 
volumes of all gaseous and vapor-forming bodies. Therefore, 
the gas densities bear the same ratio to each other as the mole- 
cular 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, their molecular 
weights. The atomic weights are compared with H — 1, ATOMS AND MOLECULES. 
77 
therefore, the densities of the elements whose molecules con- 
sist of two atoms, are equal to the atomic weights : 
Atoms. 
Molecules. 
Density. 
H 
_ j 
H2 
2 
I 
Cl 
= 35-37 
Cl, 
= 
70.74 
35-37 
Br 
= 79-7 
Br2 
= 
159-4 
79-7 
I 
= 126.5 
I2 
= 
253 
126.5 
HC1 
= 
36.37 
18.18 
HBr 
= 
80.7 
40.3 
HI 
= 
127-5 
63-7 
O 
= I5-96 
o2 
= 
31.92 
15.96 
H20 
= 
17.9 
8.9 
N 
= 14 
N2 
= 
28 
14 
nh3 
= 
17 
8-5 
P 
= 3i 
P4 
= 
124 
62 
ph3 
= 
34 
‘7 
A simpler deduction, that the molecules of the elements consist of two 
or more atoms, is the following: We proceed from the law of Avogadro, 
that an equal number of molecules is contained in equal volumes of all 
gases or vapors. This law, or better, hypothesis, cannot be proven mathe- 
matically, as was attempted; just as little as any other fundamental hy- 
pothesis*—but it possesses, as basis of the entire recent kinetic theory 
of gases, a high degree of probability. It necessarily follows from this 
law that the molecular weights of all bodies are proportional to the gas 
densities. Referred to hydrogen as unit, the empirical gas densities of 
HC1 = 18.18, of HBr = 40.3, of HI == 63.7, etc. Analysis shows, how- 
ever, that 35-37 parts of Cl are in union with 1 part II in HC1, 79.7 bro- 
mine in HBr, 126.5 iodine in HI. As the weight of one atom of H is 
made equal to 1, and 35.37 parts of chlorine are combined with it, the 
weight of a molecule of hydrogen chloride, consisting of at least one atom 
of H and one atom of Cl, must equal 36.37 ; it is, therefore, twice as much 
as its density, 18.18. Hence the molecular weights of all other bodies, as 
they bear the same ratio as the densities, must also be twice as large (referred 
to H as unit) as the latter. The hydrogen molecule is = 2, and consists 
of two atoms,as its atomic weight equals 1. The chlorine molecule weighs 
70.74 units, and consists of two atoms (CI2), if we suppose that the atomic 
weight= 35.37. Its atomic weight could, however, be only the half (or 
another sub-multiple) of 35.37; then its molecule would consist of four 
chlorine atoms (Cl* = 70.74 when Cl is made equal to 17.6), and the for- 
mula of hydrogen chloride would be IICI2. From the densities of the 
elements in gas form we only ascertain their molecular weights. Their 
atomic weights are derived from the molecular weights of their compounds, 
as we regard the smallest quantity of the element which analysis discloses 
* A mathematical proof is only possible upon the basis of another, more 
general, quantitative hypothesis (or of an axiom), which then on its side 
is not to be adduced. INORGANIC CHEMISTRY. 
in the molecule of any compound as the atomic weight. Thus, in the mole- 
cule of any compound of chlorine there are never less than 35.37 parts by 
weight of Cl. That the maximum values thus derived have not been found 
too high, but correspond to the actual relative atomic weights, follows 
from the agreement of these numbers with the atomic numbers obtained 
from the specific heat of the elements. The complete certainty of their 
correctness we reach by the law of periodicity, which is 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 (H,2), 
would occupy two volumes. We say, therefore, although in- 
correctly, 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 equal gas volumes there is contained an equal 
number of molecules. 
These convincing suppositions and conclusions deduced 
from these actual relations, form the atomic molecular doctrine, 
which is the foundation of the chemistry of to-day. As this 
doctrine completely explains the quantitative phenomena 
arising in the action of the chemical elements upon each 
other, and as it has been repeatedly confirmed by entirely op- 
posite phenomena, it is only proper and correct that it be 
designated a theory (p. 67). 
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. 
OXYGEN GROUP. 
o .= 15.96. 02 = 31.9. 
x. OXYGEN. 
Oxygen (oxygenium) is the most widely distributed element 
in nature. It is found free in the air ; in combination it exists 
in water. It is an important constituent of most of the min- 
eral and organic substances. 
It was discovered, almost simultaneously, by Priestley, in 
England, 1774, and Scheele, in Sweden, 1775. Lavoisier, in 
France, 1774-1781, first explained the important role attached 
to oxygen in processes of combustion, of respiration, and of 
oxidation. oxygen. 
79 
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: 
Hg0=Hg + 0. 
The following method is, commonly pursued in the chemi- 
cal laboratory: Potassium chlorate, a compound of potassium, 
Fig. 51. 
chlorine and oxygen, is heated in a glass retort (Fig. 51) or 
flask, and thus decomposed into solid potassium chloride and 
oxygen: 
KCIO3 = KC1 + 3O * 
The evolution of the gas proceeds more regularly and re- 
quires a less elevated temperature if the pulverized chlorate be 
mixed with ferric oxide or manganese peroxide. The liberated 
oxygen is collected over water. 
Very pure oxygen may also be obtained by heating potas- 
sium dichromate with sulphuric acid : 
K2Cr207 + 4H2S04 = Cr2(S04)s + K2S04 + 4H20 + 3O. 
Besides these, many other methods may be employed for 
the preparation of the gas: e.g., the ignition of manganese 
* The chemical equations used here and previously are only intended 
to represent the manner of the reaction, and to express the accompanying 
relative quantities by weight. It should not be forgotten that free atoms 
do not exist, but that they always occur combined in molecules. Molecu- 
larly written the equation would be: 
2 KCIO3 = 2 KC1 + 3 O2. 80 
INORGANIC CHEMISTRY. 
and barium peroxides; the decomposition of sulphuric acid 
at a high heat; the boiling of a solution of bleaching lime 
with a cobalt salt, etc. These methods, applied technically, 
will be considered more fully later. 
Properties.—Oxygen is a colorless, odorless, tasteless gas, 
which is condensed to a transparent liquid, of specific gravity 
0.978, at —130°, under a pressure of 470 atmospheres. Its 
density equals 15.96 (H = 1), or 1.1060 (air =). One litre 
of oxygen at o° C., and 760 mm. pressure, weighs 1.4296 
grams, (15.96 times more than one litre of hydrogen). It is 
only slightly soluble in water; 100 volumes of the latter dis- 
solve 4.1 volumes of the gas at o°, and 2.9 volumes at 150. It 
is more readily dissolved by absolute alcohol (28 volumes in 
100 volumes). 
Oxygen combines with all the elements excepting fluorine. 
With most of them it unites directly, accompanied by the evo- 
lution 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 phe- 
nomena of the respiration of animals are also influenced by 
the contact of the oxygen of the air—hence the earlier desig- 
nations of oxygen as inflammable air, and vital air. In pure 
oxygen the phenomena of combustion proceed more energeti- 
cally. Ignited charcoal or an ignited sliver inflame immedi- 
ately in the gas, 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). 
Fig. 52. 
Fig. 53. 
Even iron is able to burn in the gas. To execute this experi- 
ment, 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 OXYGEN. 
81 
exceedingly intense light and emits sparks. (To protect the 
vessel from the fusing globules of iron oxide, cover the bottom 
with a lajcr of sand.) Iron will burn in any flame if a cur- 
rent 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. 45). 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 
hydrogen chloride (p. 54). A mixture of hydrogen and oxy- 
gen detonates violently; most strongly if the proportions are 
1 volume of oxygen and 2 volumes of hydrogen ;• such a mix- 
ture is known as oxy-hydrogen gas. The explosibility may be 
shown in a harmless way by the following experiment: Fill a 
narrow-necked flask of 4-6 ounces, over water, 2/i with hydro- 
Fig. 54. 
gen, and 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 
complete breaking of the flask. 
The oxy-hydrogen flame is only faintly luminous; it pos- 
sesses, however, a very high temperature, answering, therefore, 
for the melting of the most difficultly fusible substances, e.g., 
platinum. 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; 
oxygen is conveyed through the inner A, and the mixture ig- 
nited at a. Such a flame impinging on a piece of burnt lime 
makes the latter glow and emit an extremely bright light— 
Drummond's LimeLight. 
The union of oxygen with other substances, is termed oxida- 
tion. This term, as well as the name oxygenium (from 
and yes-wzto), or acid producer, arises from the fact that acids 82 
INORGANIC CHEMISTRY. 
are sometimes formed in oxidation. This the combustion ex- 
periments, previously mentioned, 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, and HI, contain no oxygen. Some of the ele- 
ments yield acids by their union with oxygen, or more cor- 
rectly oxides, which form acids with water. Most of these 
are the metalloids. Thus the following corresponding acids 
are derived from the acid-forming oxides of sulphur and phos- 
phorus : 
Sulphur 
trioxide. 
S03 + H20 = h2so4 
Sulphuric 
acid. 
p2o6 + h2o = 2HPO3 
Phosphorus 
pentoxide. 
Metaphosphoric 
acid. 
With oxygen the metals usually yield oxides, which form 
hydroxides (hydrates) or bases with water: 
Pot. 
oxide. 
K20 + HaO = 2KOH 
Potas. 
hydroxide. 
CaO + H20 = Ca (OH), 
Calcium 
oxide. 
Calc, 
hydroxide. 
The salts are produced by the alternating action of acids 
and bases (see p. 58). 
Thirdly, there exist the so-called indifferent oxides, which 
yield neither acids nor bases, with water, e.g., 
N20 
NO 
Ba02 
Nitrous 
oxide. 
Nitric 
oxide. 
Barium 
peroxide. 
Oxidation is not only induced by free oxygen or bodies rich 
in it, but, frequently, also, by the halogens ; in the latter case 
the halogens first decompose the water with the elimination of 
oxygen, which then oxidizes further (compare p. 49). 
The opposite of oxidation, the removal of oxygen, is called 
reduction. Hydrogen (in statu nascendi), and substances giv- 
ing it off easily (as HI), have a reducing action. Most of the 
metallic oxides are reduced at a red heat, by hydrogen, e.g. : 
• CuO + Ha = Cu + II20. 
Copper oxide. 
Copper. OZONE. 
83 
OZONE, Og.. 
Ozone, discovered in 1840, by Scbonbein, is a peculiar 
modification of oxygen, characterized by a remarkable odor 
and great ability to react, therefore it is called 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; thus, in the decomposition of peroxides by con- 
centrated sulphuric acid, in the electrolysis of water (at the 
positive pole), in the slow oxidation of moist phosphorus, in 
the combustion of hydrocarbons, and in the action of* the 
so-called silent discharge in an atmosphere of oxygen or fir. 
In none of these instances is all the oxygen ever converted 
into ozone; only a small portion—in most favorable conditions 
5-6 per cent, this change. 
The following methods serve for the preparation of ozone: 
I. Bring several pieces of stick phosphorus into a spacious flask, cover 
them about half with water, and allow them to stand for some hours. Or 
conduct oxygen over pieces of phosphorus placed in a glass tube and 
Fig. 55. 
moistened with water. Ozone is also formed abundantly when a potas- 
sium bichromate solution is substituted for water. 
2. Pass the electric spark from an electrical machine or a Ruhmkorfif 
coil through air or oxygen. The silent discharge from a powerful induc- 
tion current is better.- For this purpose we can employ a Siemen’s induc- 
tion tube (Fig. 55) which consists of a glass tube covered without with tin 
foil, in the interior of which is a smaller tube coated upon its inner 
side. The oxygen circidates between the two tubes; the two coatings are 
in connection with the induction spiral, or the poles of a Iloltz electric 
machine. 84 
INORGANIC CHEMISTRY. 
3. Gradually add barium peroxide in small portions (or potassium per- 
manganate) to cold sulphuric acid: 
Ba02 + H2S04 ;= BaSQ4 -f H20 + O. 
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. In a long layer, ozone shows a bluish color. 
Ozonized air, subjected to powerful pressure (150 atmospheres) 
at*avery low temperature, yields drops of an indigo-blue color; 
these appear to be pure ozone. Ozone is rather stable at the or- 
dinary temperature ; when heated to 300° C., it reverts to ordi- 
nary oxygen. It is somewhat soluble in pure water ; the larger 
portion of it is, however, converted by the water into oxygen, 
without formation of hydrogen peroxide. Unlike ordinary oxy- 
gen, ozone, especially in a moist state, oxidizes strongly at 
ordinary temperatures. Phosphorus, sulphur, and arsenic are 
converted into phosphoric, sulphuric, and arsenic acids ; ammo- 
nia is changed to nitrous and nitric acid ; silver and lead are 
converted into the corresponding peroxides ; therefore paper 
moistened with a lead salt is colored brown. Iodine is sepa- 
rated from potassium iodide by it: 
2Ki 4- h2o + o = 2KOH + 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 ordinary potassium iodide starch paper 
(Schonbein) 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. The 
quantity of ozone may be approximately determined from' the rapidity and 
the intensity of the coloration; the reactive power is, however, very much 
influenced by aqueous vapor. Thallous hydrate is a-more reliable reagent 
for ozone than the potassium-iodide paper. Guaiacum tincture and paper 
saturated with a lead acetate solution may also be used to detect ozone; 
the first acquires a blue color, the second is browned. Other substances 
also blue potassium iodized starch and guaiacum, eg., chlorine, bromine, 
niti'ogen dioxide, etc., etc. To distinguish ozone from these, proceed as 
follows ( Houzeau): Take two strips of violet litmus paper, one of which is 
saturated with KI, and expose it to the action of the gas; when O3 is present OZONE. 
85 
potassium hydroxide will be formed from the KI, and color the violet lit- 
mus blue. The second paper serves to show the absence of ammonia. 
The preceding reactions of ozone are all produced by hydrogen peroxide, 
although less rapidly. The only test answering for the distinction of very 
slight quantities of ozone from hydrogen peroxide, is the blackening of a 
bright strip of silver by ozone. 
Ozone is formed from pure oxygen, and is nothing more 
than the latter condensed. The molecules consist of 3 atoms 
of O: 
3 Vols. oxygen. 
302 
yield 
2 vols. ozone. 
2O3. 
This is proved by the following experiments: In ozonizing 
oxygen its volume diminishes ; upon heating (whereby ozone 
is again changed to oxygen), the original volume is reproduced ; 
when ozonized oxygen is brought in contact with oil of tur- 
pentine or cinnamon, all the ozone is. absorbed and the vol- 
ume of the gas is 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 that the first is twice as large as the 
latter; this indicates that 1 volume of ozone yields 1 y2 vol- 
umes of oxygen. From this it follows that the specific gravity 
of ozone must be times greater than-that of ordinary oxy- 
gen, and that if the molecule of O consists of 2 atoms, the 
molecule of ozone must contain 3 atoms. This conclusion is 
confirmed by the specific gravity of ozone derived experiment- 
ally from the velocity of diffusion. The density of ozone is 
found to be 24 (H =z 1); the molecular weight of it, therefore, 
is 24 x 2 = 48, a number almost equal to the trebled at- 
omic weight of oxygen (3 X 15.97 — 47-9)- The molecular 
formula of ozone is, therefore, 03. 
A diminution in the volume of the gas does not occur in 
the action of ozone upon oxidizable bodies like KI and Hg, 
although all the ozone disappears. It would appear from this, 
that in oxidizing, ozone only acts with one atom of oxygen, 
while the other two atoms form free oxygen, which occupies 
the same volume as the ozone : 
O3 2KI — 02 -j- K20 I2. 
I vol. 
I vol. 
As a consequence of this behavior, ozone is also called oxid- 
ized oxygen ; i.e., free, oxygen (02), which has combined with 
an additional oxygen atom. 86 
INORGANIC CHEMISTRY. 
We observe, therefore, that the elementary substance oxygen 
occurs in free'condition in two different forms—allotropic 
modifications—ordinary oxygen (02) and ozone (03). We 
will learn later that very frequently substances of the same 
elementary composition possess different physical and chemical 
properties ; such bodies are called isomerides and the phenom- 
enon 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 chem- 
ical union of two bodies the particles of matter would entirely 
permeate and blend into each other, the existence of isomeric 
bodies would scarcely be 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, however, only very slowly oxidized; the 
ozone is contained in them in a peculiar, combined condition. 
In this form it acts upon some bodies like free ozone ; in 
other instances, the oxidizing action is only rendered possible 
by peculiar substances which carry the ozone. Spongy pla- 
tinum, ferrous sulphate, and the blood corpuscles are examples 
of this class. Thus, old turpentine oil, containing absorbed 
ozone, only acts on paper saturated with starch and potassium 
iodide, if a few drops of a ferrous sulphate solution have been 
added to it. 
. Since ozone is formed when electricity acts upon air, and 
indeed, probably, in all oxidation and combustion processes; 
and, 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 (i-io milligrams in ioo litres of 
air); according to recent investigations it is, however, prob- 
able that the imagined ozone reactions are frequently produced 
by hydrogen peroxide, which is very similar to ozone in re- 
action (p. 85), and is almost constantly in the air (Schone). 
Antozone, which was regarded as a third peculiar modifi- 
cation of oxygen, has been proved to be hydrogen peroxide. 87 
WATER. 
COMPOUNDS OF OXYGEN WITH HYDROGEN. 
H20 = 17.97. Density = 8.98. 
1. WATER. 
Water, the product of the union of hydrogen with oxygen 
(p. 81), is produced in many chemical processes, e.g., in the 
formation of salts from bases and acids (p. 58). 
Cavendish was the first (1781) to confirm the formation of 
water by the combustion of hydrogen. * Lavoisier first (1783) 
determined its quantitative composition. Later (1805) Gay- 
Lussac showed that it was produced by the union of two 
volumes of hydrogen with one volume of oxygen. 
Physical Properties.—It is obtained chemically pure 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 great- 
est density at -f- 40 C., the maximum contraction. The weight 
of a cubic centimeter of such water is taken as the unit of 
weight (= i gram). By further cooling the water expands— 
the opposite of most other bodies; its volume becomes 
greater, while.the specific gravity decreases. 
The following table gives the volume and specific gravity 
of wat6r for different temperatures (according to Kopp): 
Temperature. 
Volume. 
Specific Gravity. 
o° 
1.00012 
O.99988 
2° 
1.00003 
O.99997 
4° 
* I.OOOOO 
I.OOOOO 
6° 
I.OOOO3 
0.99997 
8° . 
I.OOOI I 
0.99989 
IO° 
I.OOO25 
0.99975 
12° 
I.OOO44 
0.99956 
i4° 
1.00068 
0.99932 
16° • • 
I.OOO97 
0.99903 
18° 
I.OOI3I 
0.99869 
20° 
I.OO169 
0.99831 
22° 
1.002X2 
0.99789 
240 
I.OO259 
0.99742 
By cooling water solidifies to ice. The solidification-tem- 
perature of water, or more correctly the fusing point of ice, 
is taken as the zero of Celsius’s and Reaumur’s thermometric 88 
INORGANIC CHEMISTRY. 
scales. We can, however, reduce still water considerably be- 
low the o° point without its freezing, whilst the fusing point 
of ice, like all other solid bodies, is constant (at a definite 
pressure). 
Critical Pressure.—Carnelley has recently made the interesting obser- 
vation of the dependence of the fusibility of bodies upon pressure. Just 
as there exists a definite temperature, for every gas (the critical tempera- 
ture, p. 42) beyond which it cannot be condensed, so is there a minimum 
pressure below which solid bodies are no longer fusible, but vaporize 
without fusion. If ice be placed in a vessel that is exhausted under an 
air-pump, it will be discovered that under a pressure of 4.6 mm. it ceases 
to melt", but vaporizes directly. The critical pressure of ice is, therefore, 
4.6 mm. This deportment is better observed in the case of mercuric 
chloride, HgCl2, whose critical temperature lies near 420 mm. It was 
known long ago that some bodies (like arsenic) were at once vaporized 
w'hen heated in the air or in a vacuum, and that they could only be melted 
in sealed tubes : their critical pressure, therefore, lies above 760 m. It 
has only recently been established that fusibility is universally influenced 
by pressure. 
Ill the conversion of water into ice, a considerable expan- 
sion occurs: ioo vols. H20 at o° yield 107 vols. ice; the 
specific gravity of the latter is, therefore, 0.93. Ice crystal- 
lizes in hexagonal forms, as may be distinctly observed in 
snow-flakes. 
Different bodies require different quantities of heat to bring 
them to the same temperature. The heat capacity of water 
is greater than that of all other liquid or solid bodies. It 
is customary to take the quantity of heat necessary to raise 
one part by weight of H20 from o° C., to i° 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 
equals 79 calories; that means, that for the fusion of one part 
of ice bv weight, a quantity of heat is required which is capable 
of raising one part H20 from o° to 790 C. 
Water boils upon the application of heat, 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 = ioo° of the thermometric scale of Celsius (=8o° 
Reaumur). WATER. 
89 
One volume of water, at ioo° C., yields 1696 volumes of 
vapor of the same temperature. The specific gravity of steam 
__ 17.97 __ g or = o.622(air = 1). One 
2 14-43 
litre of aqueous vapor weighs 0.8064 grams (at o°). 
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, which holds 
it in eqnilibrio. 
The following table gives the tension of aqueous vapor for various 
temperatures: 
Temperature. 
Tension. 
Temperature. 
Tension. 
—20° C. 
0.93 mm. 
40° C. 
54 9 mm. 
—io° C. 
2.09 mm. 
6o° C. 
148.8 mm 
* o° C. 
4.0 mm. 
8o° C. 
354.6 mm. 
+ io° C. 
9.1 mm. 
1 oo° C. 
760.0 mm. 
20° C. 
17.4 mm. 
120° C. 
1491.0 mm. 
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 tem- 
perature. (Compare p. 120.) 
A definite quantity of heat, requisite for the conversion of 
a liquid into a 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 
ioo° C.; i e., for the conversion of one part of water of ioo° 
C., into vapor of the same temperature, a quantity of heat 
will be absorbed capable of raising 536.5 parts of H20 from 
o° to i° will be absorbed. 
In consequence of the evaporation of water, the gases sep- 
arating 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 40). 
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 contain 90 
INORGANIC CHEMISTRY. 
foreign admixtures. The purest natural water is rain and snow 
water ; it contains upwards of 3 per cent, by volume of gases 
(oxygen, nitrogen, and carbon dioxide), and traces of solids 
(the ammonium salts of nitrous and nitric acids). If water 
that has been standing exposed to the air be heated, the dis- 
solved gases escape in bubbles. 
River and spring waters contain, on an average, from 1 to 20 
parts of solid constituents in 10,000 parts. Water having much 
lime and gypsum present in it, is ordinarily known as hard, in 
distinction from soft water, which contains less lime (see Cal- 
cium Carbonate). On boiling lime waters, most of the impurity 
deposits out. Spring water generally contains in addition 
larger quantities of carbon dioxide, which impart a refreshing 
and enlivening taste to it. Spring waters holding considerable 
quantities of solid constituents, or exhibiting special healing 
properties, are called mineral waters. These are distinguished as 
saline waters (containing sodium chloride), sulphur waters 
(hydrogen sulphide), acidulated waters (saturated with carbon 
dioxide), chalybeate waters (containing iron), and others. 
Sea water contains about 3.5 per cent, of salts, of which 
2.7 per cent, are sodium chloride. 
To purify the natural waters they are filtered (for the re- 
moval of mechanical admixtures), and for chemical purposes, 
distilled (distilled water) in apparatus of varying form. 
.Solutions.—The phenomena appearing in the dissolving of 
substances indicate that solutions are not mere mechanical 
mixtures. In every solution alterations occur in the tempera- 
ture of the liquid. The solubility of solid and liquid sub-* 
stances increases usually with the temperature, while that of 
gases diminishes. The quantity of dissolved gas is frequently 
proportional to the pressure ; other gases, on the contrary, 
which are readily soluble in water, such as the halogen-hydro- 
gen compounds, are exceptions to this rule. Heat does not 
completely expel them from their solution ; they distil over as 
liquids of definite composition (compare pp. 56 and 63). When 
they dissolve, a large quantity of heat is liberated, just as in the 
case of chemical compounds. Further, a contraction is always 
perceived in the solution of solids and liquids; the volume of 
the solution is less than the sum of the volumes of the con- 
stituents. These phenomena point to the acceptance of a cer- 
tain affinity between the dissolving bodies. Therefore, solu- 
tions, like alloys, are designated undetermined compounds, 
in contrast to the determined compounds, which are combined 
according to constant atomic weight ratios. This view is also WATER. 
91 
confirmed by the fact that frequently definite compounds con- 
taining water do exist in solution. Such compounds often 
separate, unaltered, when their solutions are evaporated; the 
water present in them is known as water of crystallization. It 
is, however, impossible to draw a sharp line between deter- 
mined and undetermined compounds, between chemical and 
physical attraction. 
The thermal phenomena, which appear when solution occurs, bear a close 
relation to chemical affinity. The halogen hydrides, easily soluble in 
water, disengage large quantities of heat in their solution, corresponding to 
the symbols: 
(HC1, aq.) = 17320; (HBr, aq.) = 19940; (HI, aq.) = 19200. 
This liberation of heat is explained by the production of the hydrates 
HC1 -)- 8H20, HBr -f- 5H20, HI -f- 5H20, which distil over unaltered (p. 
61). 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. Those forming hydrates 
liberate heat, while the non hydrate-forming solid bodies absorb heat in 
their solution, which at the same time is employed to liquefy them (latent 
heat of fusion). Thus, in the solution of the halogen compounds of po- 
tassium the following quantities of heat are absorbed: 
(KC1, aq.) = — 4400; (KBr, aq.) = — 5080; (-KI, aq.) = 5100. 
The freezing mixtures, to be described later, depend upon such an ab- 
sorption of heat. 
Chemical Properties of Water.—Water is a neutral substance, 
i.e., it possesses neither acid nor basic properties. As we have 
already observed (p. 82), 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 + K2 = 2K0H + h2. 
Other metals do not decompose it, except at elevated tem- 
peratures. Steam conducted over ignited iron gives its oxygen 
to the latter, forming ferroso-ferric oxide, while hydrogen is 
set free: 
3Fe + 4H20 = Fe304 -f 4H2. 
Chlorine decomposes water in the sunlight; the decompo- 92 
INORGANIC CHEMISTRY. 
sition is more rapid when the vapors are conducted through 
heated tubes: 
H20 + Cl, = 2HC1 + o. 
The electric current separates acidulated water* into its con- 
stituents, oxygen and hydrogen ; the first collecting at the 
positive, and the latter at the negative pole. The oxygen thus 
obtained contains ozone. 
The deportment of water at high temperature is very in- 
teresting. On pouring molten platinum into cold water, bub- 
Fig. 56 
bles of oxy-hydrogen gas escape. A similar decomposition of 
water occurs when it is led through porcelain tubes raised to 
a white heat. 
From the observations of Sainte-Claire Deville, it was found 
that the decomposition of H20 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 
temperature is known as dissociation. The following experi- 
ment illustrates this: Pass aqueous vapor through a porous 
clay tube, a, puttied into a wider non-permeable porcelain 
tube heated to a white heat in an oven (Fig. 56). The water 
suffers partial 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 
mainly through the inner tube at a. A part of the same dif- 
fuses simultaneously with the hydrogen and reunites with the 
latter. To avoid this, conduct a stream of carbon dioxide 
* Pure water appears not to be decomposed by the galvanic current. WATER. 
93 
through the wider porcelain tube; this will carry out the hy- 
drogen 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 temperature. A platinum tube may be advantageously 
substituted for the porous clay tube, because only hydrogen 
will pass through it (p. 43). 
Many other compounds, like ozone, ammonium chloride, 
phosphorus pentachloride, carbon dioxide, etc., suffer a similar 
partial decomposition when heated; they are decomposed into 
simpler molecules. 
The explanation of the dissociation phenomena is found in the kinetic 
theory of gases and heat. According to it, not only the gas molecules 
have a direct oscillating movement, inasmuch as they rebound from each 
other, like elastic balls, but even the atoms in the molecule possess heat 
vibrations. The velocity of the oscillations of molecules and atoms in- 
creases with augmented temperature ; it is, therefore, 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 conse- 
quence 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 tempera- 
ture, yield, therefore, to the decomposition. From this we discover that 
the dissociation is gradual and increases with the temperature. The law 
of dissociation is expressed by the curve of probability. 
Dissociation, i. e., the partial decomposition, increasing 
with the temperature, explains many chemical processes that 
before appeared very obscure, e.g., the mass action in reversed 
chemical reactions. We have already said that iron raised to 
a red heat decomposed water with the separation of hydrogen 
and the production of ferrous-ferric oxide. On conducting 
H over ignited iron oxides the opposite process occurs; the 
oxygen compound of the iron is reduced and water is formed: 
Fe304 -f 4H2 = 3Fe + 4HaO. 
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 Fe,04, Fe, 
H20 and H2 occur simultaneously. Such a state occurs in every dissocia- 
tion. 94 
INORGANIC CHEMISTRY. 
THE QUANTITATIVE COMPOSITION OF {WATER. 
The composition of water by weight is best determined by 
a synthesis of the same; this may be done by reducing cupric 
oxide with hydrogen: 
Cupric oxide. 
CuO 4- H2 = Cu + H20. 
Copper. 
Heat a weighed portion of cupric oxide (containing a defi- 
nite amount of oxygen), in a stream of pure, dry hydrogen, 
and weigh the quantity of H20 obtained. The operation can 
Fig. 57. 
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 com- 
pletely absorbed in the tube H. Hydrogen is led over the 
cupric oxide until it is reduced to red metallic copper, then 
allowed to cool, when F is weighed alone and G and H to- 
gether. The loss in weight of the quantity of oxy- 
gen which has combined with hydrogen to produce water. 
The increase in weight of G and H gives the quantity of 
water that was formed. The difference shows the amount of H WATER. 
95 
in water. Thus we ascertain that in 100 parts of water, by 
weight, there are: 
11.136 Parts Hydrogen. 
88.864 “ Oxygen. 
100.00 “ Water. 
Or, i part hydrogen and 7.98 parts oxygen yield 8.98 parts 
water. 
THE MOLECULAR FORMULA OF WATER. ATOMIC WEIGHT OF 
OXYGEN. 
If the molecule of water (like HC1) contains i atom H and 
i atom oxygen, then its chemical formula would be HO, and 
the atomic weight of oxygen would be = 7.98. Such a sup- 
position has, however, not been proved by any facts. 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 3.99. According to the formula H20, the atomic weight 
of O would be 15.96, etc. (see p.69). 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 69-78. In equal volumes of the gases (or 
vapors) there is an equal number of molecules. The molecular 
weights, therefore, are proportional to the gas densities, and 
are equal to double that of the densities referred to H =1. 
The density of steam is 8.98 (H = 1); the weight of the water 
molecule is therefore 17.96. Analysis, however, shows that in 
17.96 parts water 2 parts by weight are hydrogen (= 2 atoms) 
and 15.96 parts oxygen by weight. According to this, the 
molecule of water contains not more nor less than 2 atoms of 
hydrogen. That the 15.96 parts of oxygen combined with the 
latter correspond to one atom (that the atomic weight does not 
■ equal the half, in which case the molecular formula of water 
would be H202), follows from the fact that the analysis of none 
of the innumerable oxygen derivatives has shown less than 15.96 
parts oxygen in the molecule (see p. 77). The molecular for- 
mula of water, therefore, is H20 = 17.96. The gas density 
of oxygen is 15.96, the molecular weight 31.92, therefore the 
oxygen molecule consists of 2 atoms, 02 = 31.92. 
After having thus derived the molecular formula of water, 
and the atomic weight of oxygen, we deduce the following 
conclusions: 
(1) 15.96 parts of O by weight, occupy the same volume as 96 
INORGANIC CHEMISTRY. 
i part of H by weight; since 15.96 parts of the former unite 
with 2 parts of the latter in the production of water, 1 volume 
of O must combine with 2 volumes of H. 
(2) In equal volumes we have an equal number of molecules ; 
n molecules of oxygen (02) unite therefore with 2 //molecules 
of hydrogen (H2) ; the same yield 2 n molecules of water ; con- 
sequently 2 volumes of aqueous vapor: 
2 Vols. 
2nH2 + n02 = 2nH20. 
I vol. 
2 Vols. 
According to the above, 2 volumes of hydrogen and 1 volume 
of 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 8.98, and 2 parts of hydrogen by 
weight unite with 15.96 parts of 
oxygen by weight to form 17.96 
parts of water; then, the latter, 
in the form of vapor, must occupy 
two volumes. Conversely from 
these volume ratios it is shown 
that the molecule of oxygen con- 
sists of two atoms (compare 
P- 75)- 
These conclusions are con- 
firmed by the following experi- 
ments : 
1. When water is decomposed by the 
electric current in a voltameter, or, more 
suitably, in Hofmann’s apparatus (Fig. 
48, p. 73), it will be found that the vol- 
ume of the separated hydrogen is double 
that of the oxygen. This can also be 
proved synthetically. Introduce 1 vol- 
ume of oxygen and 2 volumes of hydro- 
gen into an eudiometer tube filled with 
mercury (see Air), and let the electric 
spark pass through the mixture. This 
will unite the two gases, a small quan 
tity of water forming at the same time; 
all the gas has disappeared, and the tube fills perfectly with mercury. In 
place of the eudiometer the following apparatus (Fig. 58) may be advan- 
tageously employed in this experiment (and also in many others). It con- 
sists of a U-shaped glass tube, one limb of which, open above, is provided 
below with an exit tube. The other limb really represents an eudiometer; 
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 mercury, and run into the 
Fig. 58. HYDROGEN PEROXIDE. 
97 
eudiometer limb 1 volume O and 2 volumes H. 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 th,e explosion, to convert it by heat into 
steam. The subjoined apparatus, will answer for this purpose (Fig. 59). 
Fig. s9. 
This is essentially the same as that pictured in Fig. 58, with the eudiometer 
limb closed above and surrounded by a wider tube. Through the latter 
conduct the vapors of some liquid boiling above ioo° C (aniline). These, 
then, pass through the envelope B, and are again condensed in the spiral 
tube C. The quantities of H and O used are heated to the same tempera- 
ture, 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 % of the volume of the gas mixture; and 3 volumes of 
oxv-hydrogen gas yield 2 volumes of aqueous vapor. 
The composition of water by weight may be easily deduced, knowing 
the specific gravities of hydrogen and oxygen and the ratios in which they 
unite by volume: 
1 volume of oxygen equals 15.96 parts by weight. 
2 volumes of hydrogen equal 2 parts by weight. 
The resulting H20 equals 17.96 parts by weight. 
17.96 parts water, therefpre, contain 15.96 parts oxygen and 2 parts 
hydrogen, or, in 100 parts there are 88.86 parts oxygen and 11.14 parts 
hydrogen. 
2. HYDROGEN PEROXIDE. 
In addition to water, oxygen forms another compound with 
hydrogen, known as hydrogen peroxide. It is produced by 
H202 = 33.92. 98 
INORGANIC CHEMISTRY. 
the action of dilute acids upon certain peroxides, such as those 
of potassium, calcium and barium. It is most conveniently 
obtained by the action of hydrochloric acid upon barium 
peroxide : 
Barium 
peroxide. 
Ba02 + 2HC1 = BaCl2 + H202. 
Barium 
chloride. 
Barium peroxide made to a paste with a little water (better the hydrate 
—see Barium) is introduced gradually, in small quantities, into coid hy- 
drochloric acid, diluted with three volumes of water. Hydrogen peroxide 
and barium chloride result; both are soluble in water. To remove the 
second from the solution, add to the latter a solution of silver sulphate as 
long as a precipitate is formed. Two insoluble compounds, barium sul- 
phate and silver chloride, are produced by this reagent: 
Remove the precipitate by filtration and concentrate the aqueous solution 
under the air pump. It now contains only hydrogen peroxide. 
BaCl2 + Ag2S04 = BaS04 + 2AgCl. 
In making the peroxide, carbon dioxide may be allowed to 
act on barium peroxide suspended in water: 
Ba02 -f C02 + H20 = BaC03 + H202. 
The insoluble barium carbonate is filtered off and the filtrate 
concentrated. 
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 follow- 
ing equation : 
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 liquid under the air-pump. Dry commercial 
hydrate of the peroxide of barium is not applicable for the 
above. A dilute solution of hydrogen peroxide is very readily 
prepared, if sodium peroxide (obtainable by fusing sodium in 
the air) is added to dilute tartaric acid. 
Ba02 -(- H2S04 = BaS04 -f- H202. 
Besides these, other methods exist for preparing hydrogen peroxide 
(in small quantity) ; all are dependent upon the decomposition of metallic 
peroxides. If phosphorus, covered with water, be allowed to oxidize 
(p. 82) in the air, hydrogen peroxide will be found in the water, and the 
surrounding air will contain ozone. Or, if a flask filled with air, be shaken 
with zinc and water or dilute sulphuric acid, hydrogen peroxide will be HYDROGEN PEROXIDE. 
99 
produced. It is destroyed again by the prolonged action of the zinc. Cop- 
per, lead, and other heavy metals do the same w hen agitated with more or 
less dilute sulphuric acid, and we find the same result by the oxidation of 
many organic substances, e.g., pyrogallic acid and tannin on exposure to the 
air. The explanation offered for this formation of hydrogen peroxide (and 
ozone) is, that in the oxidations, the oxygen molecules are torn asunder, 
and the nascent oxygen atoms oxidize the water to a slight degree to hy- 
drogen peroxide, and oxygen to ozone. The rare occurrence of ozone is 
due either to its difficult formation, or to the fact that it is readily decom- 
posed by the reacting bodies (zinc, etc.) This is also the case with hy- 
drogen peroxide. The appearance of hydrogen peroxide in the oxidation 
of phosphorus, seems to prove that it can be formed by the oxidation of 
water. This seems to be confirmed by its production on shaking tur- 
pentine oil with water and air, or if ozone be conducted into ether, and 
the ozonized product shaken with water (water is not directly oxidized by 
ozone). It appears probable, however, that in some oxidation reactions, 
the formation of the hydrogen peroxide is a consequence of the reduction 
of oxygen (Traube). It may, for example, be assumed that when zinc is 
shaken with air and water (or dilute sulphuric acid), the latter is decom- 
posed in such a manner that the hydroxyl group combines with the zinc to 
hydroxide, and the liberated hydrogen then yields hydrogen peroxide wdth 
oxygen : 
Zn + 2OHH + O2 =Zn(OII)2 + H202. 
A confirmation of this supposition is found in the electrolysis of water, 
where we discover H2O2 appearing at the negative pole (where hydrogen 
is found) if air or oxygen be conducted through the solution. It is veri- 
fied, too, in the production of H2O2 upon shaking palladium hydride with 
water and air: 
2Pd2H + Oj = 4Pd + H202. 
In ail these examples we can explain the formation of the peroxide by the 
action of nascent hydrogen upon oxygen. It is easily understood why 
there is no, or very little, peroxide found in the energetic evolution of 
hydrogen. Thus it arises in almost all slow oxidations in which ozone is 
produced at the same time. 
Hydrogen peroxide, 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 very 
dilute solutions, pure water freezes out. It possesses a bitter, 
astringent taste, is miscible in all proportions with water, and 
vaporizes in vacuo. Very dilute aqueous solutions can be 
boiled without decomposing the peroxide; a portion of it 
distils over with the water. 
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 INORGANIC CHEMISTRY. 
heating the point of explosion can be reached. In conse- 
quence of this ready decomposition, hydrogen peroxide oxid- 
izes powerfully, since oxygen appears (p, 76) in statu nascendi. 
It converts selenium, chromium, and arsenic into their cor- 
responding acids: sulphides are changed to sulphates (PbS to 
PbS04) ; from lead acetate solutions the peroxide is precipi- 
tated, but is again decolorized by the. excess of peroxide. 
Organic dyestuffs are decolorized and decomposed. From 
hydrogen sulphide, sulphur, from hydrogen chloride and 
iodide, chlorine and iodine are set free: 
h.2o2 + 2HI = 2H2o + I2. 
Thus hydrogen 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 potas- 
sium iodide solution, while ozone separates iodine at once ; 
but if platinum-black, ferrous sulphate, or blood corpuscles 
(see p. 86) be added to the solution, iodine immediately sepa- 
rates out, and colors added starch-paste deep blue. 
In all these cases the action of hydrogen peroxide is oxid- 
izing. Some substances, on the other hand, are reduced by 
oxygen separating at the same time; this is true of certain 
unstable oxides, peroxides, and the highest oxidations of some 
metals, like Mn207, and Cr03. Thus, argentic, mercuric, and 
gold oxides are reduced to a metallic state with an energetic 
evolution of oxygen : 
Ag20 “t" H202 — 2Ag -(- H20 -)- 02. 
Lead peroxide is changed to lead oxide : 
Pb02 + H202 = PbO + H20 + 02 
In presence of acids, the solution of potassium permanganate 
is decolorized and changed to a manganous salt (p. 101). In 
the same way chromic acid and its salts are altered to-chromic 
oxide: 
2Cr03 + 3H202 — Cr203 -f- “I- 3^2- 
Ozone and hydrogen peroxide decompose themselves into 
water and oxygen : 
o3 -f- h202 — 02 —H,0 -j- o2. 
Chlorine, in aqueous solution is oxidized to hypochlorous 
acid by hydrogen peroxide, but again reduced by an excess of 
the latter: 
ClOIi + H2Q2 = C1H + . H20 + 02. HYDROGEN PEROXIDE. 
All these reactions are generally explained by supposing that the oxygen 
atoms (also those of other elements), possess a certain affinity for each 
other; this is saturated, by their union to molecules. Those present in 
other compounds, and not firmly bound, therefore separate and unite with 
each other, and form free oxygen molecules—OO. The conclusion de- 
rived from the gas density, viz., that the molecules of the free elements 
consist of two or more atoms, is corroborated by these reactions. The 
readiness with which ozone and hydrogen peroxide react, is explained by 
their thermo-chemical behavior. In the production of ozone from oxygen, 
and of hydrogen peroxide from water and oxygen, heat is absorbed: 
conversely, equal thermal values are disengaged in their decomposition. 
Both compounds are, consequently, endothermic (see p. 65), therefore, 
little stable ; the one oxygen atom in them is very reactive. The produc- 
tion of both substances, according to the above symbols, can only be effected 
by the addition of external energy. In case of ozone this may be accom- 
plished by electricity, with hydrogen peroxide (in its production from water; 
this only occurs in small quantity) by the heat, which becomes free in the 
chief reaction. 
(02,0) = — 32,400; (H20,0) = — 23,070; 
Finally, hydrogen peroxide may be decomposed 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 this way. Such reactions, in which 
the reacting substances undergo no perceptible 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 sdver and gold upon H2O2 oxides at first 
result, but are afterwards reduced in the manner mentioned 
above, by the hydrogen peroxide. 
REACTIONS FOR THE DETECTION OF HYDROGEN PEROXIDE. 
H20., decomposes potassium iodide very slowly; in the presence of 
iron sulphate, however, iodine separates at once, and is recognized by the 
blue color it yields with starch paste. In the same way guaiacum tincture, 
in the presence of ferrous sulphate, is immediately colored blue, and an 
indigo solution is decolorized. The most characteristic test for the perox- 
ide is the following: Introduce some H2O2 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; traces of it afford an orange yellow color with 
hydrogen peroxide. 
Hydrogen peroxide is determined quantitatively by oxidation with potas- 
sium permanganate (see Manganese). The latter is added to ths solution, 
acidified with sulphuric acid until a permanent coloration occurs. The 
reaction proceeds according to the equation : 
2Mn04K 4- 3S04H2 -f 5H2Q2 = 2S04Mn + SQ4K2 + 8H2Q + 502. 102 
INORGANIC CHEMISTRY. 
Or the liquid to be examined frain 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 iodine which sepa- 
rates is then determined colorimetrically (Schone). 
Hydrogen peroxide occurs in slight quantity in the air and is detected 
in almost all rain water and in snow—but not in natural dew and frost. 
Its quantity varies from 0.05 to 1 milligram in a litre of rain. Its forma- 
tion in the air is probably induced by the action of ozone upon ammonia, 
whereby ammonium nitrite, hydrogen peroxide and oxygen result (Carius). 
Analysis shows that H202 consists of 1 part hydrogen and 
15.96 parts of oxygen; its simplest formula would therefore 
be HO. The difficult volatility of the compound, and also 
the reactions already described, cause us to believe 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 OH, called hydroxyl; these 
are combined with each other. 
S = 31.98. S2 = 63.9 (above 1000° 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. Its compounds with the metals are known 
as blendes or glances. In combination with oxygen and 
calcium, it forms calcium sulphate, the widely distributed 
gypsum. It is also present in many organic substances. 
To obtain sulphur the natural product in Sicily is arranged 
in heaps, covered with earth and then melted, or it is distilled 
from earthen retorts. To further purify this crude commercial 
product it is redistilled (in the manufactory) from cast-iron 
retorts, and when in a molten condition is run into cylin- 
drical forms—stick sulphur. If the sulphur vapors are rapidly 
cooled during distillation (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 sulphuris). 
Sulphur may be obtained by heating the well-known pyrites 
(FeS2). 
Free sulphur exists in several allotropic modifications (see 
page 86). 
i. Ordinary octahedral or rhombic sulphur exists in nature 
in beautiful, well-crystallized rhombic octahedra (Figs. 21 and 
22, p. 35). It is pale yellow, hard and very brittle; on rubbing, 
•it becomes negatively electrified. Specific gravity of this SULPHUR. 
103 
variety equals 2.05. It is difficultly soluble in alcohol and 
ether ; more readily soluble in hydrocarbons and ethereal oils. 
The best solvents are sulphur monochloride (S2C12) and carbon 
disulphide (CS2) ; 100 parts of the latter at 220 C., dissolve 
46 pai;ts of sulphur. By slow evaporation of the solutions 
sulphur crystallizes in transparent, lustrous, rhombic octahe- 
dra, like those occurring in nature. Sulphur fuses at 1x1.5° C* 
(1130 C.), to a yellow, mobile liquid, which upon further 
heating becomes dark and thick, and at‘2500 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 is obtained from the 
rhombic when the latter is heated to its point of fusion ; on 
cooling it generally assumes the monoclinic form (rhombic 
crystals separate at 90° from sulphur that has been heated 
beyond the point of fusion). The monoclinic crystals are best 
obtained as follows : Fuse sulphur 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 crucible will be covered with long, 
somewhat curved, transparent, brownish-yellow needles, or 
prisms of the monoclinic system. The same are obtained 
when a solution of sulphur in carbon disulphide is heated to 
ioo° C. in a sealed tube, and then gradually allowed to cool; 
monoclinic crystals at first separate, and later, at low temper- 
atures, rhombic octahedra. The monoclinic crystals separated 
from the solution are almost colorless and perfectly trans- 
parent. 
Prismatic or octahedral crystals may be obtained from a supersaturated 
benzene solution of sulphur, by adding small fragments of the correspond- 
ing crystals to the solution. 
This form of sulphur has a specific gravity of 1.96 and fuses 
at 1200. It is soluble in the same solvents as the rhombic 
variety. It is very unstable; the transparent prisms and 
needles become opaque and pale yellow at ordinary tempera- 
tures, and 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 rhombic modification. 
3. Soft, plastic sulphur appears to consist of two modifica- 
tions. It is obtained when sulphur heated above 230° is poured INORGANIC CHEMISTRY. 
in a thin stream into water; it then forms a soft, fusible mass, 
of a yellowish-brown color, and its specific gravity equals 1.96. 
In few days it hardens, and is converted into the rhombic 
variety. At 950 the conversion is instantaneous and accom- 
panied by the evolution of considerable heat. It is only.partly 
soluble in carbon disulphide, leaving an amorphous powder 
undissolved—amorphous insoluble sulphur. It is also produced 
when light acts upon dissolved or fused sulphur, and in the 
decomposition of the halogen-sulphur compounds by H20. 
Flowers of sulphur are for the most part insoluble in carbon 
disulphide. ioo° C. will convert the amorphous insoluble 
sulphur into the ordinary variety. 
On adding hydrochloric acid to polysulphide solutions of 
potassium or calcium, sulphur separates as a fine, white pow- 
der, known as milk mof sulphur (Lac sulphuris) : 
K2S5 + 2HC1 = 2-KC1 + 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 density of the latter 
at 500° C. has been found to equal 96 (H = 1). The 
vapor density steadily diminishes with increase of temperature 
from 700° C. onward and becomes constant at iooo° C. and 
equals 32 ; the molecular weight, therefore, is 64. Since the 
atomic weight of S (as we will see) = 32, it follows that at 
iooo0 C., the molecules of S consist of two atoms (S2 = 64 = 
32 X 2). At 500°, however, where the vapor density = 96, 
and the molecular weight 192, the molecule consists of six 
atoms(S8=6 X 32 = 192). According to this the hexatomic 
sulphur molecules dissociate (see p. 92), on further heating, 
and fall into normal diatomic molecules; the dissociation 
begins at 700° and is complete at iooo° C. Since, therefore, 
the sulphur molecules in vapor form consist of two atoms at 
very high temperatures and of six atoms at lower, we may 
assume that the molecules in the liquid and solid condition 
are more complicated, and that the various allotropic modifi- 
cations are influenced by the number of atoms contained in 
the molecules. Other solid metalloids, eg., selenium, phos- 
phorus, arsenic, carbon and silicon, occur in different modifi- HYDROGEN SULPHIDE. 
105 
cations. As yet we have no means of ascertaining the molec- 
ular size of the elements in liquid and solid conditions ; there 
is much, however, favoring the idea that when free they con- 
sist 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. When heated to 260° in the air, 
it ignites and burns with a pale bluish flame, giving sulphur 
dioxide (S02). This union with oxygen occurs gradually even 
at lower temperatures; in the dark it. is accompanied by a 
white phosphorescent flame. Nearly all the metals combine 
with it to form sulphides. By rubbing mercury, flowers of 
sulphur and water together, we obtain black mercury sulphide. 
A moist mixture of iron filings and sulphur glows after a time. 
Cu and Fe burn in sulphur vapor. The sulphides are analo- 
gous to the oxides, exhibit similar reactions, and in the main 
possess a similar composition, as may be seen from the fol- 
lowing formulas: 
H20, Water. 
KOH, Potassium hydrate. 
BaO, Barium oxide. 
CO2, Carbon dioxide. 
C03K2, Potassium carbonate. 
H2S, Hydrogen sulphide. 
KSH, Potassium sulphydrate. 
BaS, Barium sulphide. 
CS2, Carbon disulphide. 
CS3K2, Potassium sulpho-carbonate. 
COMPOUNDS OF SULPHUR WITH HYDROGEN. 
H2S = 33.98. Density = 16.99. 
X. HYDROGEN SULPHIDE. 
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 in the reduction of alkaline sulphates by decom- 
posing carbon compounds. It may be formed directly from its 
constituents, although insmallquantity, if hydrogen gas be con- 
ducted 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 hy- 
drogen sulphide : 
Ag2S + H2 — 2Ag H2S. INORGANIC CHEMISTRY. 
For its production acids are allowed to act upon sulphides. 
Ordinarily iron sulphide and diluted sulphuric acid are em- 
ployed : the action occurs at ordinary temperatures: 
The operation is performed either in Kipp’s apparatus (p. 40) 
or in the one pictured in Fig. 30. 
Hydrogen sulphide thus obtained contains admixed hydro- 
gen, in consequence of metallic iron existing in the sulphide. 
The pure gas is obtained by heating antimony sulphide with 
hydrochloric acid : 
FeS + H.2S04 = FeS04 + H2S. 
Sb2S3 + 6HC1 = 2SbCl3 + 3H2S. 
Properties.—Hydrogen sulphide is a colorless gas, having 
an odor similar to that of rotten eggs ; inhaled in large quan- 
tities it has a stupefying effect, and is very poisonous. At 
medium temperatures it condenses under a pressure of 14 
atmospheres (under ordinary pressure at —740) to a colorless 
liquid of specific gravity 0.9, which at —85° C. solidifies to a 
white crystalline mass. Its density equals 16.99 (H = 1) or 
1.177 (air = 1). Water dissolves 3-4 times its volume of gas; 
the solution possesses all the properties of gaseous hydrogen 
sulphide and is therefore called hydrogen sulphide water. 
Ignited in the air the gas burns with a blue flame, water and 
sulphur dioxide resulting : 
h2s + 30 = h20 + so2. 
With insufficient air access, or when the flune is cooled by 
the introduction of a cold body, only hydrogen burns and 
sulphur separates out in a free condition. In aqueous solution 
hydrogen sulphide is decomposed by the oxygen of the air at 
ordinary temperatures, sulphur separating as a fine powder: 
H,S + O = H20 + S. 
For this reason hydrogen sulphide becomes turbid upon 
exposure to the air. 
The halogens behave like oxygen ; the hydrides of the halo- 
gens are formed with separation of sulphur: 
II2S + I, = 2HI + s. 
This reaction serves for the production of hydrogen iodide 
(p. 61). » 
As hydrogen sulphide has a great affinity for oxygen, it 
withdraws the latter from many of its compounds, and it 
therefore acts as a reducing agent (p. 82). Thus chromic, man- HYDROGEN SULPHIDE. 
107 
ganic and nitric acids are reduced to lower stages of oxidation. 
On pouring fuming nitric acid into a vessel filled with dry gas, 
the mixture will ignite with a slight explosion. 
Hydrogen sulphide possesses weak acid properties, reddens 
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 sulphides : 
Pb -f H2S = PbS + H2. 
With the oxides and hydroxides of the metals H2S yields 
sulphides and sulphydrates : 
KOH -J- H2S = KSH -f~ H20, 
Potassium hydrosulphide. 
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 insol- 
uble in water and dilute acids; therefore they are precipitated 
by H2S from solutions of metallic salts : 
Potassium sulphide. 
CuS04 + H2S = CuS + h.2so4. 
The precipitates thus obtained are variously colored (copper 
sulphide, black; cadmium sulphide, yellow; antimony sul- 
phide, orange), and answer for the characterization and recog- 
nition of the corresponding metals. Paper saturated with a 
lead solution is at once blackened by H2S, lead sulphide being 
formed—a delicate test for H2S. 
MOLECULAR FORMULA OF HYDROGEN SULPHIDE. ATOMIC 
WEIGHT OF SULPHUR. 
The analysis of hydrogen sulphide shows that it consists of one part hy- 
drogen and sixteen (more accurately 15.98) parts sulphur. If the mole- 
cular formula of hydrogen sulphide were HS, the atomic weight of sul- 
phur would be sixteen (compare p. 95). The great analogy of the sulphur 
compounds with those of oxygen (p. 105), permits us to accept formulas 
for the former similar to those of the latter. The molecular formula of 
hydrogen sulphide would, therefore, be H*S = 34, and the atomic weight 
of sulphur would equal 32. Hence the gas density of hydrogen sulphide 
must be 3g* == 17 (H = 1), or 1.177 (air 2=* 1); this is confirmed by direct 
experiment. Conversely it follows from the gas density that the molecular 
weight of hydrogen sulphide = 34. Since the analysis of 34 parts of hy- 
drogen sulphide shows the presence of two parts of hydrogen, the molecule 
of H,2S contains two hydrogen atoms. It then follows that the 32 parts of 
sulphur combined with the latter, correspond to one atom of sulphur, 108 
INORGANIC CHEMISTRY. 
because less than 32 parts of this element have never been found in the 
molecule of any compound in which sulphur occurs (p. 95). 
From the molecular formula H2S, we further conclude that the hy- 
drogen contained in one volume of hydrogen sulphide would occupy in a 
free condition, the same volume as the latter. 
nH2S 
contains 
nH, 
This conclusion is verified experimentally as follows : In a bent glass 
tube filled with mercury (Fig. 60), introduce dry hydrogen sulphide gas; 
then in the bent portion place a piece of tin, which is heated by a lamp. 
The sulphur of the II2S, com- 
bines with the metal to form 
scdid tin sulphide, while hy- 
drogen is set free: its volume 
is exactly equal to the volume 
of the employed hydrogen sul- 
phide. The quantity of sul- 
phur, 32 parts, in vapor form, 
at 1000° C., when the density 
is 32 (p. 104) combined with 
hydrogen (2 parts) would equal 
exactly half the volume of the hydrogen ; at 500° C., however, when the 
vapor density is three times as great, it will equal ]/& volume of the hydrogen. 
1 volume H2S, therefore, consists of 1 volume H and >4 volume sulphur 
vapor (at 500°), or as ordinarily expressed : 2 volumes H2S consists of 2 
volumes II and volume sulphur vapor. Written molecularly, we have: 
I vol. 
I vol. 
Fig. 6o. 
At 5oo°C. : S6 + 6H2 = 6H2S. 
i vol. 
6 vols. 
6 vols. 
At iooo0 C., however : S2 -(- 2ll2 = 2H2S. 
I vol. 
2 vols. 
2 vols. 
Just as hydrogen peroxide H202 is formed by (p. 98) the 
action of acids upon some peroxide, so may hydrogen per- 
sulphide be obtained from metallic persulphides. Calcium 
persulphide is most suitable, and when its aqueous solution is 
poured into dilute hydrochloric acid, 
2. HYDROGEN PERSULPHIDE. 
CaSs 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 supposed that the resulting hydrogen persulphide is con- 
stituted analogous to the peroxide and consists of hydrogen disulphide, 
containing an excess of dissolved sulphur. 
As the calcium persulphide used is a mixture of CaS2, CaSs, and 
H2S2 = H2S -f s. SULPHUR. 
CaS5, it is probable that the oily liquid is a mixture of H2S2, H2S3 and 
H2S5. We must at least conclude that H2S3 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. 
It is only the latter that meets with any practical appli- 
cation. 
Sulphur Dichloride—SC12—is produced when S2C12 is saturated with 
chlorine in the cold : 
The excess of chlorine is removed by conducting a stream of C02 
through it. 
It is a dark red-colored liquid, with a specific gravity of 1.62; boils at 
64° C., with partial decomposition into S2C12 and Cl2; the dissociation, 
commences at ordinary temperatures. 
Sulphur Tetrachloride—SC14—only exists at temperatures below o° C. 
It is formed by saturating SC12 with Cl at —30° C., and readily decomposes 
S2ci2 -)- Cl, — 2SCl2. 
Fig. 6i. 
into SC12 and Cl2; the dissociation commences at —20° C., and is complete 
at + 6°. It yields crystalline compounds with some chlorides, e.g., SnCl4, 
AsC13, SbCl3. 
The most stable of the sulphur chlorides is 
Sulphur Mono-chloride—S2C12—which is formed when 
chlorine is conducted over molten sulphur. (Fig. 61.) It distils INORGANIC CHEMISTRY. 
over and condenses in the receiver E; the product is redistilled, 
to obtain it pure. 
Sulphur mono chloride is a reddish-yellow liquid with a 
sharp odor, provoking tears, having a specific gravity of 1.68, 
and boiling at 1390 C. Its vapor density equals 67 (H = 1) 
corresponding to the molecular formula S2C12 — 134.7. It 
fumes strongly in the air, and is decomposed by water into 
sulphur dioxide, sulphur and hydrochloric acid : 
2S2ci2 + 2H20 = S02 + 4HC1 + 3S. 
Sulphur mono-chloride dissolves sulphur readily and serves in 
the vulcanization of caoutchouc. 
Bromine forms analogous compounds with S. S2Br2 is a red 
liquid, boiling at i90°-209° C. When gently heated, iodine 
unites with S to form S2I2. 
Se = 78.9. Se2 = 157.9 (at 1400° C.). 
3. SELENIUM. 
This element is not very abundant in nature, and is only 
found in small quantities, principally in certain iron pyrites (in 
Sweden and Bohemia). Upon roasting this ore of iron, for 
the preparation of sulphuric acid, selenium settles out in the 
chimney dust or in the deposit of the lead chambers (compare 
Sulphuric Acid), and was found there by Berzelius, in the year 
1817. 
Like sulphur, selenium forms different allotropic modifica- 
tions. 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. Selenium crystallizes from carbon 
disulphide in brownish-red crystals. The solution of potassium 
selenide is brown-red, and when it is exposed to the air, black 
leaf-like crystals of selenium separate. These are isomorphous 
with sulphur. Upon suddenly cooling fused selenium it soli- 
difies to an amorphous, glassy, black mass, which is soluble in 
carbon disulphide and has a specific gravity of 4.28. When 
selenium (amorphous) is heated to 970 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 metal- 
lic lustre, conducts electricity, and is insoluble in carbon 
disulphide. The crystalline, insoluble modification is obtained 
by slowly cooling the molten selenium. 
Selenium melts at 2170, and boils about 700°, passing into 
a dark yellow vapor. The vapor density diminishes regularly TELLURIUM. 
111 
with increasing temperature (similar to sulphur), and becomes 
constant at 1400° C. It then equals 79 ; the molecular weight 
is, therefore, 158, i.e. the molecule of selenium at 1400° C. 
consists of two atoms (2 X 78.9 = 157.9.) 
Selenium is a perfect analogue of sulphur. In the air it 
burns with a reddish-blue flame, forming Se02 and emits a 
peculiar odor resembling rotten horse-radish. It dissolves with 
a green color in concentrated sulphuric acid, and forms 
selenious acid. 
Hydrogen Selenide.—H2Se—produced like hydrogen sul- 
phide, is a colorless, disagreeably smelling gas with poisonous 
action. In the air the aqueous solution becomes turbid and 
free selenium separates. 
With chlorine selenium forms SeCl4, and Se2Cl4 perfectly analogous to 
the sulphur compounds. SeCl4 is a solid and sublimes without decompo- 
sition. 
4. TELLURIUM. 
Tellurium is of rare occurrence, either native or in combina- 
tion with metals. It is associated with gold and silver in 
sylvanite, and with silver and lead in altaite. It is found prin- 
cipally in Transylvania, Hungary, California, Virginia, Bolivia 
and Brazil. 
The tellurium precipitated by sulphurous acid from a solu- 
tion of tellurous acid (see this) is a black powder of specific 
gravity 5.928. 
The physical properties of tellurium indicate it to be 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. It fuses at 500° and vaporizes in a 
stream of hydrogen.. When heated in the air it burns, with a 
bluish-gray flame, to tellurium dioxide (Te02). 
The vapor density of tellurium at 1380° C., has been dis- 
covered to be about 126, corresponding to the molecular 
formula Te2. 
Hydrogen Telluride, H2Te, is a colorless, very poisonous 
gas, with disagreeable odor. Twochlorides—TeGl2and TeCl4 
—and two bromides—TeBr2 and TeBr4—have been formed. 
Te = 126.8 * 
* The atomic weight of tellurium, formerly taken as 128, has been deter- 
mined to be 126.8. The law of periodicity, to be presented later, seems 
to show with considerable certainty that the true weight will yet be found 
to be lower, some less than that of iodine—126.5. INORGANIC CHEMISTRY. 
SUMMARY OF THE ELEMENTS OF THE OXYGEN GROUP. 
The elements oxygen, sulphur, selenium and tellurium form 
a natural group of chemically similar bodies. The similarity 
of the last three is especially marked, 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: 
O. S. Se. Te. 
Atomic weights, 15.96 31.98 78.9 126. 
With the increase in the atomic weight*there occurs a simul- 
taneous condensation of substance, the volatility diminishes, 
while the specific gravity, and the points of fusion and boiling 
increase, as may be seen in the following table: 
Oxygen. 
Sulphur. 
Selenium. 
Tellurium. 
Specific gravity 
Melting point 
Boiling point 
Gas density 
15.96 
I.95-2.07 
hi. 50 
440° 
32 
4* 
w io 
2 8-1 
0 0 : 
00 
6.2 
500° 
White heat. 
126 
Oxygen is a difficultly coercible gas, while the others are 
solids at ordinary temperatures. We must, however, bear in 
mind that sulphur, selenium and tellurium in a free state are 
probably composed of larger complex atomic groups (see 
P- 104). ... . . . 
Further, with rising atomic weight the metalloidal passes 
into a more metallic character. Tellurium exhibits the phys- 
ical properties of a metal; even selenium possesses metallic 
properties in its crystalline modification. In chemical deport- 
ment, however, the metalloidal character shows scarcely any 
alteration. All four elements unite directly, at elevated tem- 
peratures, with two atoms of hydrogen, to form volatile gaseous 
compounds having an acid nature ; only the oxygen derivative 
—water—is liquid at ordinary temperatures and shows a neu- 
tral reaction. The hydrogen compounds are decomposed into 
their elements at a red heat. The affinity of hydrogen for 
oxygen is greatest; therefore, the aqueous solutions of H2S, 
H2Se and H2Te are decompos'ed by the air. OXYGEN GROUP. 
113 
Thermal Relations.—A measure for the chemical affinity of the ele- 
ments of the oxygen group is afforded, as in the case of the halogens (p. 
64), by their heat of formation. By the union of 2 grams of hydrogen 
with 15.96 grams of oxygen to steam of ioo° (H2,0—vapor), 57200 ca- 
lories are disengaged. Since in the condensation of steam to water of ioo° 
9635 additional (= 17.96 X 536.5) calories become free (latent heat of 
evaporation, p. 89I, and 17.96 calories by the cooling of water for every 
degree C., then, in the production of 1 molecular weight of water of o° 
from its elements, there are disengaged, all told, 68630 calories.* 
The heat disengagement is less in the formation of hydrogen'sulphide, 
while in case of hydrogen selenide heat is even absorbed, corresponding to 
the symbols: 
(H2,0 — vapor) = 57200 (H2,S) = 4500 (H2,Se) = — 5400. 
These numbers, as in the case of the halogen hydrides, explain the 
unequal power of combination of the elements of this group with hydrogen, 
and the stability of their compounds a.t a red heat accords with it. The 
fact already mentioned, that H2S, H2Se, and H2Te are decomposed by 
oxygen, with the formation of water and elimination of the elements, and 
the fact that the free elements (S, Se) do not act upon water, are explained 
by the proposition of the greatest development of heat. What is more 
striking is the behavior of the elements already mentioned, and their hy- 
drides, towards the metals, but this is fully demonstrated by their thermal 
relations; this will be more fully shown under the different groups of the 
elements. . We see, consequently, that here, as with the halogens, the 
chemical affinity of the homologous elements for hydrogen diminishes 
successively with increasing atomic weight (decrease of negative charac- 
ter). It must, however, be borne in mind, that the heat modulus repre- 
sents no direct measure of the chemical affinity. If, for example, water 
be produced corresponding to the molecular equation 2H2 + o2=2H2o, 
the hydrogen and oxygen molecules must first be broken up into individual 
atoms, to do which a definite quantity of heat is necessary (heat of decom- 
position). The directly observed heat disengagement of 57200 calories 
only indicates that the affinity of hydrogen to oxygen (2H,0) is greater 
than the affinity of the elementary molecules; 
Or, conversely, the content of energy of the hydrogen molecules is less 
than that of the hydrogen-oxygen molecules. 
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 heat of formation of 2H to S certainly is 
more than 4500 calories. The decomposition of molecules, accompanied 
by the absorption of heat, shows that definite thermal conditions are neces- 
sary to induce and carry through every reaction ; that, for example, sulphur 
and hydrogen only unite when raised to a high heat; their union, even 
then, is only partial. 
That heat is absorbed in the breaking of the molecules, follows, among 
others, from the fact that, in the process of combustion in nitric oxide 
(NO) more heat will be set free than in oxygen (02). The energy- 
content of NO is greater than that of 02. All the affinity masses derived 
2(H2,0) = 2(H,H) + (0,0) + 2. 57200. 
* 68360 calories, according to Thomsen. INORGANIC CHEMISTRY. 
114 
from thermal data are, therefore, only relative; it is only recently that it 
has become possible to determine the heat of dissociation of the hydrogen 
molecules (H,H) = 128000, and that of carbon (see carbon dioxide). 
NITROGEN GROUP. 
Here belong nitrogen, phosphorus, arsenic, antimony, and 
bismuth. The latter possesses a decidedly metallic character. 
These elements, bismuth excepted, form gaseous derivatives 
with three atoms of hydrogen. 
N — 14.01. N2 = 28.02. 
1. NITROGEN. 
Nitrogen exists free in the air, of which it constitutes f and 
oxygen the remaining In combination, it is chiefly found 
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 absorb- 
ing oxygen without 
acting upon the nitro- 
gen. This is most 
readily brought about 
by the combustion of 
phosphorus. Several 
pieces of the latter are 
placed in a dish swim- 
ming 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 phos- 
phorus pentoxide 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 red-hot tube filled with copper turnings; the 
copper unites with the oxygen and pure nitrogen escapes. At 
ordinary temperatures the removal of O from the air may be 
accomplished by the action of phosphorus, a solution of pyro- 
gallic acid, and other substances. 
A very convenient course for the direct preparation of 
Fig. 62. NITROGEN. 
115 
nitrogen is the following: Heat ammonium nitrite in a small 
glass retort; this decomposes the salt directly into water and 
nitrogen: 
NH4N02 = N2 + 2H20: 
In place of ammonium nitrite a mixture of potassium nitrite (KN02) 
and ammonium chloride (NH4CI) may be used; upon warming, these salts 
yield, by double decomposition, potassium chloride and ammonium nitrite 
(KN02-f-NH4C1 = NH4N02 -f- KC1), which latter decomposes further. 
As potassium nitrite usually contains free alkali, some potassium bichromate 
is added to combine the same. Practically, the solution consists of 1 part 
potassium nitrite, 1 part ammonium chloride, and 1 part potassium bi- 
chromate, in 3 parts water, and is then boiled ; to free the liberated nitrogen 
from every trace of oxygen the gas is conducted over ignited copper. 
The action of chlorine upon aqueous ammonia produces 
nitrogen. The chlorine combines with the Hof the ammonia, 
forming HC1, which unites with the excess of NH3 to produce 
ammonium chloride. The nitrogen that was in combination 
with the hydrogen is set free. The following equations express 
the reactions: 
Ammonium 
nitrite. 
and 
2NH3 + 3C12 = N2 + 6HC1, 
6HC1 -f 6NH3 = 6NH4C1. 
Ammonium 
chloride. 
The apparatus pictured in Fig. 36, page 47, will serve to 
carry out the experiment. The disengaged chlorine is con- 
ducted 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 formation 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 = 14,01 (H = r) or 0.9701 (air= 1). Water dis- 
solves about 2 per cent, by volume. In its chemical deport- 
ment 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 ex- 
tinguished, and animals are suffocated by 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 respiration. 
The presence of N in the air moderates the strong oxidizing 
property of the pure oxygen. INORGANIC CHEMISTRY. 
The air, or the envelope encircling the earth, consists prin- 
cipally of a mixture of nitrogen and oxygen ; it always con- 
tains, in addition, slight and variable quantities of 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 o° C. equals, upon an average, 760 millimeters. As 
1 c.c. of mercury weighs 13.596 grams, 76 c.c. will equal 
1 °33- 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 o° C. and 760 mm. pressure) 0.0012926 
grams; 1000 c.c., therefore, or one litre, would weigh 1.2926 
grams. As one litre of H..O weighs 1000 grams, air is conse- 
quently 773 times lighter than it. Air is 14.43 t'mes heavier 
than hydrogen. The specific gravities of the gases and vapors 
were formerly referred to air (= 1) ; compared with H = 1, 
they are, therefore, 14.43 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 solu- 
tion in practice. One litre of air weighs 1.293 grains, one litre of hydrogen 
0.08958 grams. To ascertain the weight of a litre of. any other gas or 
vapor, its specific gravity referred to air=r 1 must be multiplied by 1.2926, 
or if compared with H = 1 by the factor 0.08958. 
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 portion of the air 
disappeared, and that the part remaining was no longer suitable for the sup- 
port of the above processes; hence this was called destroyed air ; and the first, 
fire air. In the second half of the eighteenth century Scheele, in Sweden, 
and Priestly, in England, found that when a certain amount of gas, set free 
by heating mercuric oxide (oxygen), was added to the so-called destroyed 
air (nitrogen) a mixture resulted possessing all the properties of atmos- 
pheric air. Although both constituents of air were thus separately ob- 
tained and air regenerated by their mixture, yet at that time views regard- 
ing the nature of both ingredients and the nature of combustion and 
oxidation processes prevailed which were perfectly false in every respect. 
It was believed that combustion and oxidation were destructive processes; 
that the combustible and oxidizable bodies enclosed within themselves a 
peculiar substance called phlogiston. The latter was said to escape as fire 
and heat (phlogiston theory of Stahl, 1723), in the processes of combustion 
and reduction. These erroneous opinions were explained and corrected 
by Lavoisier (in 1774) by the following celebrated 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 ATMOSPHERE. THE ATMOSPHERE. 
117 
The open end of the neck dipped into a mercury trough, ft S, and was 
closed completely 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 ft, forming mercuric oxide. In the course of several days, during 
which no additional decrease in the volume of air was observable on 
the application of heat, the experiment was interrupted, and the volume 
of residual gas in A and P measured. Upon comparing this with the 
volume before the ex- 
periment, it was dis- 
covered that £ volume 
of the air had disap- 
peared and combined 
with the mercury to 
red mercuric oxide. 
Lavoisier now strongly 
ignited the resulting 
mercuric oxide, and 
obtained a volume of 
oxygen equal to that 
withdrawn from the'air 
during the experiment. 
By mixing this with 
the residual volume of 
N the original volume 
of air was again recov- 
ered. Thus it was de- 
monstrated that air consists of | volumes N and volume O gas. The 
elementary character of nitrogen was first established by Lavoisier in 1787. 
It was called azote (from life and a privative!, by him. The symbol 
Az, derived from azote, is used in France and England for nitrogen. The 
name nitrogenium (from which the symbol N) was given to nitrogen be- 
cause it was a constituent of saltpetre (nitrum). 
Lavoisier made use of the above experiment for another important de- 
duction. As he determined the weight, both of the employed mercury 
and the resulting mercuric oxide, he discovered that the increase in weight 
was exactly equal to that of the oxygen withdrawn from the air, and by 
heating 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)rand that the weight of a compound body 
equals the sum of the weights of its constituents; the principle of the inde- 
structibility of matter. ' • 
Fig. 63. 
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, V, with a capacity of about 20 litres (Fig. 
64), is connected with a porcelain tube, a b, filled with me- 
tallic copper. Balloon and tubes, dlosed by stop-cocks, are 
previously emptied and weighed apart. The bent tubes, A, B, 118 
INORGANIC CHEMISTRY. 
and C, contain KOH and sulphuric acid, and serve to free 
the air undergoing analysis from aqueous vapor, carbon di- 
oxide, and other impurities. The porcelain tube, filled with 
copper is heated to a red heat, and by carefully opening the 
stop cocks u, r, and / 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 
Fig. -64. 
the cocks and weigh the balloon and porcelain tube. The 
increase in weight of the latter represents the quantity 
of oxygen in the air; the increase in V the quantity of 
nitrogen. In this manner Dumas and Boussingault found that 
in 100 parts by weight of air there are contained : 
Nitrogen 76.99 parts by weight. 
Oxygen 23.01 “ “ “ 
100.00 “ “ “ 
As we know the specific gravity of nitrogen (14.01) and of 
oxygen (15.96), we can readily calculate the volume composi- 
tion of air from that in parts by weight. We thus discover : 
Oxygen : 20.78 parts by volume. 
Nitrogen * 7922 “ “ “ 
Air 100.00 “ “ “ 
Calculating upon these numbers, we obtain 14 415 (H — 1,0 — 15.96) 
as the specific gravity of air. THE ATMOSPHERE. 
119 
The volume composition of air may be directly found by 
means of the absorptiometer. The latter is a tube carefully 
graduated, and sealed at one end. This is filled with mer- 
cury, and air allowed to enter ; the volume of the latter is 
determined by reading off the 
divisions on the tube. Now in- 
troduce into the tube, through 
the mercury, a platinum wire 
having a ball of phosphorus at- 
tached to the end (Fig. 65), (or 
a ball of coke saturated with an 
alkaline solution of pyrogallic 
acid). The phosphorus absorbs 
the oxygen of the air, and only 
nitrogen remains, the volume of 
which is read off by the gradua- 
tion. 
The eudiometric method affords 
greater accuracy. It is depend- 
ent 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). Air and hydrogen are introduced into the 
eudiometer, and the electric spark then passed through the 
wires (Fig. 67). All the oxygen in the air combines with a 
portion of the hydrogen to form water. On cooling, the 
aqueous vapor condenses and a contraction in volume occurs. 
Assuming that we had taken 100 volumes of air and 50 vol- 
umes of hydrogen, and that the residual volume of gas, after 
allowing for all corrections (p. 120), equalled 87.15 ; then of the 
original 150 volumes of mixed gas, 62.85 volumes disappeared 
in the formation of water. As the latter results from the union 
of 1 volume of oxygen and 2 volumes of hydrogen, the 100 
volumes of air employed in the analysis therefore contained 
_6_2 j_5 — 20.95 volumes of oxygen. From this air con- 
sists of 
Fig. 65. 
79.05 volumes nitrogen. 
20.95 “ oxygen. 
ioo.oo “ air. 
As the above numbers have been obtained in numerous 
analyses, and as they clearly approximate those derived from 
the weight analysis of air, it was thought that the latter at all 120 
INORGANIC CHEMISTRY. 
times and in all places—on the earth’s surface and in the highest 
regions—contained like relative amounts of oxygen and nitro- 
gen. Recent researches show, however, variations in 
volume, ranging from 20.47 to 21.01 per cent. (v. 
Jolly), and it seems that in the higher regions there 
is less oxygen than near the surface. 
Measuring Gases.—The volume of gases is influenced by 
pressure, temperature, and the moisture contained in them. 
The volume of dry gases, at 760 mm. barometric pressure and 
o° C., is accepted as the normal volume. If a gas has been 
measured under any other conditions, it must be reduced to the 
Fig. 66. 
normal volume. According to the law of Boyle and Mariotte, 
the volumes of the gases are inversely proportional to the press- 
ure; therefore, if the volume of the gas at pressure h, has been 
Vh 
found equal to V, its volume at 760 mm. equals ?- — 
According to Gay-Lussac’s law, all gases expand in propor- 
tion to the temperature. Their coefficient of expansion is 
= 0.003665 ; i. e., one volume of gas at o° occupies at i° 
the volume 1.003665. If Vt represents the observed gas volume 
at $°, Vo, however, its volume at o°, then 
Vt 
Vo = *, 
I -l~ 0.003665.$ 
and, considering the pressure, 
' Vth 
Vu = • . 
760(1 -{- 0.003665.$) 
* Vo=V— T70.o.oo366.<, consequently Vo -j- V0. 0.00366.< = Vt, and 
Vo (1 + 0.00366./) = vr THE ATMOSPHERE. 
121 
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 sul- 
phuric acid, which dries it. It is more convenient, however, 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 ; the gas is then measured and its normal volume cal- 
culated by 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. 89). 
From the great constancy of its composition air was sup- 
posed to be a chemical compound, consisting of nitrogen and 
oxygen. This supposition is, however, opposed by the follow- 
ing circumstances. All chemical compounds contain their 
constituents in atomic quantities, which is not the case with 
air. In the mixing of nitrogen and oxygen to form air there 
is neither disengagement or absorption of heat, which is always 
observed in chemical compounds. Further, the air absorbed 
by water or other solvents possesses a composition different 
from the atmospheric ; this is due to the unequal solubilities of 
nitrogen and oxygen in water. The air expelled from water 
upon application of heat consists of 34.9 volumes of oxygen 
and 65.1 volumes nitrogen. (Bunsen). These facts indicate 
that air is not a chemical compound, but a mechanical mix- 
ture 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 progressive 
movement, they distribute themselves, without limitation, into space, and 
intermingle regularly with each other. The velocity of 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 = 
i; the density of oxygen — 16; therefore, hydrogen diffuses 4 times more 
rapidly than oxygen. .The unequal diffusion of gases may be perceived if 
they are allowed to pass through very narrow 'apertures, or through porous 
partitions. The following experiment very clearly illustrates 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 cylinder is displaced and rises in 
the water in bubbles: When the separation of gas ceases tube and 
cylinder are almost fdled with pure hydrogen. On removing the larger 
hydrogen vessel the gas will escape much more rapidly into the external 
air than the latter can enter the cylinder; the internal pressure will there- 
fore be less than the external, and water ascends in the glass tube. 
In addition to N and O, air constantly contains aqueous 
vapor and carbon dioxide in very small quantities. The pres- 122 
INORGANIC CHEMISTRY. 
ence of the former can readily be recognized by the fact that 
cold bodies are covered with dew in moist air. Its quantity' 
depends on the temperature 
and corresponds to the vapor 
tension of water (see p. 89). 
1 cubic meter of air perfectly 
saturated with aqueous vapor 
contains 22.5 grams water at 
250 C. ; on cooling to o° 
1 7.1 grams of these separate 
as rain. Generally the- air 
contains only 50-70 per 
cent, of the quantity of 
vapor necessary for complete 
saturation. The amount of 
moisture in it is either de- 
termined according to phys- 
ical methods (hygrometer), 
or directly by weighing. To 
this end a definite quantity 
of air is conducted through 
a tube filled with calcium 
chloride or sulphuric acid, 
and its increase in weight 
determined. 
To detect the carbon di- 
oxide in the air, conduct a 
portion of the latter through 
solutions of barium or cal- 
cium hydrates, and a tur- 
bidity wfll ensue. To de- 
termine its quantity, pass a 
definite and previously dried amount of air through a weighed 
potassium hydrate tube, and ascertain the increase in weight of 
the latter. 10,000 parts of atmospheric air contain, ordina- 
rily, from 2.9-3.0 parts carbon dioxide. 
Besides the four ingredients just mentioned, air usually 
contains small quantities of ozone, hydrogen peroxide, and 
ammonium salts (ammonium nitrite). Finally, air contains 
microscopic germs of lower organisms; they are generally 
found in the lower air strata, and their presence influences the 
processes of the decay and fermentation of organic substances. 
Fig. 68. AMMONIA. 
123 
COMPOUNDS OF NITROGEN WITH HYDROGEN. 
NH3 = 17.01. Density = 8.5. 
AMMONIA. 
Ammonia occurs in the air in combination with some acids, 
in natural waters and in the earth, but always in small quanti- 
ties. The formation of ammonia by the direct union of nitro- 
gen and hydrogen occurs under the influence of the silent 
electric discharge. Its compounds are frequently produced 
under the most varying conditions. Thus ammonium nitrate 
is formed by the action of the electric spark upon moist air: 
N, f O + 2H20 == NH4N03 • 
Ammonium 
nitrate. 
Small quantities of ammonium nitrite result by the evapora- 
tion of water in the air: 
n2 + 211,0 = nh4no, • 
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 
Fig. 69. 
vapors which moist phosphorus forms in the air, consist of am- 
monium nitrite. Further, ammonium salts are produced in the 124 
INORGANIC CHEMISTRY. 
solution of many metals in nitric acid, in consequence of a 
reduction of the acid by the liberated hydrogen : 
HN03 + 4H2 = 3H20 + NH3. 
Ammonia is produced in large quantities in the decomposi- 
tion and dry distillation of nitrogenous organic substances. 
Even as late as the last century the bulk of the ammonium 
chloride (the most important salt technically), was obtained 
by the distillation of camel’s dung (in Egypt in the.oasis of 
Jupiter Ammon—hence the name Sal ammoniacumJ. In the 
preparation of illuminating gas by the distillation of coal, am- 
monia appears as a by-product and may be obtained by 
combining it with sulphuric or hydrochloric acid. This method 
is used almost exclusively*at present for its production. 
To prepare ammonia heat a mixture of ammonium chloride 
and slaked lime in a glass or iron flask: 
2NH4C1 + Ca(OH)2 = CaCl2 -f 2H20 + 2NII3. 
Ammonium 
chloride,- 
Calcium 
hydrate. 
The disengaged ammonia gas is collected over mercury, as it 
is readily soluble in water (Fig. 69). For perfect drying con- 
duct it through a vessel filled with burnt lime (CaO). Calcium 
chloride is not applicable for 
this purpose, as it combines 
with the gas. Inconsequence 
of its levity, ammonia, like 
hydrogen, niay be collected 
by displacing the air in in- 
verted vessels. 
Ph) sical Properties. —Am- 
monia is a colorless gas with 
a suffocating characteristic 
odor. Its density is 8.5 
(H = 1), or 0.589 (air = 1). 
Under a pressure of 6.5 at- 
mospheres (at io° C.), or by 
cooling to — 40° C., it condenses to a colorless, mobile liquid 
with a specific gravity of 0.613 at °°> and solidifies at — 8o°. 
Fig. 70. 
Ammonia gas may be condensed, just like chlorine. Take ammonium 
silver chloride (AgC1.2NHs), obtained by conducting ammonia over silver 
chloride,and enclose 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 compound is decomposed into silver chloride 
and ammonia, which condenses in the cooled limb. AMMONIA. 
125 
Ammonia gas dissolves very readily in water, with the libera- 
tion of heat. One part of water at o° and 760 mm. pressure 
absorbs 1050 volumes (= 0..877 parts by weight) ; at 150, 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 it would into a vacuum; 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 greater the ammonia content the less will 
the specific gravity of the solution be. The solution saturated 
at 140 contains about 30 per cent. NH;}, and has a specific 
gravity of 0.897. All the gas escapes on the application of 
heat. 
When the condensed liquid ammonia evaporates it absorbs a great 
amount of heat and answers, therefore, for the production artificially of cold 
and ice in Carre’s apparatus. The simplest form of the latter is repre- 
sented in Fig. 71. The iron cyl- 
is filled about half with a 
concentrated aqueous ammonia 
solution, and is connected, by 
means of the tubes from b, 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, indicates 130° C., while 
F is cooled with water. In this 
way the gaseous ammonia is ex- 
pelled from the aqueous solution 
in A, passes through b, in which 
most of the water runs back, and 
condenses to a liquid in B, of the 
receiver F. The cylinder A is 
removed from the fire, cooled with water and the vessel D constructed of 
thin sheet-metal and filled with water, placed in the cavity E, which is 
surrounded with a poor conductor, e. g., felt. The ammonia condensed 
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 Carr6 for the artificial production of ice has acquired 
great application in the arts; recently, however, it has been more and 
more replaced by the method of Windhausen. The latter depends upon 
the expansion of compressed air. 
Fig. 71. 
Chemical Properties.—A red heat and continued action of 
the electric spark decompose ammonia into nitrogen and hy- 
drogen. On conducting ammonia gas over heated sodium or 126 
INORGANIC CHEMISTRY. 
potassium, the nitrogen combines with these metals and hydro- 
gen escapes: 
Ammonia will not burn in the air; in oxygen, however, it 
burns with a yellow flame: 
NHB + 3K = NK3 + 3H. 
2NH3 + 30 = N2 4- 3H20; 
ammonium nitrite and nitrogen dioxide are formed simulta- 
neously. When a mixture of ammonia and oxygen is ignited 
it burns 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 near a flame at the 
moment of the introduction 
of the glass tube. In- con- 
tact with oxygen, the am- 
monia gas ignites and con- 
tinues to burn in it. 
The following experi- 
ment (of Kraut) shows the 
combustion of ammonia 
very conveniently. Place a 
somewhat concentrated am- 
monia solution in a beaker 
glass; heat over a lamp, 
until there is an abundant 
disengagement of gas, and 
then run in oxygen gas, by 
means of a tube dipping 
into the liquid. Upon ap- 
proaching the mixture with a flame, it ignites with a slight 
explosion. The ignition may be induced without a flame, by 
sinking a glowing platinum spiral into the mixture (Fig. 72) ; 
we then have a' number of slight explosions. The glass is filled 
at the same time with white vapors of ammonium nitrite 
(NH4N02); later, when oxygen predominates, red vapors of 
nitrogen dioxide (N02) and nitrous acid appear. 
If chlorine gas be conducted into the vessel with ammonia, 
it immediately ignites and continues to burn in the latter, with 
the production of white fumes of ammonium chloride (NH4C1). 
The chlorine combines with the hydrogen of the ammonia, 
Fig. 72. AMMONIA. 
127 
with separation of nitrogen, and yields hydrochloric acid, which 
unites to form ammonium chloride with the excess of ammonia. 
NH3 + 3C1 = 3HCI + N, 
and 3NII3 + 3HCI = 3NH4C1. 
In gaseous form, as well as in solution, ammonia possesses strong 
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 fol- 
lowing illustrates the similarity: 
NH3 + HC1 = NH4C1 KC1 
Ammonium 
chloride. 
Potassium 
chloride. 
2NH3 + H2S04 = (NH4)2S04 K2S04 
Am. sulphate. 
Potassium 
sulphate. 
NHS + H,S == NH4SH KSH. 
Am. 
sulphydrate. 
Potassium 
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. The latter, 
when acted on by strong bases, yield ammonium gas : 
2NH1C1 + CaO = 2NH3 + CaCl2 + H20. 
The metallic character of the ammonium group is confirmed 
by the existence of the ammonium amalgam and likewise by its 
entire deportment in compounds. Therefore, the ammonium 
derivatives will be considered with the metals. 
QUANTITATIVE COMPOSITION OF AMMONIA. ATOMIC WEIGHT 
OF NITROGEN. 
The quantitative analysis of ammonia shows that it consists 
of 1 part hydrogen and 4.67 parts nitrogen; hence we con- 
clude that the atomic weight of N is a multiple of the last 
number. (See p. 69.) 
H = 1 
N = 4.67 
NH = 5.67 
2II = 2 
N = 9-34 
NH2 = 11.34 
3H = 3 
N == 14.01 
NH, = 17.01 
As the density of ammonia equals 8.5 (H = 1) its molecular 
weight would almost = 17. In 17.01 parts of ammonia there 
are 3 parts, and, therefore, 3 atoms of hydrogen. That the 
# 128 
INORGANIC CHEMISTRY. 
14.01 parts nitrogen united with them correspond to one atom 
of N is a consequence, as never less than 14.01 parts of N are 
present in the molecular weight of any nitrogen derivative. 
The density of nitrogen equals 14.01, and its molecular weight 
28.02; therefore, the molecule of N consists of two atoms 
(N2). This is also concluded from the volume ratios occurring 
in the formation of ammonia. (See below.) 
From the molecular formulas NH3 and N2 follows, further, 
that 1 volume N and 3 volumes H form 2 volumes ammonia|gas, 
or that 2 volumes NH3 decompose into 3 volumes H2 and 1 
volume N2, corresponding to the molecular equation: 
N2 + 3H2 = 2NH3. 
1 vol. 3 vols. 2 vols. 
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. 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 every vol. N is 
easily shown by the volumetric method, by burning the H with 
oxygen (p. 119). 
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, in atoms of N. The nitrogen contained in 
these 2 volumes of NH3 occupies 1 volume in a free condition, and this 
contains n molecules and therefore in atoms of N. 
HYDROXYL AMINE. 
This compound, very analogous to ammonia, was discovered 
(by Lossen) in the reduction of ethyl nitrate by zinc and hy- 
drochloric 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 : 
NHsO = NHaOH. 
HN03 + 3H2 = H3NO + H20. NITROGEN. 
129 
To prepare hydroxylamine treat ethyl nitrate (120 gr.) with 
granulated tin (400 gr.) and hydrochloric acid (800-1000 c.c. 
of specific gravity 1.19, mixed with three times its volume of 
water) until solution is obtained. The strongly concentrated 
liquid is cooled and supersaturated with soda, the 
acidulated with hydrochloric acid aijd then evaporated to dry- 
ness. Hot alcohol will extract hydroxylamine hydrochloride, 
NH3O.HCI, from the residue. 
Hydroxylamine is very similar to ammonia, and like it unites 
directly with acids to form salts : 
On adding to the aqueous solution of the sulphate of hydroxyl- 
amine sufficient barium hydrate to remove all the sulphuric 
acid, an aqueous solution of the base is obtained, which, like 
the ammonia solution, possesses strong basic properties, and 
blues red litmus paper. The solution is, however, very un- 
stable, and readily decomposes into water, ammonia, and 
nitrogen : 
H3NO 4- HCl = HjNO.HCl. 
3NH3O = NH3 + 3H20 + N2. 
Upon the application of heat a portion of the hydroxylamine 
will be carried over undecomposed along with the steam, but 
most of it is broken up. The hydroxylamine solution mani- 
fests a reducing action ; it precipitates metallic silver from 
silver nitrate, white mercurous chloride, HgCl, from mercuric 
chloride, HgCl2, and cuprous oxide from cupric salts. 
Owing to its great similarity to ammonia and its various 
reactions, it is supposed that hydroxylamine represents ammonia 
in which 1 H is replaced by the hydroxyl group OH ; there- 
fore the name hydroxylamine: 
nh3o = nh2oh. 
COMPOUNDS OF NITROGEN WITH THE 
HALOGENS. 
NITROGEN CHLORIDE. 
As we have seen, nitrogen is liberated when chlorine acts 
upon an excess of ammonia (p. 126); when, however, the 
chlorine is in excess, it acts upon the previously formed ammo- 
nium chloride, to produce nitrogen chloride: 
NCI,. 
NII4C1 -f 3C12 = NC13 + 4IICI. 
For the preparation of a small quantity of nitrogen chlo- 
ride, dip a flask filled with chlorine, open end down, into an 130 
INORGANIC CHEMISTRY. 
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. 
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 decomposes by the 
slightest contact with many substances, and frequently too 
without any perceptible external cause. Its decomposition is 
accompanied by an extremely violent report. 
The formation and explosibility of nitrogen chloride may be illustrated 
in a harmless way as follows : Decompose a saturated ammonium chlo- 
ride solution with the electric current. Nitrogen chloride rising in small 
drops from the liquid will separate at the positive pole. Upon covering 
the surface of the solution with a thin layer of turpentine oil, each drop 
will explode as it comes in contact with the latter. 
Nitrogen Iodide. 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, 
which are then allowed to dry; upon the slightest disturbance 
these pieces explode with a violent noise. The iodide has the 
composition NHI2 or NIS according to the method by which it 
was prepared; it is regarded as ammonia in which the hydro- 
gen is partly or entirely replaced by iodine. 
2. PHOSPHORUS. 
This element does not occur free in nature, because of its 
very great affinity for oxygen. The phosphates, especially 
calcium phosphate, are widely distributed. By the disintegra- 
tion 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 Kunkel, in Hamburg (1669), first obtained 
phosphorus by the ignition of evaporated urine. In 1769, 
Scheele, in Sweden, showed that it could be obtained from 
bones. Its name is derived from its power of giving light in 
the dark—<pw<T<popo<;, i. e., light-bearer. 
P = 30.96. P4 == 123.8. Density — 61.9. 
To obtain phosphorus from bones the latter are burned, thereby destroy- 
ing all organic admixtures and leaving bone ashes, which consist princi- PHOSPHORUS. 
I3I 
pally of tertiary calcium phosphate (P04)2Ca3 (see Phosphoric acid). The 
ashes are digested with % of their weight of sulphuric acid, when the 
tri-phosphate becomes primary calcium phosphate, and gypsum (cal. sul- 
phate) is produced : 
Ca3(P04)2 •+ 2H2S04 = CaH4(P04)2 + 2CaS04. 
Tertiary 
calcium phosphate. 
Primary 
calc, phosphate. 
Calcium 
sulphate. 
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 expels water from the primary phosphate and the latter changes 
to calcium metaphosphate : 
CaH4 (P04)2 = Ca (P03)2 + 2H20. 
Calcium 
metaphosphate. 
The ignited residue is then raised to a white heat, in retorts of infusible 
clay. The carbon partly reduces the metaphosphate to phosphorus, by 
forming carbon monoxide with oxygen, and half of the phosphorus con- 
tained in the metaphosphate remains as calcium pyrophosphate: 
2Ca (POj)j T 5C — 2P -p 0 *T Ca2P207. 
Calcium 
pyrophosphate. 
The liberated phosphorus escapes in vapor, and is collected and con- 
densed under water in receivers of peculiar construction. To remove 
mechanically admixed impurities the phosphorus is again distilled from 
retorts and fused under water; it is then moulded into sticks. 
The crystalline ox yellow phosphorus obtained by distillation 
is a waxy, transparent, slightly yellow-colored substance, with 
specific gravity of 1.83 at io° C. At ordinary temperatures 
it is soft and tough; at o° it becomes brittle. It fuses under 
water at 44.4° and boils at 290° C. (278.3°). By the action 
of sunlight it becomes yellow, and is coated with a non-trans- 
parent, reddish-white layer. Phosphorus is insoluble in water, 
slightly soluble in alcohol and ether, and very readily soluble 
in carbon disulphide. It crystallizes from the latter solution 
in forms of the isometric (rhombic dodecahedra) system. When 
exposed to moist air, it oxidizes to phosphorous acid (H3P03); 
the white vapors which arise contain ammonium nitrite (NH4 
N02), ozone and hydrogen peroxide. Its' odor resembles that 
of ozone. In the air it phosphoresces at night. It does this 
also in other gases, but only in such as contain1 oxygen. It 
appears the phosphorescence is influenced by the formation 
and combustion of the self-inflammable phosphine, as all sub- 
stances which destroy the latter, prevent and put an end to the 
former. It is noteworthy that in pure oxygen the oxidation 
of phosphorus begins at 27°. If the oxygen be diluted by re- 132 
INORGANIC CHEMISTRY. 
moval over an air pump or by the addition of neutral gases, so 
that its quantity is not more than 40 per cent., the absorption 
will be very energetic at 20°, but cease entirely at 70. 
Another modification—the red or amorphous phosphorus— 
possesses properties entirely different from the ordinary variety. 
It is a reddish-brown amorphous powder, of specific gravity 
2.14; insoluble in carbon disulphide, non-phosphorescent, does 
not alter in the air, and is, indeed very stable. While ordinary 
phosphorus is very poisonous, this variety is perfectly harmless. 
It does not fuse at a red heat, even when subjected to strong 
pressure, and vaporizes very slowly (above 260°), although 
only partially, the vapors passing over into ordinary phos- 
phorus. 
To prepare the red variety, yellow phosphorus is heated for 
some minutes to 300°, in closed, air-tight iron vessels; there 
is a partial conversion at 250°. The resulting mass is then 
treated with carbon disulphide or sodium hydrate, to withdraw 
the 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 amorphous 
variety be heated in a glass tube, free of air, to 530°. Microscopic needles 
then sublime into the upper, less heated, portion of the tube. It is more 
easily obtained if phosphorus is heated with lead, in a closed tube to a red 
heat. The molten metal dissolves the phosphorus, and on cooling, the lat- 
ter separates in black, metallic, shining crystals. Metallic phosphorus pos- 
sesses 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 over a small piece of phosphorus, 
heated in a glass tube, the escaping gas will burn with a bright 
green flame. When .ordinary phosphorus is distilled with 
water, some passes over with the steam and, in the dark, phos- 
phoresces. This procedure serves for the detection of phos- 
phorus in poisoning by this substance. 
The density of P equals 61.9 (H — 1), or 4.29 (air = 1) ; 
the molecular weight Is, therefore, 123.8. As the atomic 
weight of P is 30.96, it follows that the molecule in the form 
of vapor consists of 4 atoms: P4 = 123.8 (30.96 X 4)- We 
saw that the sulphur molecule at 500° consists of 6 atoms (S6), 
and at 900° of 2 atoms (S2). Such a dissociation does not; 
however, occur with phosphorus; even at 1040° its vapor den- 
sity remains unaltered, although a partial dissociation does take 
place at a very intense heat. PHOSPHORUS. 
133 
When phosphorus is burned in oxygen or in air, it forms the 
pentoxide (P205). The ordinary variety inflames at 40°, and 
also by gentle friction ; the amorphous is not ignited below 260° 
The first will burn with a bright flame even under water. To 
this end heat some pieces of P in a flask w ith water, until they 
fuse, and conduct a current of oxygen through the water. 
Phosphorus combines very energetically with Cl, Br, and I 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. 
COMPOUNDS OF PHOSPHORUS WITH HYDROGEN. 
PH3, P2H4, P4H2. 
The compounds of phosphorus with hydrogen can be pre- 
pared by the action of nascent hydrogen upon phosphorus, as, 
for example, on gently heating dilute sulphuric acid with zinc 
and phosphorus (p. 140). The usual course is to heat yellow 
phosphorus with concentrated potassium or sodium hydrate, 
when spontaneously inflammable phosphine will escape and a 
salt of hypophosphorous acid enter solution. 
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 aque- 
ous 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 the delivery tube, the other 
end of which dips under warm water, to prevent any obstruction arising 
in it from phosphorus that may be carried over and solidify by cooling. 
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 (PH,) 
and hydrogen, with which is mixed a small quantity of a liquid 
substance (P2H4), whose presence imparts the spontaneous in- 
flammability to the gas. On conducting the latter through a 
cooled tube the P2H4 is condensed to a liquid, and the escaping 
gas no longer inflames spontaneously. The liquid compound 
may be isolated in a similar manner if the gas is conducted 
through alcohol or ether, which will absorb the compound 
P2H4. 134 
INORGANIC CHEMISTRY. 
Liquid Phosphine, P2H4, separated from the gas by cooling, is a 
colorless, strongly refracting liquid, insoluble in water, and boiling at 3o°. 
It 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- 
fig. 73. 
mability. In contact with some compounds, like carbon and sulphur, and 
by the action of sunlight, it decomposes into gaseous and solid phosphine: 
'Solid Phosphine, P4H2, (?) is a yellow powder, inflamed at i6o° or by a 
blow. ' It is produced in the decomposition of calcium phosphide by hy- 
drochloric acid. • • 
5P2H4 = 6PH3 + P4H2. 
Gaseous Phosphine, PH3, maybe formed, in addition to 
the manner previously described, by the action of water or 
hydrochloric acid upon calcium phosphide: 
Further, by the ignition of phosphorous and hypophosphor- 
ous acids: 
Ca3P2 + 6HC1 = 3CaCl2 + 2PH,. 
4H3P03 = PHS + 3H,P04. 
Phosphorous 
acid. 
Phosphoric 
acid. 
It is a colorless gas, with a disagreeable, garlic-like odor, and PHOSPHORUS. 
135 
is somewhat soluble in alcohol. Its density is 16.98 (H = 1), or 
1.176 (air =1). When pure—freed of P2H4—it ignites at ioo° 
C. Oxidizing agents convert it again into the spontaneously 
inflammable variety, owing to the production of P2H4. It is 
extremely poisonous. Heat and the electric current decom- 
pose PH3 into phosphorus and hydrogen. When mixed with 
chlorine it explodes violently', with production of phosphorus 
trichloride and hydrogen chlorides 
PH, + 3C12 = PCI, + 3IICI. 
Like ammonia, phosphine possesses faint alkaline properties, 
and combines with hydrogen iodide and bromide to yield 
compounds similar to ammonium chloride: 
PH3 + HI = PH4I. 
It combines with HC1 at —30° to —35°, or, at ordinary 
temperatures, under a pressure of 20 atmospheres. The group 
PH4, figuring in the role of a metal in these compounds, is 
analogous to ammonium (p. 127), and termed Phosphonium. 
Phosphonium iodide, PH4I. It is best prepared by the 
decomposition of phosphorus di-iodide (PI2), by a slight quan- 
tity of water, or by adding yellow phosphorus (1 o parts), and, after 
some hours, iodine (2 parts), to a saturated solution of hydriodic 
acid (22 parts). The liquid becomes a solid mass, consisting 
of phosphonium iodide and phosphorous acid- Phosphonium 
iodide sublimes in colorless, shining, cube-like rhombohedra; 
fumes in the air, and, with water, decomposes into PH3 and 
HI. When decomposed by potassium. hydrate it yields pure 
hydrogen phosphide, which is not spontaneously inflammable. 
PH4I -f KOH = KI + PH3 + H20. 
MOLECULAR FORMULA OF PHOSPHINE. ATOMIC WEIGHT OF 
PHOSPHORUS. 
The analysis of phosphine shows that it consists of 1 part hydrogen and 
10.32 parts phosphorus. Were its molecular formula PH, the atomic 
weight of P would be 10.32. The great analogy of phosphine to NH3, 
and that of all the P compounds to those of N, argues, however, 
for the formula PII3. The atomic weight of P, therefore, is 30.96 
(= 3 X 10.32), and the molecular weight of the phosphine is 33.96. 
H, = 3 • 
P = 30.96 
PH:J = 33.96 
This view is confirmed by the density which, according to 
its formula must be — 16.98. Direct experiment con- 
. 2 INORGANIC CHEMISTRY. 
firms this. Further, from the formula PH3 it follows that 3 
volumes of hydrogen are present in 2 volumes of the gas: 
2PH3 
2 vols. 
contain 
3H2, 
3 vols. 
or in i volume there are volumes of hydrogen. On de- 
composing the gas in an eudiometer by means of electric 
sparks, it will be found that the volume increases times; 
the gas consists, then, of pure hydrogen, while phosphorus 
separates in a solid condition. As the phosphorus molecule in 
the gaseous condition is composed of 4 atoms, the phosphorus 
(61.9 parts) separated from '2 volumes of PH3, will fill 
volume when in the form of vapor; hence, in 2 volumes of 
PH3 there are present 3 volumes of H and volume of phos- 
phorus vapor. 
Or, written molecularly 
P4 + 6H2 = 4PH3. 
1 vol. 
6 vols. 
4 vols. 
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. 
Fig. 74. 
Phosphorus Trichloride—Phosphorous Chloride—PC13. 
Conduct dry chlorine gas over phosphorus gently heated in PHOSPHORUS OXYCHLORIDE. 
137 
the retort D (Fig. 74). The phosphorus ignites in the stream 
of gas, and distils 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 740 C., 
and has a sharp, peculiar odor. Its specific grayity equals 
1.616 at o°. It fumes strongly in the air, and is decomposed 
by moisture into phosphorous and hydrochloric acids: 
PC13 + 3H20 = H3PO3 4 3HC1. 
The vapor density of the trichloride equals 68.5 (H = 1), 
corresponding to the molecular,formula PCl3 = i37.o. 
Phosphorus Pentachloride—Phosphoric Chloride— 
PC15.—This is produced by the action of an excess of chlorine 
upon the liquid trichloride. It is a solid, crystalline, yellowish- 
white compound. It fumes strongly in the air and sublimes 
without fusion when heated. It at the same time sustains a 
partial decomposition into trichloride and chlorine. 
At lower temperatures (in an atmosphere of chlorine) the vapor density 
of the pentachloride has been found to be 103.9, corresponding to the 
molecular formula PC15 — 103.9. At increased temperatures the 
vapor density steadily diminishes, anil a gradual decomposition occurs—■ 
dissociation (p. 93) of the molecules PCI5 into the molecules PC13 and Cl,2. 
The dissociation is complete at 236°, and then equals the vapor density 
51.9; i. e., the vapor then fills a volume twice as large as at a lower tem- 
perature. The breaking up of PC15 into PC13 and Cl2 explains this: 
I vol. PC15 = i vol. PC13 -f- 1 vol. CI2. 
That such a decomposition of the penta- into trichloride and chlorine 
does really occur, is proven among other things by the originally colorless 
vSpor gradually assuming the yellow color of chlorine as the temperature 
rises. The decomposition products—PC13 and Cl2—may be separated from 
each other by diffusion (p. 121). 
PC15 acts very energetically with water. With a little water it 
forms the oxychloride and hydrochloric acid : 
PC15 + H20 = PC130 + 2HCI. 
Phosphorus Oxychloride—POCl3, is a colorless liquid, 
fuming strongly in the air, with a specific gravity, at 120, of 
1.7. It boils without decomposition at 11 o°. Its vapor den- 
sity equals 76.5, corresponding to the molecular formula POCl3 
= 153.0. Water decomposes it into metaphosphoric and 
hydrochloric acids: 
POCl3 4- 2U.fi = HP03 4. .3HCI. 
Phosphorus oxychloride is most practically obtained by the 
distillation of PC16, with an excess of phosphorus pentoxide: 
3PC15 + P205 - 5POCI3. 138 
INORGANIC CHEMISTRY. 
Or chlorine gas is conducted into a mixture of 3PCI3 and 
P205: ' 
3pci3 + PA + 3C12 = 5POCI3. 
Its production on conducting ozonized air through phos- 
phorous chloride (Remsen) is quite interesting. PC13 unites 
with S in the same manner when heated to 130°. It forms 
phosphorus sulpho-chloride—PSC1,—an oily liquid, sinking in 
water. It is decomposed by the latter into SH2, 3CIH 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- 
getic, it is best to proceed as follows: Dissolve the phosphorus 
in carbon disulphide, gradually add the calculated amount of 
Br or I, and then distil off the volatile solvent. 
Phosphorus Tribromide, PBr3, is a colorless liquid, boiling at 1750, 
and having a specific gravity of 2.7. The pentabroiride, formed by the 
gradual addition of 2Br to PBr3 is a yellow, crystalline substance, which 
fuses when heated, and breaks up into PBr3 and Br2. Water decomposes 
both compounds, as it does the corresponding chlorides. Phosphorus 
oxy-bromide (POBr3), is a colorless crystalline mass, fusing at 450, and 
boiling at 195°. 
Phosphorus Chlor-bromide, PCI3Br2, is produced by the union of 
PC13 with Br2 in the cold. It is a yellowish-red mass, which decomposes at 
350 C., into PCI, and Br2. • 
Phosphorus Tri Iodide—PI3, forms red crystals, fusing at 550 and 
distils, with partial decomposition, at a higher temperature. The so-called 
phosphorus iodide, PI2, or P2I4 (corresponding to P2H4), crystallizes in 
beautiful orange-red needles or prisms, and fuses at 1 io°. A little water 
decomposes it into phosphorous acid, PHS and HI. The last two bodies 
then form phosphonium iodide, PH4I (p. 135). 
The recently discovered Phosphorus Pentafluoride—PF15—is inter- 
esting. It results upon heating PC13 or PC15 with arsenic trifluoride, AsF13. 
3PC15 + 5AsF13 = 3PF15 + 5AsC13. 
It is a colorless gas that fumes in moist air and is decomposed by water 
into phosphoric acid and hydrogen fluoride. Its density is 63, correspond- 
ing to the molecular formula PF16 = 125.9. 
It is rather remarkable that although phosphorus pentaiodide could not 
be obtained, the stability of the compounds, PBr5, PC15,PF15, gradually 
increases with the diminution of the atomic weight of the combined halo- 
gens.' PF15 can be gasified without decomposition. ARSENIC WITH HYDROGEN. 
139 
As = 74.9. As4 = 299.6. Vapor density, 149.8. 
3. ARSENIC. 
Arsenic is a perfect analogue of phosphorus, but possesses a 
somewhat metallic character. In its free state it is similar to 
metals. 
Arsenic is found free in nature, although it occurs more fre- 
quently in combination with sulphur (realgar, orpiment), with 
oxygen (arsenolite, As203), and with metals (mispickel, FeSAs, 
cobaltite, CoAsS). To prepare it, heat mispickel with some 
iron, and free arsenic will sublime. Or, in the customary 
way of isolating metals from their oxides, heat the trioxide 
(arsenolite) with charcoal: 
As203 -)- 3C = 2As -\- 3CO. 
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 2100— 
220°, forms steel gray hexagonal rhombohedra, with strong 
metallic lustre; its specific gravity equals 5.7. 
Arsenic volatilizes at 180° without previously fusing ; it will, 
however, fuse if heated under great pressure in a sealed tube 
(p. 88). Its vapor possesses a lemon-yellow color. The 
vapor density is 149.8 (H = 1), the molecular weight, there- 
fore, 299.6. As its atomic weight equals 74.9, it follows that 
the molecule in the state of gas, like that of phosphorus, con- 
sists of four atoms (As4 = 299.6 = 4 X 74-9)* 
Arsenic does not change in dry air. When 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 will inflame when projected into 
•chlorine gas. It yields arsenides with the metals. 
It is remarkable that arsenic, belonging to the nitrogen group and gen- 
erally forming compounds which in constitution are quite different from 
those of sulphur, should be analogous to the latter in its metallic combina- 
tions. 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 OF ARSENIC WITH HYDROGEN. 
Arsine, AsH3 = 77.9. Like nitrogen and phosphorus, 
arsenic furnishes a gaseous compound containing 3 atoms of 140 
INORGANIC CHEMISTRY. 
hydrogen. It is obtained pure by the action of dilute sul- 
phuric acid or hydrochloric acid upon an alloy of zinc and 
arsenic: 
As2Zn3 -|- 6HC1 = 3ZnCl2 2AsH3. 
It also results in the action of nascent H (zinc and sulphuric 
acid), upon many arsenic compounds, as, e.g., the tri-oxide: 
As203 + 6H2 = 2AsH3 -f- 3H20. 
Arsine is a colorless gas, of strong, garlicky odor, and ex- 
tremely poisonous action ; it may be condensed to a liquid at 
—40°. Its density equals 38.9 (H = 1), or 2.69 (air — 1). 
It burns with a bluish-white flame when ignited, and evolves 
white fumes of arsenic tri-oxide : 
It is decomposed at a dull glowing-heat or by the electric 
spark into arsenicand hydrogen. On conducting the gas through 
9. heated tube the arsenic deposits itself behind the heated part 
as a metallic coating (arsenic mirror). On holding 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 PH3; its basic properties are 
very slight, and it does not furnish any derivatives with the 
halogens. 
2AsH3 -)- 302 — As203 4-' 3H20. 
According to analysis, arsine consists of I part by weight of hydrogen 
and 24.97 parts arsenic. If, because of its analogy to PH:j, we ascribe to 
it the formula AsH3; then the atomic weight of arsenic would be 74.9 
(3 x 24.97) an(l the molecular weight of AsH3 would = 77.9. Hence 
the density must be — 38.9, which is confirmed by experiment. The 
formula, too, shows that 3 volumes of hydrogen are present in 2 volumes 
of As H3: 
2AsH3 contain 3H2. 
2 vols. 3 vols. 
We can satisfy ourselves of this by decomposing the gas by electricity in 
an eudiometer (see p. 136). 
Marsh's Method for the Detection of Arsenic.—The detection 
of arsenic is very important, because, of the poisonous nature 
of the element. The method of Marsh is based upon the for- 
mation and the characteristic properties of arsine. It is as 
follows: Hydrogen is generated in a flask a (Fig. 75), by the 
action of dilute sulphuric acid upon zinc, and a portion of the 
solution to be tested for arsenic is introduced through the fun- 
nel-tube. The liberated gas, a mixture of hydrogen and arsine, 
is dried in the calcium chloride tube c and escapes through the ARSENIC WITH HYDROGEN. 
141 
difficultly fusible glass tube d, which is contracted at several 
points. 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 with a bluish-white 
flame, if arsenic is present, and at the same time disseminate a 
white vapor. The dark arsenic spots are obtained by holding 
a cold porcelain dish in the flame. If the tube d be heated 
(as shown in Fig.’ 75), an arsenic mirror will be formed upon 
the adjacent contraction. The slightest traces of arsenic may 
be detected by this method. 
Fig. 75. 
Besides the ordinary arsine, AsH3, we might expect the existence of 
As2H4 and As4H2, corresponding to the liquid and solid phosphines (P2H4 
and P4H2). The first is not known;'its derivatives exist, and contain 
hydrocarbon groups instead of hydrogen. An example of this class is 
cacodyl, As2(CH3)4 = (CH3)2As-As(CH3’>2. Nitrogen affords similar com- 
pounds—(CH3)2N-NH2 and (CH3)NH-NH2, derived from diamine 
(N2H4 = H2N-NH2) which is not known in a free condition. 
The solid arsine, As4H2, is obtained by the action of nascent hydrogen 
upon arsenic compounds in the presence of nitric acid. It forms a reddish- 
brown powder, which decomposes when heated. 142 
INORGANIC CHEMISTRY. 
COMPOUNDS OF ARSENIC WITH THE HALOGENS. 
These are perfectly analogous to the corresponding phos- 
phorus compounds, and are the result of the direct union of 
their constituents. The iodide is the only known representa- 
tive of the compounds with the formula AsX5 (see p. 136). 
The metallic character of arsenic is shown by the fact that 
arsenic chloride, like other metallic chlorides, may be obtained 
by the action of hydrochloric acid upon the oxide; 
Arsenic Chloride is evolved when a. solution of As203 is 
boiled with concentrated hydrochloric acid. 
Arsenic Tri-chloride—AsC13. A colorless, oily liquid, 
fuming' in the air, and having a specific gravity of 2.2. It 
solidifies at —30° and boils at 1340. The vapor density equals 
90.50 (H = 1), corresponding to the molecular formula AsC13 
— 181.0. The chloride dissolves in a small quantity of water 
withput change, while much water converts it into oxide and 
hydrochloric acid: 
As203 + 6HC1 = 2AsC13 + 3H20. 
Arsenic Tribromide, AsBr3, is a white crystalline mass, 
fusing at 20°, and boiling at 220° C. The Tri-iodide, Asl3, 
forms red crystals; the Trifluoride, AsF13, is a liquid fuming 
strongly in the air. It results in the distillation of AsC13 or 
As203 with calcium fluoride ahd sulphuric acid. Arsenic 
pentaiodide, Asl5, melts at 70°, and is very soluble in water 
and alcohol. 
2AsC13 + 3H20 = As2©3 + 6HC1: 
4. ANTIMONY. 
Sb = 119.6. 
The metallic character exhibited by arsenic, becomes more 
distinct with antimony, which at the same time retains its 
complete analogy to the metalloidal elements, arsenic and 
phosphorus. Antimony is a perfect metal so far as its physical 
properties are concerned. 
It (Stibium) occurs in nature chiefly in union with sulphur, 
as stibnite, Sb2S3, and with sulphur and metals in many ores. 
It is almost always accompanied by arsenic. To prepare anti- 
mony, stibnite is roasted in a furnace, i.e , heated with air 
access, whereby the sulphur burns, and antimony trioxide 
remains: 
Sb2S, + 90 = Sb2Os + 3S02. ANTIMONY. 
The residual oxide is ignited together with carbon, which 
reduces it to metal (general procedure for the separation of 
metals). Antimony may also be obtained by heating its sul- 
phide with iron, which combines with the sulphur : 
Sb2S3 + 3Fe = 2Sb + 3FeS. 
The resulting commercial crude antimony is further purified 
in the laboratory by fusing it with nitre, whereby the admixed 
arsenic, sulphur, and lead are removed. Chemically pure anti- 
mony is obtained by reducing the pure oxide. 
It is a silver-white, and very brilliant 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. 
It fuses at 430°, and distils at a white heat. It is not altered 
in the air at ordinary temperatures; but when heated it burns 
with a blue flame, yielding white vapors of antimonic oxide, 
Sb2Os. Like phosphorus and arsenic it combines directly with 
the halogens; powdered antimony inflames in chlorine gas. It 
is insoluble in hydrochloric acid; nitric acid oxidizes it to 
antimonic oxide. 
Hydrogen Antimonide—Stibine—(SbH3) is produced 
like arsine, and is very similar to the latter. It is always 
obtained mixed with hydrogen. It is a colorless gas of pecu- 
liar odor, and when ignited, burns with a greenish-white flame, 
disseminating white vapors of antimonic oxide. A red heat 
decomposes it into antimony and hydrogen. In Marsh’s ap- 
paratus (Fig. 75, p. 141) it affords 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) and by its slight volatility in a 
current of hydrogen. 
When a solution of SbCl3 is decomposed by the galvanic current, metal- 
lic antimony separates and is strongly impregnated with hydrogen. It 
seems to form a sort of alloy with the latter. The same product is more 
readily obtained if antimony be made the negative electrode in the gal- 
vanic decomposition of water. When antimony in this condition is heated 
(or on breaking the rod) the hydrogen is separated with explosion (com- 
pare 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 144 
INORGANIC CHEMISTRY. 
by the solution of the oxide or sulphide in strong hydrochloric 
acid : 
Sb2S3 + 6HC1 = 2SbCl3 + 3lI2S. 
This solution is evaporated to dryness and the residue dis- 
tilled. 
It is a colorless, crystalline, soft mass (Butyrum Antimonii), 
fusing at 730 and boiling at 2230. Its vapor density equals 
112.8 (H = 1), corresponding to the molecular formula, 
SbCl3 = 225.7. In the air it attracts water and deliquesces. 
It dissolves unchanged in water acidified with hydrochloric 
acid. Much water decomposes it; the solution becomes tur- 
bid and a white powder—powder of algaroth—separates 
SbCl3 + H20 = SbOCl + 2HC1. 
The composition of this powder varies with the conditions 
under which it is formed, but generally corresponds to the 
formula 2 (SbOCl.). Sb203. Pure Antimony Oxychloride, 
SbQCl, obtained by heating SbCl3 with alcohol, occurs in 
colorless crystals and is further decomposed by water. 
While the metallic chlorides are not decomposed by water 
at ordinary temperatures, the ready decomposition of the 
halogen derivatives of antimony, shows that this element yet 
possesses a partial metalloid al character. 
Antimonic Chloride—Pentachloride—SbCl5, results from 
the action of an excess of chlorine upon antimony or the tri- 
chloride. It is a yellowish liquid which fumes in the air, 
becomes crystalline when cold, fuses at — 6°, and crystallizes 
in three different forms. Heat partially decomposes it, like 
PCI5, into SbCl3 and Cla: 
. SbCl5 = SbCl3 + Cl,. 
Water converts it into pyroantimonic acid (H4Sb20T), and 
hydrochloric acid. 
Antimony Tribromide—SbBr3—is a white, crystalline 
substance, fusing at 940 and distilling at 270°. The Tri- 
iodide, Sbl3, is a red compound, crystallizing in three distinct 
forms. 
X. vol. 
I vol. 
1 vol. 
• In the group of nitrogen, phosphorus, arsenic, and anti- 
mony, we must also include Bismuth—Bi = 207 ; it forms 
similarly constituted compounds, e. g., BiCl3, Bils, BiOCl. Its 
metallic character, however, considerably exceeds its metal- 
loidal. Thus, it does not unite with hydrogen, and the oxide ELEMENTS OF THE NITROGEN GROUP. 
145 
(Bi203), similar in constitution to the acid-forming As203, pos- 
sesses only basic characters. We will, therefore, consider bis- 
muth and its derivatives with the metals. 
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, just like the elements of 
the chlorine and oxygen group, and this gradation is inti- 
mately connected with the atomic weights. As the latter 
increase the substance condenses, the fusibility and volatility 
decrease, and the metallic character becomes more prominent: 
N. 
P. 
As. 
Sb. 
Atomic weight 
14.OI 
30.96 
74-9 
II9.6 
Specific gravity 
1.8-2.1 
47-5-7 
' 67 
Fusion point 
44° 
red-white heat 
Vapor density 
O.972 
4-32 
10.3 
Excepting bismuth, which is perfectly metallic in its nature, 
the elements of this group form gaseous compounds with three 
atoms of hydrogen. 
Ammonia (NH3) possesses strongly basic properties, and 
combines with all acids to yield ammonium salts; phosphine 
(PH3) combines with HBr and HI to form salt-like compounds. 
AsH3 and SbH3 no longer show basic properties. Arsenic and 
antimony, as well as the two preceding elements, combine with 
the hydrocarbons (e. g, CH3 and C2H5) and form compounds 
that are analogous in constitution and similar in character to 
the hydrides. These compounds [As (CH3)3 and Sb (CH3)3] 
will be described in Organic Chemistry; they possess basic 
properties and yield salts corresponding to the ammonium 
salts. 
The oxygen derivatives of these elements exhibit a similar 
gradation. With increase of atomic weight, corresponding to 
the addition of metallic character, the oxides that form strong 
acids in the lower series acquire a more basic nature. 
The thermal xelations 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 (HsN—gas) 11,890 calories are 
set free ; and yet the union occurs almost solely in an indirect way. This 
is probably explained by the fac* that the decomposition of ammonia (by INORGANIC CHEMISTRY. 
heat and electricity) is more easily brought about than the disgregation of 
the hydrogen and nitrogen molecules (compare p. 113). It appears (as in 
the group of the halogens and of oxygen) that the heat disengagement 
grows gradually less in the formation of PH3, AsH3 and SbH3. The in- 
creasing decomposability argues in favor of this view. On the other hand, 
in the formation of the chlorides with increase of atomic weights and 
metallic character the heat disengagement successively increases, corre- 
sponding to the symbols: 
(N,C13) = — 38,100 (P,C13) = 75.300 (As,C13) = 71,400 
(sb, eg = 91,400 
(P,C15) = 104,990 (P,CI3,0) = 142,600 (Sb, Cl5) = 104,000. 
The great instability of nitrogen chloride finds explanation in the re- 
markable heat absorption taking place in its formation. 
Less heat is produced in the formation of the iodides of the metals than 
with their chlorides. This is true also of phosphorus : 
(P.C1S) = (P,Brs) = 42,600 (P,I3) = 10,900. 
CARBON GROUP. 
The two non-metals, carbon and silicon, and the metal, tin, 
belong to this group. These unite with four atoms of hydro- 
gen or four of the-halogens. 
1. CARBON. 
Carbon occurs free 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 vegetable 
kingdoms, and is especially contained in the fossilized products 
arising from, the slow decomposition of vegetable matter—in 
turf, in brown coal, bituminous coal, and in anthracite. In 
combination with hydrogen it forms the so-called mineral oils 
—petroleum and asphaltum. It occurs, further, as carbon di- 
oxide (C0.2) in the air; and in the form of carbonates (marble, 
calcite, dolomite) comprises many minerals and entire rock 
formations. 
It is found in different allotropic modifications when free; 
these may be referred to the three principal varieties—dia- 
mond, graphite, and amorphous carbon. In all these forms it 
is a solid, even at the highest temperatures; non-fusible and 
non-volatile. This deportment can only be explained by the 
supposition that its free molecules are composed of a large 
number of carbon atoms combined with each other. (See p. 
104.) All the modifications of carbon are quite stable, but not 
very reactive. When burned all yield carbon dioxide. 
c = 11.97 CARBON. 
147 
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, that are mostly 
rhombic dodecahedra, rarely octahedra. Ordinarily, it is perfectly color- 
less and transparent; sometimes, however, it is colored by impurities. Its 
specific gravity equals 3.5. It does not soften any unless exposed to the 
most intense heat—between the poles of a powerful galvanic battery. It 
is then converted into a graphitic mass. When heated in oxygen gas it 
burns to carbon dioxide. It is scarcely attacked at all when acted upon 
by a mixture of nitric acid and potassium chlorate. 
2. Graphite is characterized by its oxidation to graphitic acid when it is 
heated with a mixture of potassium chlorate and nitric acid. Like amor- 
phous carbon, it is oxidized to mellitic acid by an alkaline solution of po- 
tassium permanganate, or when it is made the positive electrode in the 
electrolysis of alkaline solutions. Native graphite is found in the oldest 
rock formations, and of especially good quality at Altai in Siberia. It 
occurs, too, in considerable quantities at many places in the United States. 
It is occasionally found crystallized in six-sided forms, but usually as an 
amorphous, grayish-black, glistening, soft mass, used in the manufacture of 
lead pencils. The specific gravity is 2.25. It conducts heat and electricity 
well. When away from air-contact it is not altered even at the highest 
temperatures. It usually burns when heated in an atmosphere of oxygen, 
but with more difficulty than the diamond, forming carbon dioxide, and 
leaving about 2-5 per cent, of ash. To purify the poorer and more im- 
pure kinds of graphite, the latter is pulverized and heated with a mixture 
of KClOj and H2S04; the product is washed with water, and the residue 
ignited (Brody’s Graphite). 
Graphite may be obtained artificially by fusing amorphous carbon ydth 
iron; when the latter cools, 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 convert it in the cold 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, called metallic carbon, deposits in the manufacture 
of gas in the retorts, and is very hard, possessing metallic lustre, and con- 
ducting electricity well; hence its use in galvanic batteries. Coke, resulting 
from the ignition of bituminous coal, forms a sintered mass, conducting heat 
and electricity well. Charcoal is very porous, and can absorb many gases 
and vapors; 1 volume of it condenses 90 volumes NHS, 55 volumes H2S, 
and 9 volumes O2. At ioo°, and under the air pump, the absorbed gases 
are again liberated. Charcoal will also take up many odorous substances 
and decaying matter; hence is employed as a disinfectant. Animal Char- 
coalis obtained by the carbonization of animal matter (bones, blood, etc.), 
and possesses the power of removing many coloring substances from their 
solutions; hence it serves in the laboratory and in commerce for the de- 
colorizatkm of dark solutions. 
All these varieties o£ carbon contain smaller or larger quantities of ni- 
trogen, hydrogen, and mineral substances, which remain as ash after com- 
bustion. Hydrochloric acid will withdraw almost all the mineral con- 
stituents. INORGANIC CHEMISTRY. 
The fossil coal varieties, bituminous coal, lignite and turf,’ are the pro- 
ducts 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., and brown coal 70 per cent., of carbon. The fossil 
coal richest in carbon, the last product of the alteration, is anthracite. 
This has lost all its organic structure, and contains 96-98 per cent, of 
carbon. 
COMPOUNDS OF CARBON WITH HYDROGEN. 
With hydrogen, carbon forms an unlimited number of com- 
pounds, into which all other elements, especially oxygen and 
nitrogen, can enter. These derivatives of carbon have been 
termed organic compounds, because they were formerly ob- 
tained exclusively from vegetable and animal organisms, and 
the idea was entertained that they were produced by the in- 
fluence of forces other than those forming the mineral sub- 
stances. At present, most carbon derivatives are prepared ar- 
tificially'from the elements by simple synthetic methods; we 
are aware that they do not differ essentially 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 derivatives of carbon are so 
numerous, and possess so many peculiarities, that it appears 
necessary, from a practical stand-point, to treat them apart 
from the other compounds, in'a separate portion of chemistry, 
which we, pursuing the old custom, term organic chemistry. 
We then designate the chemistry of all other bodies as Inorganic 
Chemistry. .Only the simplest carbon-compounds will be con- 
sidered here. 
It is only under the influence of the electric arc that the 
direct union of carbon and hydrogen may be effected; the 
product is acetylene (C2H2). All other hydrocarbons are ob- 
tained indirectly in various ways. 
Methane—Marsh Gas—CH4.—This simplest hydrocar- 
bon, 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). It maybe obtained 
synthetically by conducting vapors of carbon disulphide and 
hydrogen sulphide over ignited copper filings: 
CS2 + 2ll2s + 8Cu = 4Cu.2S + CH4. METHANE. 
149 
For its- preparation, heat a mixture of sodium acetate with 
sodium hydrate: 
C2H3Na02 + NaOH = CHt + Na2-C03. 
Methane is a colorless, odorless gas, insoluble in water; it 
can be condensed by pressure and cold. When ignited it burns 
with a faintly luminous flame. It affords a violently explosive 
mixture (fire-damp of the miners) with two volumes of oxygen 
(or ten volumes of air). 
CH4 202 — C02 -f- -eH20. 
I vol. 
2 vols. 
I vol. 
2 vols. 
The quantitative analysis of methane shows that for every 
1 part of hydrogen in it there are 2.9 parts carbon. Were the 
formula CH (analogous to hydrochloric acid) then the atomic 
weight of carbon would be 2.9. If it corresponded to the 
formula of water (H20) then carbon would equal 5.98, etc. 
(see p. 95): 
MOLECULAR FORMULA OF METHANE. ATOMIC WEIGHT OF 
CARBON. 
H = i 
C = 2.99 
CH = 3.99 
2II = 2 
C = 5.98 
CH2 = 7.98 
3H = 3 
C = 8.97 
CH3 = 11.97 
4II = 4 
C = n.97 
CH4 = 15.97 . 
In this case the analysis yields (as in former instances) no 
Conclusive answer. We derive the molecular weight of methane, 
according to Avogadro’s law, from the density. The latter 
equals 7.98 (H = 1) hence the molecular weight is 15.97. In 
15.97 parts by weight of methane there are 4 parts by weight, 
hence 4 atoms, of hydrogen, and 11.97 parts carbon. The 
atomic weight of C is, then, presuming that only 1 atom of it 
is present in methane, 11.97. 
4 atoms hydrogen H4 = 4 
1 atom carbon C = 11.97 
Methane.molecule CH4 = 15.97 
That the atomic weight of carbon is really 11.97, is proven 
by the fact that of all its innumerable derivatives,'not one con- 
tains less than 11.97 parts, by weight, of this element. It follows, 
with certainty, from the periodic system of elements (p. 77). 
From the formula CH4 it follows that in 1 volume of methane 
there are 2 volumes of hydrogen (CH4 contains 2H2). This is 
I vol. 2 vols. 150 
INORGANIC CHEMISTRY. 
proved indirectly by the combustion of methane with oxygen 
in an eudiometer (see p. 119). Four atoms of hydrogen yield 
two molecules of H20 ; 1 atom of C yields 1 molecule of CO*. 
Hence the volume relation in the combustion of CH4 in oxygen 
is expressed by the equation : 
I vol. 
CH4 -f- *202 = C02 -f- 2H20. 
2 vols. 
I vol. 
2 vols. 
In two volumes of aqueous vapor there are 2 volumes of 
hydrogen ; hence in one volume of CH4, there are 2 volumes 
of H2. The result of the eudiometric analysis confirms these 
conclusions. 
Ethane—C2H6, is formed when hydrogen in statu nascendi 
acts upon ethyl chloride: 
C2H5C1 + 2H = C2H6 + HC1. 
Or by the action of potassium or sodium upon methyl iodide: 
2CH3I + Na2 = C2H6 + 2NaI. 
This is a colorless gas, insoluble in water, and when ignited 
it burns with a feebly luminous flame. Its density equals 14.97 
(H = 1) or 1.037 (air =1) corresponding to the molecular 
formula C2H6 = 29.94. 
Besides methane (CH4) and ethane (C2H6) there exists a 
long series of hydrocarbons of the general formula CnH2n+2* 
(e.g., C3H8, C4H10, C5H,2 etc.), in which each member 
from the preceding and next following by 1 C and 2 H (CH2). 
Bodies belonging to such a series greatly alike in their chem- 
ical behavior are termed homo/ogues. In addition to this 
series of saturated hydrocarbons others exist, with less hydro- 
gen, and by the addition of the latter, pass into the saturated, 
and may, therefore, be termed unsaturated. ■ The first unsatu- 
rated series is composed according to the formula CnH2„, the 
second according to CnH2n_2, etc. The lowest member of the 
series CnH2n is ethylene (see Chemical Structure, p. 169). 
Ethylene—C2H4, is formed in the destructive distillation 
of wood, bituminous coal, and many carbon compounds, hence 
is contained in illuminating gas. It is most easily obtained by 
the action of sulphuric acid upon alcohol, whereby the acid 
withdraws H20 from the latter: 
c2H6o - H20 = c2h4. 
Alcohol. 
Ethylene. It is a colorless gas, of weak, ethereal odor, and condenses 
at — iio° to a liquid.. Its density equals 13.97 (H = 1) or 
0.969 (air = 1), corresponding to the molecular formula 
C2H4 = 27.94. It burns with a bright, luminous flame, decom- 
posing first into marsh gas and free carbon : 
THE NATURE OF FLAME. 
15 I 
c2h4 = ch4 + c. 
The CH4 then burns and heats the particles of carbon in the 
flame to incandescence; these are then consumed to carbon 
dioxide (C02). 
The unsaturated compound, ethylene, unites directly with 
two atoms of chlorine and bromine : 
C2H4 + Cl, = c2h4gi2. 
The resulting compounds, C2H4C12 and C2H4Br2, are oily 
liquids; hence the name olefiant gas, for ethylene. 
The lowest member of the second unsaturated series is 
C 2H2. 
Acetylene—C2H2—is produced in the dry distillation of 
many carbon compounds, and is present in coal gas, to which 
it imparts a peculiar penetrating odor. Its density = 12.97 
(H =1) corresponding to the formula CaH2 25.94. It 
combines directly with 2 and 4 atoms, of chlorine and 
bromine. 
The three hydrocarbons considered above, methane (CH4), 
ethylene (C2H4), and in slight amount acetylene (C2H2), con- 
stitute, together with H and carbonous oxide (CO), ordinary 
illuminating gas, which is produced in the dry distillation of 
bituminous coal, lignite, or wood. The illuminating power is 
influenced by its quantity of ethylene and acetylene (and their 
homologues.) 
We are aware that every chemical union which occurs in a 
gaseous medium, and is accompanied by the evolution of light 
is designated combustion. * We observe herein, that some 
bodies, like sulphur and phosphorus, yield a flame when burned 
in the air or in other gas; such substances are converted 
into gases or vapors at the temperature of combustion. 
Carbon burns without a flame, becomes incandescent, because 
it is non-volatile. The carbon compounds, wood, bituminous 
coal, and tallow, are, indeed, not volatile, but burn with a 
THE NATURE OF FLAME. 152 
INORGANIC CHEMISTRY. 
flame because under the influence of heat they develop com- 
bustible gases. Flame is, therefore, nothing more than a com- 
bustible gas heated to incandescence. We have observed, too, 
that the combustibility is only a relative phenomenon ; if hy- 
drogen burns in oxygen and chlorine, oxygen and chlorine, 
conversely, will burn in hydrogen (p. 54). 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 ordi- 
nary lamp-chimney (Fig. 76) is closed at its 
lower end with a cork, through which two 
tubes enter ; 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 issu- 
ing 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 increasing the supply 
of gas, the flame becomes larger, the globe 
fills with illuminating gas, while the air is 
crowded out. The gas flame is. extin- 
guished, and an air-flame appears upon the 
wider tube, b, as the entering air continues 
to burn, in the atmosphere of illuminating 
gas. The excess of the latter escaping 
from the upper portion of the globe may 
be ignited, and we then have a gas-flame 
above, while within the globe we have an 
air-flame. On again lessening the gas flow 
the air-flame will distribute itself, extend 
to the exit of the tube a, and then the gas- 
flame will appear upon the latter, while the flame above the globe is ex- 
tinguished. In this manner, we may repeat the return process of flames 
at will. That the air actually burns in the air-flame may be plainly 
proved if we introduce a small gas-flame from c, through the wide me- 
tallic tube b; the little flame will continue to burn in the air-flame, but 
will be extinguished if it be introduced higher up into the atmosphere of 
illuminating gas. 
Fig. 76. 
We say ordinarily that only those bodies are combustible 
which, because of their power to unite with oxygen, burn in ;tn 
atmosphere of this gas or in air. If we imagine, however, an 
atmosphere of hydrogen, or illuminating gas, then bodies rich 
in oxygen must be combustible in these. In fact, nitrates, THE NATURE OF FLAME.’ 
153 
chlorates, etc., burn in an atmosphere of illuminating gas 
with‘the production of an oxygen flame. This maybe de- 
monstrated as follows: An Argand-lamp chipmey (Fig. 77) is 
closed at its lower end by a cork, bearing- a gas-conducting 
tube.. The gas which escapes through the opening of the sheet 
covering, a, is ignited. Then the substance (potassium or ba- 
rium chlorate, etc.) is introduced into the flame on an iron 
spoon provided with a long handle, heated to the temperature of 
Fig. 77. 
decomposition (disengagement of oxygen), and the spoon then 
plunged through the opening into the gas atmosphere. The 
substance burns with a brilliant light, as the resulting oxygen 
flame is brightly colored by the vaporizing and reduced me- 
tallic salts. 
The brilliancy or luminosity of a flame is influenced by the 
nature of the substances contained in it, also by its temperature 
and density. Incandescent gases shine very faintly per se; this 
is especially true when they are diluted. Thus hydrogen, am- 
monia, and methane burn with a pale flame. Even sulphur burns 
in the air with a flame slightly luminous. If, on the contrary, 
sulphur or phosphine be permitted to burn in oxygen, or 154 
INORGANIC CHEMISTRY. 
arsenic and antimony in chlorine gas, an intense display of 
light follows. This depends on the fact that the flame is not 
diluted by the nitrogen of the air, is therefore more condensed, 
develops a higher temperature, and thre combustion products 
(S02, P205, PC13) or the evaporating substances are not imme- 
diately gasified. That the density of the flame of gases exer- 
cises a great influence upon the luminosity is proved by the 
fact that hydrogen, compressed into a smaller space with oxy- 
gen burns writh intense light display.' 
A slightly luminous flame may be rendered intense by intro- 
ducing solid particles into it. For example, if hydrogen be 
passed through liquid chromium oxychloride (Cr02Cl2) it 
burns with a bright luminous flame, because the volatile 
Cr02Cl2 in it is changed by the oxygen of the flame into 
solid, Jion-volatile chromium oxide, Cr203, whose 
particles are heated to incandesence by the hydro- 
gen flame. The illuminating pow’ej of the various 
hydrocarbons and carbon compounds is similarly 
explained. Marsh gas, CH4, and ethane, C2H6, 
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 parti- 
cles glow in the flame. (See p. 151). 
Let us consider the flame of an ordinary stearin 
candle: On approaching the wick wuth a flame the 
stearin melts, is drawn up by the fibres and con- 
verted into gaseous hydrocarbons, which ignite, and 
by their chemical union with the oxygen of the air, 
produce the flame. The unaltered gases exist in the 
inner non-volatile zone a (Fig. 78); theycapnot burn 
because of lack of air access. If the lower end of a thin glass 
tube be inserted here the gases wall rise in it, and may be 
ignited at the upper end. There is a partial combustion of 
the gases in the middle, luminous part,/-, e, g; ethylene, C2H4, 
breaks up into CH4 and C: the first burns completely, while 
the C is heated to a white heat, because there is not sufficient 
oxygen present for its combustion. The presence of carbon 
particles in the luminous part may be easily proved by placing 
a glass rod or a wire net in it; it will at once be coated with 
soot. 
In the outer, very feebly luminous and almost invisible 
mantle, b, c, d, of the flame, which is completely surrounded by 
Fig. 78. THE NATURE OF FLAME. 
air, occurs the perfect combustion of all the carbon to carbon 
dioxide. 
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 lumi- 
nous but very hot flame. Upon this 
principle is based the construction 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, c, is 
screwed into the lower portion, and 
in the figure is only separated for the 
sake of explanation. The gas enters 
through the narrow opening, a, from 
the side gas tube, and mingles with 
air in the tube c, which enters 
through the openings of the ring, b. 
In this way we obtain a flame that is 
but faintly luminous, although afford- 
ing an intense heat. 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 
contains an 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. 
The construction and application of the ordinary and the 
gas blowpipes depend upon the same occurrences; they are, 
however, replaced at present by gas lamps. 
Fig. 79. 
The non-luminosity of the Bunsen burner flame, due to addition of air, 
depends on a more complete combustion of the separated carbon or of the 
yet undecomposed hydrocarbons. The flame, in consequence, is smaller, 
more intense, and the combustion extends itself even to the inner cone of 
the flame. It is more difficult to render the flame non-luminous by pure 
oxygen, because it is then not diluted by nitrogen, and is, therefore, much 
smaller, the temperature much higher, and the flame gases are more con- 
densed. 
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, 156 
INORGANIC CHEMISTRY. 
takes place only in the outer cone. In consequence of the dilution there 
are present fewer combustible particles in an equal space, and these can be 
more completely consumed by the oxygen of the air, which enters more 
readily ; further, the temperature is lowered, and probably does not acquire 
the decomposition temperature of ethylene (C2H4) in the adjoining cone, 
which is being continually renewed. The simple extension of an illumin- 
ating flame upon a plate, will render it non luminous, because then the 
air comes in contact with a larger flame surface. On heating a gas made 
non-luminous by the admixture of nitrogen, and then letting it burn, its flame 
becomes luminous because the increased temperature can induce the de- 
composition of ethylene. 
In rendering flame non-luminous by carbon dioxide, we must also con- 
sider that the same is converted, by the particles of carbon, into carbon 
monoxide: 
C02 + C =’ 2CO. 
Indeed, but a few per cent, of C02 in a gas-flame suffices to consider- 
ably diminish its luminosity, 
I vol. 
c2h4 4- co, = ch4 + 2C0, 
I vol. 
I VoL 
2 vols. 
while the presence of nitrogen is far less detrimental. 
Every substance requires a definite temperature for its igni- 
tion—temperature of ignition. When a substance is once 
ignited it generally burns further, because additional particles 
Fig. 8i. 
Fig. 8o. 
are raised to the temperature of ignition by the heat of com- 
bustion. By rapid cooling (e.g., by the introduction of a 
piece of metal into a small flame) every flame may be extin- CARBON WITH THE HALOGENS. 
guished. 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- 
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. 
157 
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, PHS, etc., in such manner as to 
unite with the hydrogen to form hydrogen chloride, etc., while 
the other element is either set free or is also combined with the 
chlorine. Chlorine and bromine act similarly upon the hydro- 
carbons ; here hydrogen is displaced, atom after atom, by chlo- 
rine, forming HC1 and chlorine derivatives: 
CH4 + Cl2 = CH3C1 + HC1 
CH4 + 2C12 = CH2C12 4. 2HCI 
CH4 + 3C12 = CHClg + 3HCI, etc. 
Such a process is termed substitution, and the products substitu- 
tion products. In this way we obtain from methane, CH4, the 
products, CHgCl, CH2C12, CHC13 (<chloroform) and finally 
CC14, carbon tetrachloride: The last compound is a color- 
less, ethereal liquid, boiling 770. Its. vapor density equals 
76.7 (H = 1), corresponding to the molecular formiila.CCl4 
= 153.4. The chlorides of carbon are not decomposed by 
water. 
The compound C2C16, hexachlorethane, obtained by the ac- 
tion of chlorine upon ethane, 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 1220 G. Bromine 
and iodine yield similar compounds; they will be treated more 
extensively in organic chemistry. 158 
INORGANIC CHEMISTRY. 
The heats of formation of the previously mentioned hydrocarbons 
(from amorphous carbon and hydrogen) are deduced from the heat of 
combustion, and equal: 
(C,H4) =21,750 (C2,II6) = 25,670 (C2,H4) = — 2710 (C2,H2) = — 48,290. 
The absorption of heat in the formation of acetylene and ethylene is 
explained by the fact that the solid carbon molecules must first be gasified 
and separated into atoms in order to unite with hydrogen. The disgrega- 
tion heat requisite for this equals very probably upwards of 38,300 C. for 12 
parts by weight of amorphous carbon (76,600 C. for 24 parts); then the 
preceding heats of formation must be enlarged about this much, in order 
to express the true heal of formation (from gaseous carbon atoms) of the 
hydrocarbons. 
Beginning with acetylene, C2H2, we discover that its conversion into 
ethylene, C2H4, and into ethane, C2H6, ensues with heat disengagement: 
In accordance with this acetylene readily unites with hydrogen (in statu 
nascendi, or by action of platinum sponge, etc.), forming ethylene and 
ethane. 
' The heats of formation of the carbon chlorides approach those of the 
hydrogen derivatives very closely : 
(C2H2,H2) = 45,580 (C2H4,H2) = 28,380. 
(C,C14—gas) = 21,030 (C2,C14—gas) = —1150. 
The affinity of chlorine for carbon is, therefore, very nearly the same as 
that of hydrogen. 
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 a free condition. Combined with oxygen as silicon 
dioxide (Si02), and in the form of salts of silicic acid (silicates) 
it comprises many minerals and almost all the crystalline rocks. 
It may be obtained in a free condition by heating silicon 
fluoride (SiFl4), or sodium-silicon fluoride (Na2SiFl6) with 
metallic sodium: 
The ignited mass is treated with water, which dissolves the 
sodium fluoride and leaves the silicon as a brown, non-lustrous, 
amorphous powder. It burns to silicon dioxide (Si02) when 
heated in the air. 
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 consists of black, shining 
Na2SiFl6 + 4Na = 6NaFl -f Si. HYDROGEN SILICIDE. 
159 
octahedra, of specific gravity 2.49, and of very great hard- 
ness. Upon ignition in the air or oxygen it is pot oxidized ; 
it is not attacked by acids. On boiling it with a sodium or 
potassium hydrate solution it dissolves, forming a silicate and 
liberating hydrogen: 
Heated in chlorine gas, silicon burns to the chloride. 
Hydrogen Silicide—SiH*—the analogue of CH*, is pro- 
duced like arsine by dissolving an alloy of silicon and mag- 
nesium in dilute hydrochloric acid : 
Si -f 4KOH = K4Si04 + 2Ha. 
SiMg2 + 4HCI = SiH4 + 2MgCl2. 
The escaping hydride contains admixed hydrogen, has a 
disagreeable odor, ignites spontaneously in the air, and burns 
to the dioxide and water: 
SiH4 + 202 = Si02 + 2H20. 
Perfectly pure silicide, free of H, is obtained by heating a 
compound, SiH (O. C2H5)3, which will be treated in Organic 
Chemistry. In the air at ordinary pressure, it ignites only upon 
Fig. 82. 
warming; if, however, the gas, by diminution of pressure 
or by the addition of H, is diluted, it becomes spontaneously 
combustible at ordinary temperatures. A red heat decomposes 
the hydride into amorphous silica and hydrogen. When mixed 
with chlorine it inflames and probably forms substitution pro- i6o 
INORGANIC CHEMISTRY. 
ducts similar to those of methane (CH4). Pure hydrogen silicide 
condenses to.a liquid, at -50 and a pressure of 70 atmospheres. 
Silicon Chloride—SiCl4—results from the action of chlo-. 
rine upon silicon, or by conducting chlorine over an ignited 
mixture of the dioxide and carbon (Fig. 82) : 
SiOa -f- 2C + 2OI2 = SiCl4 + 2CO. 
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 separately upon the Si02, when they act simultaneously 
the reaction is induced by the mutually supporting affinities of 
carbon for oxygen and of chlorine for silicon. 
The silicon chloride which distils over is a colorless liquid, 
having a specific gravity of 1.52, and boiling at 570. It fumes 
in the air, and is decomposed by water into silicic and hydro- 
chloric acids: 
SiCl4 -(- 4II2O = H4Si04 + 4HCI. 
This compound may be employed in determining the atomic weight of 
silicon. Analysis shows that it contains seven parts silicon for every 35.37 
parts of chlorine. Supposing, from the great analogy of the silicon com- 
pounds to those of carbon, that its formula is SiCl4, the atomic weight of 
the silicon would be 28. 
Si = 28 
Cl4 = 141.48 (4X35-37) 
SiCl4 = 169.48 
This supposition is confirmed by the vapor density of the compound. 
This equals 84.7 (H — 1), hence the molecular weight is 2 X 84.7 = 169.4. 
As the analysis shows that there are 141.4 parts chlorine in 169.4 parts of 
silicon chloride, the atomic weight of the metal must be 28. 
Silicon Bromide—SiBr4, and Silicon Iodide—Silj, are 
formed in the same manner as the chloride.- The first is a 
colorless liquid, of specific gravity 2.8, becoming solid at —120 
and boiling at -{- 1530. The iodide forms colorless octahedra, 
fusing at 120°, and boiling at 290°. Like the chloride, both . 
are decomposed by water. 
Besides these compounds, which may be viewed as hydrogen silicide, 
in which all the hydrogen is replaced by halogens (see page 157), others 
exist, in which only a part of this element is replaced. Thus, silicon 
chloroform, SiHCl3, corresponds .to the chloroform (CHC1S) derived from 
methane. It is produced by the action of phosphorus pentachloride, or 
antimony chloride, on hydrogen silicide: 
SiH4 + 3SbCl6 = SiHClj + 3SbClg + • 3HCI. SILICON FLUORIDE. 
Or upon heating silicon in dry hydrogen chloride gas: in this case a mix- 
ture of SiCl4 and SiHCl3 results.. These compounds may be separated by 
fractional distillation. Silicon chloroform is a colorless liquid, of specific 
gravity l.6tand boils at 35°~37°. The vapor density equals 67.5 (H = 1), 
corresponding to the molecular formula, SiHCl3 = 135.11. It fumes in 
the air, and decomposes with water into silicic and hydrochloric acids. 
The silicon bromoform, SiHBrs, and iodoform, SiHJ3, are very similar to 
the chloroform ; these correspond to the analagous carbon compounds. 
The compound Si2I6, analagous 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 
SiO • 
Si02 + 4HFI = 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. 62) is dis- 
engaged, and this 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 3.60 (air = 1), 
or 52 (H = 1), corre- 
sponding to the mole- 
cular formula SiFl4 = 
104. Its deportment 
with water is very char- 
acteristic; it is decom- 
posed thereby into sili- 
cic acid (H4Si04) and 
hydrogen silico-fluo- 
ride : 3SiFl4 + 4H20 
= H4Si04 + 2H2SiFl6. 
For the execution 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 SiH2Fl6 
remains dissolved in the water. As the separating silicic acid 
may easily obstruct the opening of the glass tube, the latter is 
allowed to project a slight distance into mercury. The solid 
silicic acid is separated from the aqueous solution by filtration. 
Fig. 83. 162 
INORGANIC CHEMISTRY. 
Hydrogen Silico-Fluoride—H2SiFl# (or 2HFl,SiFl4) is 
only known in aqueous solution. Upon evaporating at a low 
heat it decomposes into SiFl4 and 2HFI. 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. 
From its chemical nature we must include Tin, Sn = 117.5 
in the $ame group as carbon and silicon. It forms perfectly 
analogous compounds, as SnCl4, SnBr4, SnFl4. It bears the 
same relation to the above elements as antimony does to those 
of the nitrogen group. This fact is plainly visible in the 
atomic weights: . 
N = 14.0 P = 30.9 As = 74.9 Sb = 119 6 Bi = 207. 
C = 11.9 Si = 28. Sn = 117.5 Pb = 206.4. 
The element of the carbon group corresponding to the arsenic 
of the nitrogen group, if it does indeed exist, is unknown. 
This explains the abrupt break from silicon to tin. Tin is 
decidedly more metallic in its nature than antimony, and forms 
the transition to the real metals. In its physical properties tin 
is a true metal: unlike the metalloids it does not unite with 
hydrogen. It and its higher analogue, Lead, will, therefore, 
be treated with the metals. 
lit the preceding pages we have considered four groups of 
elements, comprising all the so-called metalloids (with the 
exception of boron). In each group the last members, pos- 
sessing the highest atomic weights, exhibit very distinct metal- 
lic properties, especially when in their free state. This is 
clearly 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. The affinity for hydrogen, on the other hand,-dimin- 
ishes with increase in metallic character; the hydrides of 
iodine, tellurium, antimony, and arsenic, are very unstable 
and decompose readily into their constituents; finally, tin and 
bismuth do not combine with hydrogen. 
The remarkable relations between the atomic weights of the ATOM AND MOLECULE. 
163 
elements of the four groups are seen in the following table. 
The deportment of the elements is expressed by the same. 
C = 11.9 N = 14.0 O = 15.9 'FI = 19.. 
Si = 28. P = 30.9 S = 31.9 Cl = 35-37- 
As = 74.9 Se = 78.9 Br = 79.76. 
Sn == 1175 Sb = 119.6 Te = 126. I = 126.54. 
We will give a fuller consideration of these relations in the 
presentation of the Periodic System of the Elements. 
ATOM AND MOLECULE. 
In the study of the hydrogen derivatives of the elements of 
the four groups already mentioned, we arrived at the establish- 
ment of the following formulas: 
CH4. 
SiH4. 
NH3. 
SHs. 
AsHj. 
SbH3. 
OH,. 
sh2. 
SeH2. 
TeH2. 
F1H. 
C1H. 
BrH. 
-IH. 
These formulas possess a fundamental importance, and serve 
for the deduction of other more important conclusions and 
generalizations, and for this reason we again here introduce in 
a connected form all the actual relations and considerations that 
were the basis of their derivation. 
The fact, proved by the analysis and synthesis of chemical 
bodies, that the elements combine with each other according 
to constant multiple proportions, is most simply explained by 
the supposition of elementary atoms, which can combine 
among themselves in different quantity (page 6"8). "The exist- 
ence of isomeric bodies can only be concluded from the atomic 
constitution of matter (p. 86). The experimental combining 
weights, however, offer no assistance in determining the 
number of atoms in a compound, and, therefore, the true 
relative atomic weights of the elements (p. 69); the ques- 
tion can only be solved upon a basis of other actual relations. 
The specific gravities of the bodies in the gaseous and vapor 
form, and the relations according to which the gases combine 
by volume render an important service in 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, an equal number 
of molecules is contained in equal volumes of all gases. Hence 
molecules are the smallest particles of matter—masses—which 164 
INORGANIC CHEMISTRY. 
occur separated in gases; their relative weights, therefore, are 
directly expressed by their relative gas densities (p. 77). 
Molecules of compound bodies are composed of atoms; there- 
fore, the relative atomic weights may be derived from the mole- 
cular weights. The atomic weight is the smallest quantity of an 
element which is contained in the molecular weight of a?iy one 
compound. Although the atomic weights of the above con- 
sidered elements have been deduced from a small number of 
compounds, yet all further investigations have confirmed the 
observation that these elements are not contained in a smaller 
quantity in any molecule. Thus oxygen is present in all mole- 
cules, in at least 15.96 parts by weight; nitrogen not less than 
14.01; carbon not less than 11.97—if the smallest quantity of 
hydrogen contained in a molecule, i. e., an atom, equals x. 
From a comparison of the densities of the elements with 
those of their derivatives we concluded (p. 75)—upon the 
basis of the proposition, that an equal number of molecules is 
contained in equal gas-volumes—that the molecules of the 
elements consist of two or more atoms. Thus the molecules 
of hydrogen, 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 Cl2 02 N2 S6 P4 As4. 
The molecular quantities of all bodies in the gaseous state 
occupy equal volumes. Referring the densities (specific grav- 
ities) to hydrogen as a unit (=i), the molecular weights are 
double the densities, and conversely, the latter are half the 
molecular weights (p. 76). 
The existence of allotropic modifications of the elements 
confirms (p. 86) the hypothesis that the molecules of the ele- 
ments are composed of several atoms. In the gaseous condi- 
tion the allotropy is known only in the case of oxygen and 
sulphur. Ordinary oxygen has two atoms, ozone three ; the 
sulphur molecule, at rooo°, contains two atoms; at 500°, how- 
ever, six atoms. Although we possess no means of determin- 
ing the molecular value 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 complex groups 
of atoms. Further, the energetic action of the elements in 
the moment of their formation (see p. 76) argues for their 
complexity when free. 
The phenomena of the so-called status nascens find their explanation, 
in the majority of cases, in the thermo-chemical changes accompanying ATOM AND MOLECULE. 
165 
them. If, e. g., in the action of zinc upon nitric acid, where, from analogy 
to other acids, hydrogen must be set free, the same does not occur, but 
the nitric acid is reduced to ammonia and nitrogen oxides—while free 
hydrogen does not exert any action on nitric acid—then it is the heat dis- 
engaged by the formation of zinc and ammonium nitrates, which, in ac- 
cordance with the principle of greatest heat liberation, represents the true 
cause of the reaction. However, the supposition that the free atoms act 
more energetically than the molecules is not refuted by this; 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 
purely chemical procedure, leading to the same end. As mole- 
cules are the smallest quantities which can exist in a free con- 
dition, it is very, probable the same quantities appear in the 
chemical reactions. Indeed, the study of the latter leads us 
to the same molecular quantities as are derived from the den- 
sities. For example, the compound CH3C1 results when chlo- 
rine acts upon marsh gas ; we hence infer that four atoms of 
H are present in the molecule of the latter hydro-carbon : 
From ethylene—C2H4—we obtain C2H:jCl; the composition 
of the former, therefore, cannot be-expressed by the simpler 
formula, CH2. With iodine, ammonia yields NHI2; with hy- 
drogen chloride NH4C1 : 
CII4 -f Cl2 = CH3C1 -f HC1. 
NII3 + 2l2 = NHI2 + 2III. 
NH, + HC1 = NH4C1. 
Both reactions prove that three atoms of hydrogen are pres- 
ent in ammonia. The compound nature of the elementary 
molecules is concluded in the same manner; 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 following equations : 
MnO,2 + 4HCI = MnCl2 + 2H20 + Cl2) 
NH+N02 = 2H2G + N2, 
Ag20 H- H202 = Ag2 4- H20 + 02. 
Although this chemical method for obtaining the molecular 
value is less simple, it is much more general than the one de- 
pending on the determination of the densities, because it can be 
applied in the case of the non-volatile bodies. Thus, for example, 
the composition of hydrogen peroxide is expressed by the for- 
mula, HO ; the methods of its production and chemical rear- 
rangements point, however, with great probability, to the 
doubled molecular formula, H202. 166 
INORGANIC CHEMISTRY. 
Hence the most varied relations carry us to the same con- 
clusions as to the existence-and the quantity or value of the 
atoms and molecules. The atomic-molecular formulas used in 
chemistry to-day are only the expression of actual relations; 
they are entirely independent of speculative abstractions, but 
forcibly call out these. 
Important generalizations result from the molecular formulas 
deduced above. We have the four following groups of hydro- 
gen compounds (forms of combination or types) : 
THE VALENCE OF THE ELEMENTS. CHEMICAL STRUCTURE. 
CH4 
SiII4 
NH3 
PH3 
AsHs 
SbII3 
OH2 
SH2 
SeH2 
TeH2 
F1H 
ClH 
BrH 
IH 
In each group the affinity of the elements for hydrogen di- 
minishes, step by step, with increasing atomic weight. Yet 
the number of hydrogen atoms which are combined with one 
atom of the other elements is constant for each group. Hence, 
we must ascribe a particular function of affinity to each ele- 
ment, in its relation to hydrogen. This is called its valence or 
atomicity. The elements of the fluorine group are univalent or 
monatomic ; the elements of the oxygen group bivalent or dia- 
tomic : nitrogen and its analogues, trivalent; carbon and silicon, 
finally, are tetravalent elements—if we accept the valence 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 valence of the elements is frequently designated by 
lines or Roman numerals placed above the symbols : 
I II III IV 
Cl O N C 
The mutual union of affinity units is indicated by one line: 
H 
I 
N 
/ \ 
H H 
H 
I 
II—C—H 
I 
H 
H—Cl 
Hydrogen 
chloride. 
H—U—H 
Water. 
In these formulas the atoms of oxygen, of nitrogen, of car- 
bon—indeed, of all the multivalent elements—represent, as it 
were, the nuclei to which the hydrogen atoms attach them- 
Ammonia. 
Methane. VALENCE. 
167 
selves ; their valence units are, as it were, the points of attach- 
ment for the valence unit of hydrogen. 
The hydrogen atoms in these molecules can be replaced or 
substituted by other elements (p. 158). According to this, 
the monovalent halogen atoms each'replace one hydrogen 
atom: 
H 
II/ 
o 
\ 
H 
H 
II/ 
o 
\ 
Cl 
Cl 
H/ 
o 
\ 
Cl 
H 
III/ 
N 
\ 
I 
Cl 
III/ 
Sb—Cl 
\ 
Cl 
Water. 
Hypochlorous 
acid. 
Chlorine 
iodide. 
Nitrogen 
iodide. 
Antimony 
trichloride. 
It is more convenient to employ brackets instead of the 
lines : 
r h 
IV j Cl 
C Cl 
Cl 
f Cl 
IV j Cl 
C j Cl 
Cl 
r « 
IV cl 
Si j Cl 
Cl 
Chloroform. 
Carbon 
tetrachloride. 
Silicon 
chloroform. 
By the replacement of hydrogen in the preceding hydro- 
gen compounds by the monovalent metal potassium, we obtain 
K—Cl 
Potassium chloride. 
H 
/ 
o 
K. 
r H 
N K 
(k 
c{h 
IS 
Di-pot. amide. 
Potassium hydroxide. 
Potassium methyl. 
The bivalent elements, like oxygen, sulphur, replace two 
hydrogen atoms in the compounds of that element: 
Ill/Cl 
Sb=zO 
IV /H 
C—H 
/O 
IV/ On 
C\0n 
IV r On 
Sijo11 
Antimony 
oxychloride. 
Carbon 
dioxide. 
Silicon 
dioxide. 
Finally, trivalent nitrogen can replace three atoms of hy- 
drogen, or three of the halogens : 
Methylene 
oxide. 
III 
IV r N 
c{H 
or, 
I IV III 
H—C=N 
Ill 
IV f N 
c{ci 
Hydrogen cyanide. 
I IV III 
Cl—C= N 
Chlorcyanogen. 
or, 
In all these compounds the same valence peculiar to the 
elements appears. 168 
INORGANIC CHEMISTRY, 
Like valence is designated by the word equivalence. One atom of .chlo- 
rine is equivalent to one atom of hydrogen; 35.37 parts, by weight, of 
chlorine are, then, equivalent to 1 part, by weight, of hydrogen. One 
atom of oxygen is equivalent to two atoms of hydrogen; consequently, 2 
parts, by weight, of hydrogen are equivalent to 15.96 parts, by weight, of 
O, or iH to 7.98 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 XT4 == 4.66 parts nitrogen, etc. These quantities, equal to x 
part, by weight, of H are termed equivalent weights, and were formerly 
employed instead of the atomic weights. As may be observed from the 
preceding, the equivalent weights are parts of the atomic weights corre- 
sponding 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 resulting groups 
or residues— 
II 
—O—H 
II 
—S—H 
III 
—n=h2 
IV 
-c=h3 
Hydroxyl. 
Sulphydrate. 
Amide, 
Methyl. 
can plainly not exist in a free condition, as the one affinity 
unit of the element having higher valence, is not saturated. 
When set free, these groups, therefore (like the elementary 
atoms) unite with the free affinities and enter into complicated 
compounds. Thus, for example, we obtain the bodies: 
II II 
HO—OH 
Hydrogen 
peroxide. 
II II 
HS—SH 
Hydrogen 
persulphide. 
III III 
H2P— PH2 
Liquid 
hydrogen 
phosphine. 
IV IV 
H3C—CH3 
Dimethyl 
or 
ethane. 
Carbon is particularly inclined to such combinafion. Upon 
removing an atom of H from dimethyl or ethane (C2H6) the 
so-called ethyl group remains: 
IV 
IV 
in which one carbon affinity is unsaturated; this can again 
unite with the methyl group CH3. The resulting compound is: 
C2H5 or CHS - CH2 - 
C3H8 or H3C — CH2 — CH3. or C2H5. CH3 . 
IV IV IV 
By the continuation of this process of a chain-like union, as 
it were, of the carbon atoms, we obtain a whole series of hy- 
drocarbons (C4H10, C5H)2, etc.) with the general formula 
CnH2n+2 (compare page 150). . 
Not only similar residues or groups, but dissimilar also, com- 
bine in this way: CHEMICAL STRUCTURE. 
169 
Ill II 
H2N — OH 
IV II 
H3C — OH 
IV III 
H3C — NH,. 
• Hydroxylamine. 
Methyl hydroxide. 
Methylamine. 
Such combinations are generally effected by reactions of double decom- 
position. Thus methyl hydroxide (wood-spirit) results from the action of 
methyl Iodide (CH3I) upon silver hydroxide (AgOH): 
CH3I + AgOH = Agl + CHjOH; 
methylamine by the action of methyl iodide and ammonia: 
Dimethyl is produced when sodium acts upon methyl iodide: 
CHjI + NH3 = CH3NH2 + HI. 
2CH3I + Na4 = 2NaI -f C2H6. 
By the withdrawal of sodium the methyl groups are liberated, and then 
combine with each other. 
'Further, the atoms with several valences can unite with two 
and three affinities (double and triple union): 
Ill III 
N = N 
\/ 
O 
IV IV 
HaC = CH, 
IV IV 
HC = CII 
Ethylene 
(p. 150.) 
Acetylene, 
(p. 15M 
Hyponitrous 
oxide. 
By complete mutual union result the free elements : 
H — H 
Hydrogen. 
0 = 0 
Oxygen. 
N = N 
Nitrogen. 
p = p 
I I 
p — p 
This manner of linking or combination of the atoms in the 
molecule, not with their entire mass but with single affinities, 
is designated chemical constitution ox chemical structure of com- 
pounds; the formulas representing them are called constitution 
or structural formulas. Of course the actual position of the 
atoms in space (of which we have no knowledge) is not indi- 
cated by the chemical structure. The fundamental principle 
of chemical structure consists in this, that the affinity unit of 
one atom unites itself with the affinity unit of another atom. The 
following circumstances, however, complicate these simple re- 
lations : Among the elements of the nitrogen group we saw 
that P and N combine with 3 and 5 atoms of chlorine and the 
other halogens; that sulphur, selenium, and tellurium take up 
2 and 4 atoms of chlorine and bromine; that iodine unites 
with 1 and 3 atoms of chlorine and 5 of fluorine. Only the 
tetravalent elements, carbon and silicon, are capable of com- 
bining with 4 hydrogen atoms, and with not more than 4 of 
the halogens: 
Phosphorus. 170 
INORGANIC CHEMISTRY. 
IV 
CC14 
III 
PC13 
V 
PC15 
II 
SC12 
IV 
SC14 
I 
IC1 
III 
ICI3 
V 
IF15 
Hence, it appears that the elements (excepting carbon) do not 
express such a constant valence in their relation to chlorine (and 
to the halogens) as they do to hydrogen. P and its analogues 
appear to be tri- and pentavalent; the elements of the sulphur 
group, di- and tetravalent; iodine finally appears to be mono-, 
t-ri-, and pentavalent. 
This higher valence of the metalloids shows itself more dis- 
tinctly and frequently in the more stable oxygen derivatives. 
We are acquainted with the following oxygen compounds of 
the elements of the four groups already mentioned ; the mem- 
bers of each group afford perfectly analogous compounds: 
IV 
co2 
III 
NA 
V 
NA 
ii 
(SC12) 
IV 
so2 
VI 
so3 
ci2o 
III 
ciA 
V 
iA 
VII 
IA 
Here the valence of iodine (and of the halogens) reaches 
seven, and the elements of the sulphur group six affinity 
units. The elements of the nitrogen group are not more than 
pentavalent, both in their relation to chlorine and oxygen. 
Carbon, finally, does not show more than four affinities for H, 
Cl, or O. As we will observe later, these relations are more 
apparent in the hydroxyl derivatives of the oxides, the acids. 
Hence we conclude, that valence is not an absolute property 
belonging per se to and for ■ the elements solely (like the atomic 
weights) but that it appears as a function of the mutual action 
of the various elements. We can, in general, distinguish 
two valences; the hydrogen valence and the halogen or oxygen 
valence. The hydrogen valence is constant for all elements; 
for Cl = i, for 0= 2, for N = 3, for C = 4. The valence 
of the most of the elements appears to be variable as regards 
oxygen and chlorine. It varies, indeed, as seen from the for- 
mulas above given for 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 the maxi- 
mum valence from C = 4 to Cl = 7 argues for the actual oc- MOLECULAR COMPOUNDS. 
171 
currence of a change of valence. It is perfectly immaterial 
whether we ascfibe a variable valence to the elements or accept 
the maximum as the’true measure, and regard the lower com- 
pounds as unsaturated, because we possess no conceptions upon 
the nature of valence. Later, we will discover that this altera- 
tion of valence finds full expression and generalization in the 
periodic system of the elements, which is based upon the group- 
ing of the elements according to their atomic weights. 
In the manner just represented, the idea of valence is viewed from a 
purely empirical standpoint. Another opinion prevails. This, denying 
the alterability, regards the valence as an absolute constant property of the 
elementary atoms. According to this idea, the.true valence, or atomicity, 
is only derivable from the hydrogen compounds; the halogens are abso- 
lutely monovalent, the elements of the O group divalent, and those of the 
N group trivalent, etc. Different suppositions are proposed in order to 
carry out the constant atomicity for all compounds. A chain-like union is 
assumed in order to explain the oxygen compounds. This is similar to 
the C atoms in the carbon compounds. It is represented in the following 
examples: 
I II II II I 
Cl—0—0—O—Cl 
I II II II II I 
Cl—0—0—0—0—H 
Chlorine trioxide. 
y/? 
s I 
\o 
n/°\ 
s o 
Perchloric acid. 
n/o—o—n 
S\ 
No—O—H 
Sulphur dioxide. 
Sulphur trioxide. 
Sulphuric acid. 
II III II II II III II 
O = N—O—O—O—N = 0 
Nitrogen pentoxide. 
II III II 11 I 
O = N—O—O— II 
Nitric acid. 
Hence it would appear that the oxygen atoms can unite without end, in 
a chain-like manner, similar to the C atoms in the carbon compounds. 
However, a maximum combining affinity of different groups for oxygen 
does actually occur; this varies regularly from group to group (for Cl = 7, 
for S = 6, for N = 5, for C = 4). Consider it as-we may, on the suppo- 
sition of constant affinity, the reason for the different number of linking 
O atoms must be based on the nature of the other element. The fact, then, 
that the highest oxides or their hydrates (HC104, H2S04, HNOs) are more 
stable than the lower (p. 181), argues decidedly against the chain-like 
union of the oxygen atoms, such as occurs in. the unstable peroxides. 
To explain the other compounds according to the constant atomicity 
theory, a difference is conceived between atomic and molecular compounds. 
The former are such as can be explained by constant atomicity. All others 
are regarded as molecular compounds, resulting from the union of two or 
more molecules, upon the basis of newly acquired molecular affinities. 
Thus the compounds PC15, SC14, ICI3 are viewed as addition products of 
atomic compounds with chlorine molecules: 
PCI,, Cl2 SCI,, Cl2 IC1, Cl2. 
Phosphorus 
pentachloride. 
Sulphur 
tetrachloride. 
Iodine 
trichloride. 172 
INORGANIC CHEMISTRY. 
It was thought that a proof that these bodies were differently constituted 
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 molecules PC15, SC14 is only 
gradual, increasing with the temperature, and that they do exist at lower 
temperatures, undecomposed, in a vapor form (compare p. 137). True 
atomic compounds, like H2S04 and HNO3, frequently separate into simpler 
molecules, in their conversion into vapor (compare Sulphuric Acid). 
Phosphorus pentafluoride, PF15, a gas at ordinary temperatures, is a 
strong aigument in favor of the pentavalence of phosphorus ; iodine, too, 
affords a volatile pentafluoride, IF15. It is noteworthy that, as a general 
thing, the metalloids yield more stable and higher compounds with the 
lower halogens (fluorine and chlorine) than with bromine and iodine, 
which possess higher atomic weights (p. 51). The idea that molecular 
compounds can also exist in vapor form, would, render their distinction 
from the atomic compounds purely arbitrary—not present in the nature of 
the bodies themselves. But other gaseous compounds exist, which can, in 
no light, be regarded as molecular. Thus, the usually tetravalent or hexa- 
valent tungsten (WC16, WOCl4) forms a gaseous pentachloride, WC15, and 
molybdenum, perfectly analogous to tungsten, yields a pentachloride, MoC15. 
Further, the pentavalentrvanadium (VdOCl3) gives rise to a gaseous tetra- 
chloride, VdCl4. 
The salts of ammonia, according to the theory of constant atomicity, 
are not regarded as ammonium compounds 
V V V 
NH4Q, no3nh4, S04(NH2)2, 
(p. 127)', but as addition products of ammonia with acids, 
NHj.HCl, NHs,N03H, (NH3)2,S04H2, 
which would make the analogy of the same with metallic salts appear 
rather doubtful. Further, the properties of many compounds like 
V V IV IV V 
POCl3, (C2H5)3PO, (CHs)3S.OH, (CH3)2SO, (CH3)4N OH, 
and others, cannot be interpreted by the constant valence theory. The 
existence of potassium permanganate, Mn04K, is not compatible with a 
constant di- or tetravalence of manganese. 
Yet, up to the present, the acceptance of molecular additions could not 
be entirely dispensed with, especially for the so-called water of crystalli- 
zation compounds ; the proposed effort, however, continues to deduce all 
such compounds on the basis of higher valence of the elements. While, 
consequently, the constant valence theory comprises only the so-called 
atomic compounds, the extended valence idea' draws all others into the 
circle of generalizations. 
We must, first of all, remember that the nature of chemical union and 
the cause of the valence of the atoms are entirely unknown to us, and 
that, therefore, neither the idea of a variable, nor yet of a constant valence 
constitutes a final explanation. So long as we do not possess an hypothe- 
sis upon the real causes, our task is restricted to collecting the varying 
combination relations of atoms into single points of view. The supposi- 
tion of a constant valence, formerly given preference as the simpler, has 
shown itself to be insufficient. Because it maintains real differences, where OXYGEN COMPOUNDS OF THE METALLOIDS. 
173 
such are not perceptible, it departs from the ground of induction. The 
idea of a variable valence, which takes all facts into consideration, is not 
prejudicial. Indeed, it finds a pregnant analogy in the deportment of the 
hydrocarbon radicals. By the elimination of hydrogen from the saturated 
molecules (C2H6. for example), we obtain radicals or groups of increasing 
valence (C2H5, C2H4, C2H3, C2H2, ....), just as the valence increases 
from fluorine to carbon (FI = 19 ; O = 16 ; N= 14; C= 12). The group 
C2H2, however, is bi- and tetravalent; the group C2H3, mono- and trivalent. 
Everything, however, indicates that the chemical elements do not repre- 
sent final individuals, but are composed of one or several primordial sub- 
stances. 
The principles of chemical structure, presented above, show 
themselves most distinctly, and with the greatest regularity, in 
carbon. The constitution of the innumerable varieties of car- 
bon compounds is explained by the tetravalent nature of the 
carbon atoms, and their ability to combine with each other by 
single affinities. In the other, the so-called inorganic com- 
pounds, the valence and structure relations are more compli- 
cated, and are far less investigated, but even in them so many 
regularities show themselves that the actual material is greatly 
simplified thereby, and is made more comprehensible. The 
doctrine of valence and structure is the first attempt to refer 
the facts underlying the law of multiple proportions to the 
functions of the elementary atoms. As this theory only com- 
prises actual relations, it cannot be negatived, but only further 
developed. 
The foundation of the theory of atomicity and structure was 
laid down by A. Kekule (1857-1859). It constitutes a repre- 
sentation and further development of the type theory of Ger- 
hardt, at the basis of which lay, 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 : 
12Of -f H20 = 2HI04; S03 + H20 = H2S04; 
Iodine 
heptoxide. 
Periodic 
acid. 
Sulphur 
trioxide. 
Sulphuric 
acid. 
PA + 3H20 = 2H,P04 
Phosphorus 
pentoxide. 
Phosphoric 
acid. 
Conversely, these oxides may be obtained by the removal 
of water from the acids; therefore, they are ordinarily termed 174 
INORGANIC CHEMISTRY. 
anhydrides of the corresponding acids ; I207—anhydride of 
periodic acid ; S03 as sulphuric anhydride; P205as phosphoric 
anhydride, etc. 
Salts are formed by the replacement of the hydrogen of the 
acids by metals. The acids are distinguished as ?nonobasic, 
dibasic, tribasic, or monohydric, dihydric, trihydric, etc., de- 
pending upon the number of their hydrogen atoms replaceable 
by metals. 
Like the metalloids, the metals possess different valences; 
the monovalent metals (Na, K, Ag), replace each one hydro- 
gen atom ; those of higher valence replace more. Therefore 
the metals of higher valence can unite several acid resi- 
dues. This explains the import, e.g., of the following chemical 
formulas : 
II 
S04Ca 
Potassium sulphate. 
K2S04 
Potassium sodium sulphate. 
S04KNa 
Calcium sulphate. 
no3k 
(N08)4Cu 
Ill 
(NOs)3Bi 
The salts of sulphuric acid are called sulphates, those of nitric 
acid, nitrates, those of phosphoric acid, phosphates, etc. The 
symbols of the metals are sometimes written at the beginning 
and sometimes at the end of the chemical formulas of the 
salts; in the second case we mean to indicate that the metal 
atoms are in union with the oxygen. 
Potassium nitrate. 
Copper nitrate. 
Bismuth nitrate. 
Excepting fluorine, all the halogens combine with oxygen 
to form anhydrides and acids which are analogously consti- 
tuted, although not all the members are known for each hal- 
ogen. The acids have one atom of hydrogen that can be 
replaced by metals, and hence are monobasic. Chlorine forms 
the following anhydrides and acids: 
OXYGEN COMPOUNDS OF THE HALOGENS. 
Anhydrides. Acids. 
C120 HCIO —Hypochlorous acid. 
0203 HC102—Chlorous acid. 
(C1205) HClOj—Chloric acid. 
(C1207) HC104—Perchloric acid. 
The anhydrides C1.205 and C1.207 (corresponding to the 
compounds I.205 and I,207) are unknown. The compound 
C1204 exists, and must be regarded as a mixed anhydride of 
chlorous and chloric acid. HYPOCHLOROUS OXIDE. 
175 
The following formulas express the chemical structure of 
these compounds: 
I II I 
Cl—O—Cl 
I II 
Cl—O—H. 
Hypochlorous 
anhydride. 
Hypochlorous 
acid. 
Ill III 
oci—O—CIO 
III 
(CIO)—OH. 
Chlorous anhydride. 
V V 
0,C1—O—CIO, 
Chlorous acid. 
(CIO,)—OH. 
Chloric anhydride. 
Chloric acid. 
VII VII 
0,C1—O—CIO, 
VII 
(CIO,)—OH. 
Perchloric anhydride. 
Ill V 
OC1—o—cio2 
Perchloric acid. 
Chlorous-chloric anhydride. 
In the acids we assume the presence of the monovalent group 
OH (hydroxyl or water residue) in which hydrogen can be 
replaced by metals, by the action of metals or bases. The 
group (C102 or CIO,) combined with hydroxyl is called an 
acid residue or radical.. In the anhydrides two acid radicals 
are united by an oxygen atom ; water converts them into 2 
molecules of acid: 
OClx .H CIO—OH 
>o + o<; = 
OCK \H CIO—OH. 
The salts of perchloric acid are called perchlorates; those 
of chloric acid, chlorates ; those of chlorous acid, chlorites ; 
and those of hypochlorous acid, hypochlorites. 
Hypochlorous Oxide—0,0—Hypochlorous anhydride, 
is produced by conducting dry chlorine gas, in the cold, over 
precipitated and dried mercuric oxide : 
HgO + 2C12 = HgCl2 +C120. 
The disengaged gas is condensed in a bent glass tube, cooled 
by a freezing mixture. 
Hypochlorous oxide is a reddish-brown liquid, resembling 
chlorine. It boils at -f-200, and passes into a yellow vapor. 
The vapor density is 43.3 (H = 1), corresponding to the 
formula, C120 = 86.7. This oxide is very unstable, and in the INORGANIC CHEMISTRY. 
course of a few hours decomposes, yielding chlorine and oxy- 
gen. It has strong oxidizing and bleaching properties. It 
dissolves in water to hypochlorous acid : 
Hydrochloric acid decomposes it into water and chlorine: 
C120 + H20 = 2HOCI. 
C120 + 2HC1 = H.p + 2C12. 
Hypochlorous Acid—HCIO—is only known in aqueous 
solution. It is obtained by conducting chlorine into water in 
which there is suspended freshly precipitated mercuric oxide ; 
the solution cannot be distilled. 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: 
Cl, + H20 = 2IIC1 + O 
2C1QH = 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 + Cl2 = NaCl + NaOCl + H20. 
Upon heating their solutions with dilute nitric acid the free 
acid distils over. 
On shaking the aqueous solution of hypochlorous acid with mercury, 
there is produced a white precipitate of HgO, HgCh, soluble in hydro- 
chloric acid (salts of hypochlorous acid form HgO). This behavior serves 
to distinguish hypochlorous acid from chlorine, which under like circum- 
stances 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 KC10S, nitric acid and reducing substances- like 
arsenic trioxide is warmed : 
2HCIO3 — C1203 -f- H20 + 02; 
a yellowish-green gas escapes, which can be condensed at 
—20°. Chlorine trioxide is a reddish-brown liquid, boiling 
about o° and decomposing rapidly. Its vapors, when heated 
to 50°, explode with violence. The oxide is soluble in water, 
forming chlorous acid C102H or CIO.OH, unknown in a free 
condition. With alkalies it yields salts called chlorites. 
Chlorine Tetroxide—Cl204—is the mixed anhydride of CHLORIC ACID. 
177 
chloric and chlorous acid, as water and the alkalies decompose 
it into these two acids: 
CIO }° + H*° = clo-OH + cio2.OH 
It is formed when sulphuric acid acts upon potassium chlo- 
rate (KC103); a dark yellow gas escapes, upon the application 
of heat and at —20° condenses to a reddish-yellow liquid, 
boiling at -f- 90. 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 the oxide 
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 yellow phosphorus into water, 
contained in a measuring glass, then allowing sulphuric acid 
to touch the bottom of the tube, drop by drop, by means of a 
pipette. By the action of the disengaged tetroxide the phos- 
phorus will burn under water with a brilliant light. 
When concentrated sulphuric acid is added 'to a mixture of 
potassium chlorate and sugar, a violent combustion occurs. 
Chlorous acid. 
Chloric acid. 
The vapor density of the oxide is 33 (H = 1); the gaseous molecules, 
therefore, possess the formula C102 (= 67.3). It is very probable, that 
the molecules have a double formula, C1204, at a lower temperature; this is 
confirmed by the manner in which water decomposes the oxide, and by 
its perfect analogy to nitrogen tetroxide, whose N204 molecules have been 
proved to dissociate into N02. 
Chloric Acid—HC103, or C102.0H—is obtained by de- 
composing an aqueous solution of barium chlorate with sul- 
phuric acid : 
(C103)2Ba + S04H2 = BaS04 + 2HCIO3. 
Barium chlorate. 
Barium sulphate. 
The barium sulphate separates as a white insoluble powder, 
and can then be filtered off from the aqueous solution of the 
acid. This is concentrated, under an air-pump, until the 
specific gravity becomes 1.28, and it then contains about 40 
per cent of chloric acid; it is oily and, when heated to 40°, 
decomposes into chlorine, oxygen, and perchloric acid,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 + 5HC1 = 3H20 + 3C12. 178 
INORGANIC CHEMISTRY. 
The chlorates are produced, together with chlorides, by the 
action of chlorine, in the presence of heat, upon many bases 
(compare Potassium chlorate): 
Perchloric Acid—HC104, or C10g.0H. This is the most 
stable of all the oxygen derivatives of chlorine. As previously 
stated, it is produced by the decomposition of chloric acid, but 
is more easily obtained from its salts. Upon heating potassium 
chlorate to fusion, oxygen escapes, and potassium perchlorate 
results: 
6KOH + 3C12 '= 5KCI + KCIO3 -f 3H20. 
2KCIO3 = KC104 + KC1 + Oa. 
Upon warming the perchlorate with four parts sulphuric acid, 
perchloric acid distils over: 
2C104K + H2S04 = K2S04 + 2HC104. 
The pure acid is a mobile, colorless liquid, fuming strongly 
in the air; its specific gravity is 1.78 at 150. It boils at iio°. 
It cannot be preserved, since after a few days it decomposes 
with violent explosion. It also explodes in contact with phos- 
phorus, paper, carbon, and other organic substances. It pro- 
duces painful wounds when brought in contact with the skin. 
It 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 + 2H20:—is a thick, oily liquid, 
resembling sulphuric acid, and boils unchanged at 208°. It 
may also be obtained by evaporating the aqueous solutions of 
perchloric and chloric acids. When the crystalline hydrate is 
distilled it breaks up into anhydrous perchloric acid and the 
second hydrate: 
2C104H.H.20 = 4- C104H.2H20. 
Bromine yields the following oxygen compounds: 
HBrO Hypobromous acid. 
HBrOg Bromic acid. 
HBr04 Perbromic acid. 
The corresponding anhydrides are not known. The acids 
are perfectly analogous to. the corresponding chlorine com- 
pounds. 
Hypobromous Acid—HBrO—is formed when bromine 
water acts upon mercuric oxide; the aqueous solution can be BROMIC ACID IODIC ACID. 
179 
distilled in vacuo, and posssesses all the properties of hypo- 
chlorous acid. 
Bromic Acid—BrOaH. Bromates are formed by the 
action of bromine, in the heat, upon the aqueous solution of 
the alkalies or of barium hydrate; an aqueous solution of the 
acid can be obtained from the barium salt by decomposing the 
latter with sulphuric acid. A more practical method of getting 
the free acid is to let bromine act upon silver bromate or oxi- 
dize bromine with hypochlorous acid : 
The aqueous solution may be concentrated in vacuo until its 
content reaches 50.6 per cent. BrOsH and then closely cor- 
responds to the formula BrOsH -f 7H20. When heated it 
breaks up into bromine, oxygen, and water. 
Perbromic Acid—Br04H—is said to be formed in the 
action of bromine vapor upon perchloric acid: 
5C120 + Br2 + H20 = 2Br03H + ioCl. 
and is perfectly similar to the latter. 
C104H + Br = Br04H + Cl. 
Iodine forms the following anhydrides and acids: 
I205 HI03 — Iodic Acid. 
(I207) HIOt — Periodic Acid. 
Iodic Acid—HI03. Its salts (iodates) are formed in the 
same manner as those of chloric and bromic acids, by dissolv- 
ing iodine in a hot solution of potassium or sodium hydrate: 
The free acid can be obtained by the oxidation of iodine with 
strong nitric acid, or by means of chlorine; further, by the 
action of iodine upon chloric or bromic acids, whereby the 
iodine directly eliminates the chlorine and bromine: 
6KOH + 3la = SKI + IKOa + 3H20. 
2HCIO3 + J2 = 2HIO3 + Cl2. 
Upon evaporating the aqueous solution the free iodic acid crys- 
tallizes in colorless rhombic tablets of specific gravity 4.63. 
The solution possesses strong oxidizing properties. When iodic 
acid is heated to 170° it decomposes into water and iodic an- 
hydride: 
2HIO3 = I20 i + H20. 
It is decomposed, similarly to chloric acid, by hydrochloric 
acid: 180 
INORGANIC CHEMISTRY. 
Reagents, like H2S, S02 and HI, reduce it to iodine. Periodic 
acid sustains similar decompositions. 
Iodic Anhydride—I2Os—is a white crystalline powder, 
which dissolves in water to form iodic acid. It decomposes at 
300° into iodine and oxygen. 
Periodic Acid—HI04—is produced by the action of io- 
dine upon perchloric acid: 
2l03II + ioIICl = I, + SC12 + 6H20. 
2lIC104 + I2 = 2HI04 + Cl2 
Upon the evaporation of the aqueous solution, the acid crys- 
tallizes out with two molecules of water (HI04, 2H20—com- 
pare below). In the air, the crystals deliquesce, fuse at 130°, 
and at a higher temperature decompose into water and pe- 
riodic anhydride, the latter at once breaking up into oxygen 
and iodic anhydride: 
2(HI04 + 2H2o) = i205 + o2 + 5h2o. 
The existence of the hydrates of periodic and perchloric acids, as well 
as of many others (See Sulphuric and Nitric acids), which we .once re- 
garded as molecular compounds (p. 171), is interpreted at present by the 
acceptance of hydroxyl groitps, directly combined with the element of 
higher equivalence: 
VII 
C104H -f- H„0 = C102(0H)3— trihydrate or tryhydric acid. 
VII 
CIO.H -|- 2H,0 = CIO (OH)5— pentahydrate or pentahydric acid. 
VII 
C104H -f- 3H20 = Cl (OH)7— heptahydrate or heptahydric acid. 
The extreme hydrates, Cl(OH)7and 1(011)7, in which all seven affinities 
of the halogen atom are attached to hydroxyl groups, are not known, but 
probably exist in aqueous solution. As they give up water, and one atom 
of O becomes simultaneously united with two bonds to the halogen, they 
yield the lower hydrates—even to the monohydrate C1030H. Perchloric 
acid continues monobasic in the polyhydrates, since but one hydrogen 
atom is replaced by metals : 
CI06H5 + KOH = C10iK 4- 3H20. 
On the other hand, periodic acid (IOsOH) is not only monobasic, but as a 
pentahydrate (IO(OH)5) can, like the polybasic acids, furnish also poly- 
metallic salts, as: 
VII f (OH)3 
10 \ (ONa)2 
VII / (OH), 
IO \ (OAg)3 
VII 
IO(O.Na)5 
VII 
10(OAg)5. 
Salts also exist which are derived from condensed polyiodic acids, as: OXYGEN COMPOUNDS OF THE HALOGENS. 
181 
/(0H)4 
X(OH)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 are, in fact, heptads in their highest combinations. 
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 for hydrogen di- 
minishes with increasing atomic weight from FI to I (see page 
64), the affinity for oxygen is the exact reverse. Fluorine is 
not capable of combining with oxygen ; the chlorine and bro- 
mine compounds are very unstable, and are generally not 
known in free condition ; the iodine derivatives, on the con- 
trary, are the most stable. In accord with this is the fact that 
in the higher oxygen compounds chlorine and bromine are set 
free by iodine, while in the hydrogen and metallic compounds 
of the halogens the direct reverse is the case, viz., that iodine 
and bromine are replaced by chlorine. 
Further, the oxygen compounds exhibit the remarkable pe- 
culiarity that their stability increases with the addition of oxy- 
gen. The lowest acids, HCIO, HBrO, HC102, are very un- 
stable, even in their salts; they possess a very slight acid char- 
acter, and are, too, separated from their salts by carbon di- 
oxide. The most energetic and most stable are the highest 
acids, HC104, HBr03, HIOs, in which 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 
chain-like grouping of the oxygen atoms (according to the 
constant atomicity theory, See p. 171). 
The peculiar behavior of the oxygen compounds of the ha- 
logens, their variable stability and decomposition, and their 
modes of formation, find a clearer explanation in their thermo- 
chemical relations. All oxide compounds of chlorine and 
bromine are endothermic, i.e., heat is rendered latent in their 
production from the elements (compare p. 65). They do not 
result, therefore, by direct union of the elements; further, 
they are not very stable, decompose readily with elimination 
of oxygen, and then oxidize strongly. The heat, appearing 182 
INORGANIC CHEMISTRY. 
in the formation of chlorine monoxide, and of the hypothetical 
pentoxides, C1205 and Br2Os (in their production from the 
elements and solution in water), corresponds to the symbols: 
(C12,0 — gas) = — 18,040; (CI2,05,Aq.) = —20,480. 
(Br2,05,Aq.) = — 43,500. 
In the formation of iodine pentoxide and iodic acid heat is 
liberated: 
(I2,°6) = + 44,86o; (I.O..H) = + 57,880. 
This explains its stability in comparison with the chlorine 
and bromine compounds, and also the direct production of 
iodic acid by the oxidation of iodine. When the pentoxides 
are compared with each other, it is seen that the most heat is 
rendered latent in the formation of bromine pentoxide, Br205 
—the affinity of bromine for oxygen, consequently, is the 
lowest, that of iodine the greatest. This is also evident from 
the heat of formation of the acids, in dilute aqueous solution, 
or of the potassium salts in solid condition: 
(Cl,03,H,Aq.) = 23,940; (Br,03,H,Aq.) =_• 12,420; (I,03,H,Aq.) = 
55,710; (C1,03,K) = 94,600; (Br,Os,K) = 87,600; 
(I,03,K) = 128,400. 
We now understand why chlorine and bromine are separated 
from chloric and bromic acids by iodine, with formation of 
iodic acid, while bromine does not act upon chloric acid. 
Later, in the groups of sulphur and of phosphorus, we will 
observe that the middle members, selenium and arsenic, show 
a less liberation of heat in their oxygen compounds—their 
affinity, therefore, is slighter than that of their analogues. 
OXYGEN COMPOUNDS OF THE ELEMENTS OF 
THE SULPHUR GROUP. 
The elements sulphur, selenium, and tellurium combine with 
two atoms of H, and also yield oxygen acids, which contain 2 
H atoms: 
In these acids i and 2 atoms of H can be replaced by metals; 
hence they are dibasic. By the replacement of 1 atom of H 
we get the so-called acid or primary salts, while the neutral or 
secondary salts are obtained by the replacement of both hy- 
drogen atoms: 
HaS (so2h2) h2so3 h2so4. 
S04KH 
SOtKr 
Acid potassium sulphate. 
Neutral potassium sulphate. OXYGEN COMPOUNDS OF SULPHUR. 
183 
I. OXYGEN COMPOUNDS OF SULPHUR. 
Hyposulphurous acid. 
(so2h.2) 
so, 
S03H, 
Sulphurous anhydride. 
so, 
Sulphurous acid. 
so4h2 
Sulphuric anhydride. 
Sulphuric acid. 
In addition to these compounds there are others of more 
complicated nature. They will be studied later. 
The structure of the former may be expressed by the follow- 
ing formulas: 
IV 
O = S = O 
IV OH 
O = S< 
X)H* 
Sulphur dioxide. 
Sulphurous acid. 
VI 
O = S = O 
& 
vi OH 
O = S< 
roH 
Sulphur trioxide. 
Sulphuric acid. 
Sulphur Dioxide, S02, or sulphurous anhydride, is formed 
by burning sulphur or sulphides in the air: S + 02 = S02. 
I vol. I vol. 
The combustion may also be effected by the action of me- 
tallic oxides (copper oxide, manganese peroxide) which give 
up their oxygen quite readily. It is most conveniently pre- 
pared for laboratories by heating sulphuric acid with mercury 
or copper: 
2H2SQ4 + Cu = CuS04 + S02 + 2HaO. 
The acid is similarly decomposed by heating it with carbon. 
Copper sulphate. 
2S04H2 -f C = 2S02 + C02 + 2H20. 
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 mercury. 
* The structure of sulphurous acid must probably be expressed by the 
IV 
formula, H — SO2 — OH, according to which 1 atom of H is connected 
with sulphur, but the other is contained as hydroxyl. This appears from 
the carbon derivatives of sulphurous acid. Probably both structural cases 
exist in compounds as two isomeric series of neutral ethers of the acid are 
known. 184 
INORGANIC CHEMISTRY. 
It is a colorless gas, with a suffocating odor. Its density is 
31.9 (H = I), corresponding to the molecular formula S02= 
63.9. It condenses at —i5°,or at ordinary temperatures under 
a pressure of two atmospheres, to a colorless liquid, of specific 
gravity 1.45, which crystallizes at — 76° and boils at — io°. 
Upon evaporation the liquid sulphur dioxide absorbs much 
heat; so that if some of the liquid is poured upon mercury, in 
a clay crucible and the evaporation accelerated by blowing air 
upon it, the metal will solidify. Water dissolves 50 volumes 
of sulphur dioxide gas with liberation of heat.* The gas is 
again set free upon application of heat. The solution shows 
all the chemical properties of the free gas. 
Sulphur dioxide has great affinity for oxygen. The gases 
combine when dry; if their mixture be conducted over feebly 
heated platinum sponge* sulphur trioxide results : 
2 vols. I vol. 
2S02 + 02 = 2SO3. 
In aqueous solution the dioxide slowly absorbs O from the 
air, and becomes sulphuric acid: 
so2 + h2o + o = h2so4. 
Aqueous sulphur dioxide is converted more rapidly into sul- 
phuric acid by the action of Cl, Br, and I; 
SOsH2 + H20 + Cl2 = S04H2 + 2HC1. 
Here the decomposition of a molecule of water is effected in 
consequence of the affinity of the halogen for hydrogen and of 
sulphurous acid for oxygen. On adding sulphurous acid to a 
dark-colored iodine solution the latter is decolorized. 
Similarly, sulphurous anhydride and its solution withdraw 
oxygen from many compounds rich jn that element; hence it 
deoxidizes strongly and passes over into sulphuric acid. Thus 
chromic acid is reduced to oxide, and the red solution of per- 
manganic acid is decolorized with formation of manganous 
salts. Many organic coloring substances, like those of 
flowers, are decolorized by it. This property is what leads to 
its application in the bleaching of wools and silks, which are 
strongly attacked by the ordinary chlorine bleaching agents, 
(p. 49.) 
The dioxide may be deoxidized; thus by H2S sulphur is sep- 
arated out by strong reducing agents: 
S02 + 2H2S = 2ll20 + 3S. 
* 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. HYPOSULPHUROUS ACID. 
185 
If, however, both gases are strongly diluted by other neutral 
gases, the action is but very slow. 
A mixture of equal volumes of S02 and CI2 unites in direct sunlight to 
thionyl chloride, S02C12. (P. 193.) When sulphur dioxide acts upon 
warmed phosphoric chloride, the products are phosphorus oxychloride, 
and the compound SOCb: 
Chlorthionyl—SOCl2—may be viewed as sulphur dioxide in which 
one atom of O is replaced by two atoms of chlorine. It is a colorless 
liquid with a sharp odor, and boils at 78°. Water decomposes it into 
hydrogen chloride and sulphurous acid : 
so2 + PC15 = POCl3 +. S0C12. 
SOCl2 + 2H20 = S02 4- H20 + 2HC1. 
Sulphurous Acid—H2S03—is not known in free condi- 
tion, but is probably present in the aqueous solution of. S02. 
On cooling the concentrated solution to o°, colorless cubical 
crystals separate, which have the composition (SO,2 -f- *5H20) 
or (S03H2 + i4H20). If the aqueous solution is allowed to 
stand for some time, especially in sunlight, sulphur separates 
with formation of sulphuric acid: 
3so2 + 2ii2o = 2so4h2 + 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 from its 
salts by stronger acids it decomposes into its anhydride and 
water: 
Na2S03 -f 2HCI = 2NaCl + SO, + H20. 
Pyrosulphuroua Acid, S205H2, bears the same relation to sul- 
phurous acid, as pyrosulphuric acid S2OtH2 to sulphuric acid (p. 191). 
The acid is only known in its potassium salt—S2Q5K2 (see Potassium sul- 
phite). 
Hyposulphurous Acid—H2S02 or S204H2. On adding zinc to the 
aqueous solution of sulphurous acid the metal dissolves without liberation 
of hydrogen. A yellow solution is obtained, which decolorizes indigo and 
litmus solutions energetically. Schiitzenberger has shown that this prop- 
erty is due to the hyposulphurous acid contained in the solution, formed 
there by the action of the H set free by the zinc upon a second molecule 
of SO3H2: 
H2S03 4- Zn = S03Zn + If2, and 
H2SOs + H2 = S02H2 4- HjO. 
The pure aqueous solution is obtained by the decomposition of its salts. 186 
INORGANIC CHEMISTRY. 
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 of a similar behavior 
of the lower oxygen derivatives of chlorine and bromine. 
The salts are more stable than the acid. The sodium salt is obtained 
by the action of zinc filings upon a concentrated solution of primary sodium 
sulphite. Its composition is not established with certainty; it corresponds 
to either the formula S02HNa or S204Na2. The salt solutions abstract 
oxygen very rapidly from the air and change to sulphites. 
Two peculiar oxides of sulphur, which, however, do not afford any cor- 
responding acids and salts, but resemble the peroxides more, are sulphur 
sesquioxide and sulphur heptoxide. 
Sulphur Sesquioxide—S203. Is obtained by the solution of flowers 
of sulphur in anhydrous sulphuric anhydride; it separates out in blue 
drops, which solidify to a mass resembling malachite. It decomposes grad- 
ually, more rapidly on warming, into S02 and sulphur. It is very violently 
broken up by water, with formation of sulphur, S02,S04H2 and polythionic 
acids. It dissolves with a blue color in concentrated sulphuric acid. 
Sulphur Heptoxide—S2Ot—is produced by the action of a silent elec- 
tric discharge of great tension upon a mixture of S02 and oxygen. It sepa- 
rates, in oily drops, which solidify to a crystalline mass at o°. Upon standing 
but especially upon warming, it gradually decomposes into S03 and oxygen: 
s2o7 =2SOs + o. 
It fumes strongly in the air, and with water decomposes into sulphuric 
acid and oxygen: 
Its solution in concentrated sulphuric acid is tolerably stable. It ap- 
pears, also, in the electrolysis of sulphuric acid, and upon the addition of 
H202 to strongly cooled sulphuric acid. 
S2Ot + 2HjO = 2S04H2 +o. 
Sulphur Trioxide—S03—or sulphuric anhydride, is pro- 
duced, as previously described, by the union of S02 and 
oxygen, aided by platinum black; or when S02 and air are 
conducted over ignited oxide of iron (Wohler). It is most 
conveniently obtained by heating fuming (Nordhausen) sul- 
phuric acid (p. 191) ; the escaping white fumes are condensed 
in a chilled receiver. Sulphur trioxide exists in two different 
(polymeric) modifications. In the one form obtained by cool- 
ing the vapors, there is produced a white, asbestos-like mass 
which, after fusion, crystallizes in long, colorless prisms; it 
melts at 160 and boils at about 46°. The vapor density agrees 
with formula SOs. By keeping it below 250 it passes in another 
so-called solid modification, which does not fuse until above 
50°, and passes into the liquid variety. SULPHURIC ACID. 
According to later investigations of Weber neither modification is 
the pure anhydride, but contains water. He obtained the pure anhydride 
by subjecting the asbestos-like variety to repeated and careful distillations 
in a closed tube. It is a readily mobile liquid, of specific gravity 1.940 at 
16°, but solidifies to long, transparent, needles, resembling saltpetre. The 
crystals fuse at 14.8° and boil at 46.2°. By the addition of a small quan- 
tity of moisture the transparent crystals pass into the asbestos-like needles 
of the ordinary anhydride. 
Sulphuric oxide fumes strongly in the air, and attracts mois- 
ture with avidity. When thrown on water it dissolves with 
hissing, to form sulphuric acid (S03 H20 = H2S04). 
When the vapors are led through heated tubes they are 
decomposed into S02 and oxygen. 
This acid has long been known and is extensively applied 
in technology, etc. Besides the reactions already mentioned, 
it arises in the oxidation of sulphur by nitric acid. It was ob- 
tained formerly by heating ferrous sulphate (FeS04); at pres- 
ent, however, it is almost exclusively manufactured in large 
quantities, after the so-called English lead chamber process. 
This method is based upon the conversion of S02 into S04FI2. 
Sulphur or pyrite (FeS„), is roasted in ovens, and the disen- 
gaged S02 immediately conducted, 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 
collects upon the floor of the same. 
SULPHURIC ACID—H2S04. 
The lead chamber process is very complicated, being influenced by the 
quantity of the reacting substances and the temperature, and as yet is not 
fully explained. It is most simply represented as follows : in the presence 
of water, the nitric acid oxidizes the SO2 to sulphuric acid, and the former 
is reduced to nitrogen oxide or nitrogen dioxide: 
The oxygen of the air (which entered the chambers simultaneously with 
the S02) and the steam convert the NO again into nitric acid: 
3S02 -(- 2HNO3 + 2H2Q = 3H2S04 + 2NO. 
2NO + 30 + H20 = 2HNO3, 
and this converts a fresh portion of S02 into sulphuric acid. Or, the nitric 
oxide forms N203and N02 by union with oxygen, and these, in the presence 
of steam, oxidize sulphur dioxide to sulphuric acid: 
S03 + H20 -f N2O3 = S04H2 + 2NO 
and 
so2 + H20 -f no2 = so4h2 + NO. 
The regenerated nitric oxide is again subjected to the same transforma- 
tions. In this manner apparently one and the same quantity of nitric acid, 188 
INORGANIC CHEMISTRY. 
by sufficient air access and water, changes an unlimited amount of S02 
into sulphuric acid ; the nitrogen oxide (and other oxides of nitrogen) acts, 
as it were, as a carrier of oxygen. In fact, in this process, in addition to 
the NO, small quantities of N20 and N are formed from the nitric acid, 
and these not being oxidized by the air, escape, with excess of the latter 
from the chambers. For the continuation of the process, the regular ad- 
dition of a definite amount of nitric acid is acquired. 
In practice, the active nitrogen oxides (N203 and N02) are carried along 
and withdrawn from the action by means of the escaping nitrogen and excess 
of air. To avoid any further loss of nitric acid by this means, the escaping 
Fir,. 84. 
brown gases are conducted through the so-called Gay-Lussac tower. This is 
constructed from lead sheets, and filled with pieces of coke, over which con- 
centrated sulphuric acid constantly trickles. The acid completely absorbs 
the nitrogen oxides N203 and N02, with formation of nitrosylsulphuric 
acid (see p. 2051. The nitrogen oxides can be regained from the acid— 
the so-called nitroso-acids—collected at the bottom of the tower, and made 
useful in the production of sulphuric acid in the chambers. This is effected 
at present, in the so-called Glover tower, which is constructed of lead 
plates and fire-proof bricks, and inserted between the sulphur ovens and 
lead chambers. In this the nitroso-acid (diluted with the previously ob- 
tained chamber acid) is allowed to run over fire-brick, while the hot gases 
of combustion from the sulphur ovens stream against it. This cools the 
hot gases to the required •temperature (70-80°), water evaporates from the 
chamber acid, and, at the same .time, the nitrogen oxides are set free (see 
p. 207), and carried into the lead chambers. Hence, the Glover tower SULPHURIC ACID. 
189 
serves, not only for complete utilization of the nitrogen oxides, but also 
for the concentration of the chamber acid. 
The chamber process may be illustrated by the following laboratory ex- 
periment: A large, glass flask (Fig. 84) A replaces the lead chamber; in 
its neck are introduced, by means of a cork, several glass tubes, which 
serve to introduce the various gases. In «,S02 is developed by heating a 
a mixture of H2S04 and Hg or copper strips. The flask b contains some 
dilute nitric acid and copper turnings, from which NO is evolved. 
Water is boiled in c to afford steam. Air enters through d while the ex- 
cess 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 
(P- i87)- 
The regenerated nitric oxide yields NOa with the oxygen of the air, and 
converts another portion of SO,2 into sulphuric acid. In time aqueous sul- 
phuric acid collects upon the bottom of the vessel. If, at first, only S02, 
NO and air enter without the steam, we get (by aid of the moisture of the 
air) the compound S02 j (the so-called nitrosulphonic acid) which 
covers the walls of the vessel with a white crystalline sublimate (comp. p. 
206). These crystals, known as lead-chamber crystals, are also formed 
in the technical manufacture of sulphuric acid, when an insufficient quan- 
tity of steam is conducted into the chambers. Water decomposes them 
into sulphuric acid and nitrogen oxides. 
The acid collecting in the chambers (chamber acid) pos- 
sesses, when the operation has been properly conducted, the 
specific gravity 1.5 (50° according to Beaume) ; it contains 
about 60 per cent. H2S04 and 40 percent. H20. For concen- 
tration the chamber acid is first heated in open pans until the 
specific gravity reaches 1.72 (6o° Beaume). The lead vats are 
strongly attacked by further evaporation, hence the acid is 
finally heated in glass vessels, or, better, platinum retorts, un- 
til the residual liquid has acquired the specific gravity, 1.83 
(65.5° Beaume). It is now entered upon trade under the 
name crude sulphuric acid (Acidum sulphuricum cruduni). It 
still contains about 8 per cent, water and traces of lead and 
arsenic. By further concentration we can obtain 95-96 per 
cent. H2S04 (extra concentrated acid). 
By the distillation of the crude English acid an aqueous so- 
lution at first distils over (*/$ distillate), but at 330° we obtain 
almost pure H2S04 {Acidum sulphuricutn or destillatum). 
This has the specific gravity 1.854 at o° or 1.842 at 120, and 
contains about 1.5 per cent, water. On cooling this to—350 
white crystals separate, which, after repeated recrystallization 
fuse at + 10.50 ; this is the anhydrous acid, H2S04. When 
this is heated, white fumes of S03 escape at 40° ; the liquid 
begins to boil at 290°, and at 330° the acid, with 1.5 per 
cent. H20, again distils over. 190 
INORGANIC CHEMISTRY. 
From these data it is obvious that sulphuric acid, even at a gentle heat, 
sustains a partial decomposition (dissociation) into S03 and H20,which 
again unite to form sulphuric acid when they cool. The dissociation is com- 
plete 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 49 (9/); the empirical 
formula, found to be half as large, is explained by the decomposition of 
the molecule of H2S04 into the molecules SOs and II20, 
so4h2 = SOs 4t h2o. 
Concentrated sulphuric acid is a thick, oily liquid. On 
cooling a sulphuric acid containing about 15 percent, water to 
o°, large six-sided prisms of the hydrate S04H2 + H20 sepa- 
rate ; these fuse at T- 8.5°, and give up water at 205°. The 
second hydrate, S04H2 + 2H20, corresponding to the maximum 
contraction, has the specific gravity 1.63, and yields water at 
1950. The concentrated acid possesses an extremely great 
affinity for water, and absorbs aqueous vapor energetically, 
hence is applied in the drying of gases and in dessicators. It 
unites with water with the evolution of considerable heat, and, 
for this reason, it is practically recommended, in mixing the 
acid, to pour the latter in a thin stream into the water, and not 
the reverse, as otherwise explosive phenomena occur. In 
mixing sulphuric acid with water, a contraction of the mixture 
takes place ; its maximum corresponds to the hydrate S04H2 
+ 2H20. 
1 volume. 
1 volume. 
i volume. 
The existence of the hydrates of sulphuric acid is explained, as in the 
case of periodic acid, by the supposition of hydroxyl groups : 
S04H„ -4- 2H..O = S(OH). Hexahydroxyl sulphuric acid. 
SO;H; + H20 = S(OH)J Tetra “ 
SQ4H2 = S02(0H)2 Normal sulphuric acid. 
The tetra- as well as the hexahydroxylsulphuric acid yield only salts of 
the normal dibasic acid, when they are acted upon by bases. Salts, in 
which several H atoms are replaced by metals, 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. The with- 
drawal of H and O from such substances leaves the carbon. 
This explains the charring action of H2S04 upon wood, sugar, 
and paper. When sulphuric acid acts upon alcohol (C2H60), 
ethylene, C2H4 (p. 150), results. 
By conducting H2S04 over red hot porous bodies, it is de- 
composed into sulphur dioxide, water, and oxygen : PYROSULPHURIC OR DISULPHURIC ACID. 
I9I 
H2S04 = SO, + HaO + o. 
This decomposition affords us a method for manufacturing 
oxygen technically ; the sulphur dioxide is absorbed by water 
and afterwards converted into H2S04. When heated with S, 
P, C, and some metals (Hg, Cu), the acid is reduced to di- 
oxide (p. 183). Nearly all the metals are dissolved by it, 
forming salts ; only lead, platinum, and a few others are 
scarcely attacked at all. It is a very strong acid, and, when 
heated, expels most other acids from their salts; upon this 
depends its application in the manufacture of hydrochloric 
and nitric acids. The barium salt (BaSQ4) is characterized 
by its insolubility in water and acids; therefore, sulphuric 
acid added to solutions of barium compounds produces a white 
pulverulent precipitate, which serves to detect small quantities 
of the acid. 
Pyrosulphuric, or Disulphuric Acid—H2S2Ot.—On 
withdrawing one molecule of water from two molecules of the 
acid there results the compound S207H2, whose formation and 
structure may be represented by the following formula: 
so,<°S _ so,/°H 
/on-".0 - /" 
5U«\OH su!\OH 
As this contains two hydroxyl groups it is a dibasic acid; 
yet its manner of formation shows that it possesses an anhy- 
dride character. Later, we will observe that almost all poly- 
basic acids, like phosphoric acid, PO(OH)3, silicic acid, SiO 
(OH)2, and chromic acid, Cr02 (OH)2, are capable, by the con- 
densation and the elimination of several molecules of water, 
of forming like derivatives, which bear the name Poly- or Pyro- 
acids. 
The disulphuric acid is contained in the so-called fuming ox 
Nordhausen sulphuric acid (Acidum sulphuricum fumans), 
which is obtained by heating dehydrated ferrous sulphate— 
green vitriol (FeS04). It is a thick, oily, strongly-fuming 
liquid, of specific gravity 1.85-1.9. When it is cooled, large 
colorless crystals of H2S207 separate; these fuse at 350. Heat 
breaks it up into sulphuric acid and sulphur trioxide, which 
volatilizes: 
Conversely, disulphuric acid may be obtained by dissolving 
S03 in sulphuric acid. The production of fuming sulphuric 
acid also depends on this, as it may be regarded as a solution 
of S03 (or S207H2) in excess of sulphuric acid. 
S207H, = S04H2 + S03. 192 
INORGANIC CHEMISTRY. 
Technically, fuming sulphuric acid is obtained from pyrites (FeS?)— 
(at present only in Bohemia). The decomposition of the pyrites in the 
air affords ferrous sulphate and ferric oxide. The first can be dissolved 
out with water. The solution is evaporated, and the residue roasted in a 
reverberatory furnace, whereby the ferrous salt is changed to ferric 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)S = FeA + JSO3 
Ft‘2°3 + S03 + S04H2- 
The residue, consisting of red ferric oxide, finds application as colco- 
thar (caput mortuum) in polishing and as a paint. 
Of late, solid, fuming sulphuric acid, S207H2, has been prepared accord- 
ing to Winkler’s method. The mixture of S02 + O, obtained by heating 
English sulphuric acid (p. 190) and absorbing the steam produced at the 
same time by sulphuric acid in a coke tower, is conducted over ignited 
platinized asbestos (p. 184) and the resulting S03 collected in concentrated 
sulphuric acid. 
Disulphuric acid yields sulphuric acid with water. Its salts 
may be obtained by heating the primary salts of the latter 
acid : 
Primary pot. 
sulphate. 
2S04HK = h2o + K2§2Ot. 
Pot. disulphate, 
By the application of more heat, the disulphates decompose 
into SOs, and sulphates: = S04K2 + S03; hence these 
reactions may serve for the formation of S03. 
Sulphuric Acid Chlor anhydrides.—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 
S°2 {Cl + 2H*° = s + 2 HC1- 
\)H 
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 
/Cl 
first, SO/ — Sulphuryl Hydroxy-chloride or Chlorsulphonic 
Acid—results when 1 molecule of PC15 acts upon I molcule of H2S04:— 
Cl 
s°2 { QH + PCI5 = S02 + POCl3 + HC1. 
\>H PYROSULPHURYL CHLORIDE. 
193 
The resulting POCl3 acts upon two additional molecules S04H2, with 
formation of metaphosphoric, hydrochloric, and chlorsulphonic acids. 
It is formed, too, by the direct union of S03 with HC1. The most prac- 
tical method for its formation consists in conducting chlorine gas through 
S04H2 (15 parts), and gradually adding PC13(7 parts). Or, HC1 gas is led 
into solid fuming sulphuric acid (S207II2), as long as absorption occurs, 
and then it is distilled (Otto). 
Chlorsulphonic acid is a colorless, strongly fuming liquid of specific 
gravity 1.776 at 180, and bails at 1550. Its vapors possess the normal 
density at a temperature a little beyond the boiling point, but at 180° 
sustain dissociation, which is complete at 440°, and corresponds to the 
equation: 
2S03C1H = S03 -f HaO -)- SOa + Cl*. 
(Cl 
The salt S02 < results from the union of SC3 with KC1. 
The second chloranhydride, S02C12, or Sulphuryl-chloride'* forms when 
PC15 acts upon SOa; by heating S03HC1 to i8o° : 2 S03HC1 = S02C12 -)- 
S04II2; and also by the direct union of S02 with Cl2 in sunlight, or when 
these gases are conducted into camphor. A colorless, suffocating, strongly 
fuming liquid, of specific gravity 1.708 at o°, results. It boils at 70°. 
Water decomposes it energetically into.sulphuric and hydrochloric acids 
(also the case with chlorsulphonic acid). A little water will first change 
it to chlorsulphonic acid; 
S°2\C1 + H2° = S°2 \0H + HC1- 
Its vapor density is normal at 184°; at higher temperatures it gradually 
diminishes, and at 440° corresponds to the decomposition equation: 
so2ci2^so2 + ci2. 
Pyrosulphuryi Chloride, S2()5CI2, is the chloranhydride of pyrosul- 
phuric acid. It is obtained by several reactions, chiefly, however, by the 
action of PC15 or P205 upon chlorsulphonic acid : 
/cl 
2S°/oH H*° = S 2/° 
s°,\Ci 
It is a thick liquid, fuming in the air; has a sp. gr. = 1.858 at o°, and, 
when perfectly pure, boils at 153°. At 183-210° it shows a normal den- 
sity ; it is dissociated at higher temperatures, and this is complete at 440°, 
corresponding to the equation; 
s,o5ci2 = so3 + so2 + Cl2. 
It dissolves gradually in water, without hissing, and decomposes into 
S04H2 and HC1; it at first yields chlorsulphonic acid with a little water. 
Thionyl chloride—SOCl2—(p. 185), may be regarded as a chloranhy- 
dride of sulphurous acid. 
* The group SO2 combined with 2OH in 1I2S04, is known as Sulphuryl. 194 
INORGANIC CHEMISTRY. 
Amid Derivatives of Sulphuric Acid. When NH3 acts on S03, the 
compound S032NH3 arises, which is to be regarded as ammonium 
sulphaminate —SO,/ q \2{ j B *s a white powder, which may be crys- 
allized from water; its solution is not precipitated by barium salts. 
When ammonia acts upon S02 q the ammonium salt of disulphimid 
/SO O H 
acid, HN/s()2’oh, forms; this is tribasic, as all three H-atoms can be 
replaced by metals. 
By this name (from 6eh>v, sulphur) is understood the complex 
acids of sulphur, containing two or more atoms of the latter. 
The following are known : 
POLYTHIONIC ACIDS. 
S2OsH2 — Thiosulphuric acid. 
S206H2 — Dithionic acid. 
S306H2 — Tri “ “ 
S406H2 — Tetra “ “ 
S506H2 — Penta “ “ 
The general chemical character of these acids is represented 
simply and distinctly in the following structural formulas. We 
suppose that they contain one or two univalent groups, S03H, 
VI 
or —S02 OH, in which one affinity of sulphur is unsaturated. 
This group is known as the sulpho group; it is also present in 
organic sulpho-acids, and corresponds to the acid-forming 
carbon group, COOH, called carboxyl. From this group 
(written in another form) are derived the above-observed 
acids: 
H.SOa.OH 
ho.so2.oh 
0/so2.oh 
u\so2.oh 
The following structural formulas express the polythionic 
acids: 
Sulphurous acid. 
Sulphuric acid. 
Disulphuric acid. 
S/“,H 
Thiosulphuric acid. 
HS.SO.H 
<J°$ 
Trithionic acid. 
o /S03 H 
a2\SOs H 
Dithionic acid. 
q /SOsH 
°»\SOsH 
Tetrathionic acid. 
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 univalent sulpho groups. In thio- 
sulphuric acid, only iH-atom is replaced by sulpho ; the dithi- 
onic acid, on the other hand, results by the direct union of 
two sulpho groups, with their free affinities. THIOSULPHURIC ACID —TETRATHIONIC ACID. 
195 
/ SH 
Thiosulphuric Acid, H2S203=S2 generally known 
as hyposulphurous acid, can be considered as sulphuric acid in 
which the oxygen of an hydroxyl group is replaced by sulphur. 
It is not known in a free condition, since as soon as it is liber- 
ated from its salts by stronger acids, it at once decomposes 
into S02, S and H20 : 
Its salts, called hyposulphites, are of practical importance (com- 
pare sodium hyposulphite). They are formed by the direct 
addition of sulphur to sulphites: 
S2OsNa2 + 2HC1 = 2NaCl + S02 + S + H20. 
NaaS03 + S = Na2S203; 
similar to the formation of sulphates by the addition of O to 
the sulphites. 
A very interesting formation of thiosulphuric acid is that of 
the action of iodine upon a mixture of sodium sulphite and 
sodium sulphide: 
f NaSO,2.ONa S02.ONa 
\ +I.= I +2NaI. 
(NaSNa SNa 
Sodium hyposulphite. 
Conversely sodium hyposulphite is split up by sodium amalgam 
into S03Na2 and Na2S. 
Dithionic Acid—II2S206—'sonly known in aqueous solution. When 
concentrated in vacuo or when heated it decomposes into sulphuric acid 
and sulphur dioxide. Its manganese salt results from the action of sul- 
phur dioxide upon Mn02 suspended in water: 
Mn02 + 2S02 = MnS206. 
Barium hydrate converts this into the barium salt, from which the free 
dithionic acid is obtained by means of sulphuric acid. 
Tritbionic Acid—H2S306—is not known in a free condition. Its salts 
are produced when an aqueous solution of primary potassium sulphate is 
digested with flowers of sulphur: 
6HKS0, + 2S = 2KaS306 + KAO, + 3H20. 
Separated from its salts by other acids it decomposes into H2S04, S02 and S. 
Its production by the action of iodine upon a mixture of sodium sulphite 
and hyposulphite is especially interesting: 
SCVONa 
NaS02.0Na , j — j 
+ 2 ~ + 2iNal- 
S02.0Na 
Tetrathionic Acid—H2S4Og. Its salts are produced when iodine acts 
upon solutions of the hyposulphites : INORGANIC CHEMISTRY. 
KS.SO..K S.SO.K 
+ I, = I + 2KI. 
KS.S03K S.SOjK 
Potassium tetrathionate. 
The free acid decomposes, upon concentration, into H2S04, S02 and 2S. 
Pentathionic Acid—H2S5Oe—results when H2S is conducted into an 
aqueous solution of sulphur dioxide: 
5so2 + 5h2s = h2s5o6 + 4h20 + 5S 
further, by the action of S2C12 upon barium hyposulphite: 
s.so, s.so. 
/ \ / \ 
S,CL + Ba Ba = S Ba + BaCl, 4- S. 
\ / \ V 
s.so3 S.S03 
Recent investigations indicate the existence of pentathionic acid as 
doubtful. . 
The polythionic acids are distinguished from sulphuric acid 
by the solubility of their barium salts. 
2. OXYGEN DERIVATIVES OF SELENIUM AND TELLURIUM. 
Selenium dioxide. 
Se02 
Se03H2 
Selenious acid. 
(Se03) 
Se04H2 
Selenium Dioxide—Se02, or selenious anhydride, is 
produced when selenium burns in the air or in oxygen. It 
consists of long white needles, which sublime at about 320° 
without fusing. It dissolves readily in water, forming seleni- 
ous acid, H2Se03. The latter is also obtained by dissolving 
the metal in concentrated nitric acid. When the solution 
is evaporated it crystallizes in large, colorless prisms, which 
decompose, on heating, into the anhydride and water. Sul- 
phurous oxide reduces selenious acid, with separation of free 
selenium: 
Selenium trioxide. 
Selenic acid. 
Selenic Acid—H2Se04—is obtained by conducting chlo- 
rine gas into an aqueous solution of selenious acid. 
H,ScOs + 2S02 + H20 = 2H2S04 + Se. 
The solution may he concentrated until it attains a specific gravity of 2.6 
when it becomes an oily liquid, similar to sulphuric acid, and contains 95 
per cent. HaSeO*. If the solution be heated above 280°, the acid breaks 
up into Se02, O and H20. Selenic anhydride is unknown. 
H2Se03 + H20 + Cl2 = H2Se04 -f 2HCI. TELLUROUS ACID TELLURIC ACID. 
197 
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 when tellurium is burned, and forms a white crys- 
talline mass, fusing at a red heat and subliming in water. It is almost 
insoluble in water. 
Tellurous Acid—H2TeOs—is produced when the metal is dissolved 
in concentrated nitric acid. Water will precipitate it from such a solu- 
tion in the form of a white amorphous powder. On warming it readily 
decomposes into Te02 and water. 
Telluric Acid— H2Te04. Potassium tellurate is produced when tel- 
lurium or its dioxide is fused with saltpetre. The acid (telluric) is ob- 
tained from this salt by means of sulphuric acid, and crystallizes from its 
aqueous solution, in large colorless prisms, with 2 molecules of H20 
(H2TeOt -(- 2H20), which are expelled at ioo° C., and the acid remains 
in the form of a white powder. The latter is not very soluble in water, 
and manifests but a slight acid reaction. When carefully heated, the acid 
breaks up into water, and the trioxide Te03, which is a yellow mass insol- 
uble in H20, and by further application of heat decomposes into Te02 and 
oxygen. 
The affinity of the elements of the oxygen group for the 
halogens seems to increase with rise of atomic weight from 
oxygen to tellurium. This is just the reverse of what we 
observed for the halogens with hydrogen; OCl2 is very un- 
stable and is formed with heat absorption, SC12 and SC14 only 
exist at lower temperatures, while SeCl4, TeCl2 and TeCl4 even 
exist as gases. The thermal relations in the formation of the 
oxygen compounds indicate, on the contrary, that in the latter 
the affinity of selenium is the least, as is true of bromine in 
the halogen group (p. 182). 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.) — 78,770 (Se,02,Aq.) = 56,790 (Te,02,Aq.) = 81,190 
(S,03,Aq.) = 142,400 (Se,03,Aq.) = 77,240 (Te,03,Aq.) = 107,040 
The same is seen in the formation of the anhydrous di- 
oxides : 
In all these compounds, consequently, the affinity of selenium 
to oxygen is the least, and this explains the reduction of sele- 
nious by sulphurous acid, as well as the slight stability of 
selenic acid. 
(S,02—gas) — 71,070 (Se,02—solid) = 57,700. 
In the union of sulphur with oxygen ter form sulphur tri- 
oxide (S,03—gas) 103,260 C. are disengaged, while the heat 
of formation of the dioxide equals only 71,070 C. Yet, in 
the combustion of sulphur we get S02 almost exclusively, and 198 
INORGANIC CHEMISTRY. 
only a slight quantity of S03. This is explained by the SO;4 
decomposing readily at high temperatures into S02 and oxygen 
(p. 187), whereas S02 does not sustain decomposition. We 
must remember also that a greater number of atoms or mole- 
cules work together in the formation of complex compounds, 
and, therefore, at least, in case of gases, the simpler deriva- 
tives can be produced more easily. We can, therefore, com- 
prehend that reactions, like + 302 — 2 S03, N2 -f- 3H2 — 
2NH3 or C2 -j- 4H2 = 2CH4, do not occur ordinarily. Then, 
too, it is necessary to consider that, in all bodies exhibiting a 
high heat of formation and decomposable by heat, the heat of 
formation must be simultaneously removed or conducted away 
so that its production can really occur. Many phenomena, 
apparently contradicting the principles of thermo chemistry, 
are thus explainable, for example, that the reactions in which 
but little heat is disengaged very often take place more 
easily than those with greater heat disengagement. It may 
be the so-called catalytic action of platjnum and other metals, 
e. g., in the reaction SO,2 + O = S03, are due to such a re- 
moval of heat. Another instance is the greater reactivity of 
many bodies if they constitute a galvanic chain ; in this in- 
stance the chemical energy is conducted away in the form of 
electricity; as the union of hydrogen and oxygen at ordinary 
temperatures by the formation of a polarization current. 
OXYGEN DERIVATIVES OF THE ELEMENTS OF 
THE NITROGEN GROUP. 
The halogens combine with one atom of hydrogen and also 
afford oxygen acids containing one atom of H. The elements 
of the sulphur group contain two atoms of H in the hydrogen 
derivatives and oxygen acids. In accord with this we find 
that the elements of the N group combine with 3 atoms of H, 
and form acids which also contain three atoms of the same 
element. 
HC1 
hcio4 
H2S 
h2so4 
PHs 
po4h, 
Perchloric acid. 
HC103 
Sulphuric acid. 
H2SOs 
Phosphoric acid. 
PO3H3 
Chloric acid. 
Sulphurous acid. 
Phosphorous acid. 
The acids containing three atoms of H, designated normal 
or Ortho Acids (as HsP04, 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: NITRIC ACID. 
199 
Orthophosphoric acid. 
h:!po4 
Metaphosphoric acid. 
hpo3 
Orthoarsenious acid. 
H3As03 
Metaarsenious acid. 
HAs02 
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): 
(HsN04) 
hno3 
Orthonitric acid. 
(H3N03) 
Ord. Nitric acid. 
UNO, 
The further exit of water affords the true anhydrides (p. 173). 
Orthonitrous acid. 
Ord. Nitrous acid. 
x. OXYGEN DERIVATIVES OF NITROGEN. 
Nitrogen pentoxide. 
N2o5 
NO3H 
Nitric acid. 
N2Os 
no.2h 
Nitrogen trioxide. 
n20 
Nitrous acid. 
(NOH)2 
In addition to the above compounds we have: Nitrogen 
tetroxide (N204), the mixed anhydride of nitrous and nitric 
acids, and two oxides, nitrogen dioxide (N02) and nitric oxide 
(NO), which do not yield acids. 
The following formulas express the structure of these com- 
pounds: 
Hyponitrous oxide. 
Hyponitrous acid. 
N = N 
Nitrogen. 
in in 
N = N 
o 
hi hi 
ON — O — NO 
Nitrogen trioxide. 
V V 
O..N — O — N02 
Nitrogen pentoxide. 
Hyponitrous 
oxide. 
V 
02N — OH 
Ill 
ON — OH 
Ill 
(ON - H)2 
Ill V 
ON — O — NQ2 
Nitric acid. 
Nitrous acid. 
Hyponitrous 
acid. 
Nitrous-Nitric 
anhydride. 
The salts of nitric acid are called nitrates; those of nitrous, 
nitrites. 
This acid occurs in nature only in the form of salts,—potas- 
sium, sodium, and calcium nitrates (compare these)—which 
NITRIC ACID.—UNO.,. 200 
INORGANIC CHEMISTRY. 
have resulted from the decay of nitrogenous organic substances 
in the presence of strong bases (the alkalies). It is sometimes 
present in the air as ammonium salt. The free acid is formed 
in very slight quantity by conducting the electric sparks 
through moist air. 
To prepare nitric acid heat potassium or sodium nitrate with 
sulphuric acid, when the nitric acid will distil over and sodium 
sulphate remain: 
2NaNOs + H,S04 = Na2SO* + 2HN03 and 
NaNOg + H2S04 = HNaS04 + HNOa. 
This process may be conducted in the distillation apparatus 
figured on page 51. The quantity, by weight, of sodium nitrate 
and sulphuric acid corresponding to the second equation must 
be employed, because, with less acid a higher temperature is 
requisite to complete the reaction, and, in consequence, the 
nitric acid that is produced will be partially decomposed. 
Pure anhydrous nitric acid is a colorless liquid of specific 
gravity 1.54 at o°, fumes in the air, and at —40°, solidifies to a 
crystalline mass. At ordinary temperatures it undergoes a par- 
tial decomposition (like H2S04) into water, oxygen, and nitro- 
gen dioxide N02, which dissolves in the acid, with a yellowish- 
brown color; the colorless acid therefore becomes colored 
upon standing, and in sunlight soon turns yellow. At 86° the 
acid commences boiling and sustains a partial decomposition; 
the first portions are colored yellow by the dissolved nitrogen 
dioxide, but subsequently, some aqueous acid distils over. 
Nitric acid is completely decomposed into nitrogen dioxide, 
oxygen, and water, when its vapors are conducted through red- 
hot tubes: 
The acid mixes in all proportions with water. Upon dis- 
tilling the dilute aqueous solution, only pure water passes over 
at first; the boiling temperature gradually rises, and at 1210 a 
solution goes over, which contains 68 per cent. HN03, and has 
a specific gravity of 1.414 at 150. This is the ordinary con- 
centrated nitric acid of trade. When this is distilled with 5 
parts sulphuric acid, an almost anhydrous acid is obtained, 
which may be freed of N02 contained in it by conducting a 
stream of air through it. 
2HN03 = 2NO2 + h2o + o. 
Generally, the anhydrides of acids distil at temperatures lower than the 
acids themselves (S03 is more volatile than H2S04). The higher 
boiling-point of the aqueous nitric acid, in relation to the anhydrous, is 
probably explained by the fact that the hydrate HNO3 -|- HjO, i. <?., the NITRIC ACID, 
201 
normal nitric acid, N04H3 = NO(OH)s (p. 199), is present in this solu- 
tion. The liquid boiling at 1210, 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 
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. It acts as a very 
strong oxidizing agent, destroys organic coloring substances 
and decolorizes a solution of indigo very readily. In so doing 
the nitric acid itself is deoxidized to the lower oxidation pro- 
ducts of nitrogen (NO and N03). Some substances even reduce 
the acid to ammonia. Thus, for example,-if zinc be brought 
into dilute nitric acid (5-6 per cent.) the metal will be dis- 
solved without the liberation of hydrogen. The latter, in 
statu nascendi, acts at once upon the excess of acid and reduces 
it to ammonia, which forms an ammonium salt with the acid ; 
hence, in solution, we have ammonium nitrate in addition to 
the zinc nitrate: 
2HNO3 4- Zn = Zn(N03)2 -f H2 and 
2HN03 + 4H2 = NOgNH4 + 3II2O. 
If the aqueous nitric acid be more dilute (containing more 
than 10 per cent. NOsH) it will be reduced by zinc and other 
metals, not to ammonia, but to the nitrogen oxides, N20, NO, 
N203, and N304. The more concentrated the acid, the higher 
will the oxides be. 
The reduction of nitric acid to ammonia by nascent hydro- 
gen occurs more easily in alkaline solution. If an alkaline 
solution of nitrates be treated with zinc or aluminium filings, 
all the N of the nitric acid will be converted into ammonia : 
Hydroxylamine (p. 128) and ammonia are produced when 
nitric acid acts on tin. 
Nitric acid—N02OH—and its hydrates, NO(OH)g and 
N(OH)5 (see above), usually form salts of the form NOgMe, 
with 1 eq. of the metals; these are called nitrates, and are all 
soluble in water. 
HNOj + 4H2 = NH, + 3H2O. 
Red Fuming Nitric Acid (Acidum Nitricum Fumans) is 
the name given a nitric acid containing much nitrogen diox- 
ide in solution. It is obtained by the distillation of 2 mole- 
cules of HN03 with 1 molecule of sulphuric acid (p. 200), 
or better, by the distillation of commercial nitric acid with 202 
INORGANIC CHEMISTRY. 
much sulphuric acid. It generally has the specific gravity 
i.5-1.54, and possesses greater oxidizing power than the color- 
less 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 is due to the presence of free chlorine and the 
two chlorine derivatives (N02C1 and NOC1), which may be 
considered the chloranhydrides of nitric and nitrous acids. 
Nitroxyl Chloride—N02C1—the chloranhydride of nitric acid, results 
by the union of NO2 with chlorine, and according to the ordinary method 
of forming chloranhydrides (see p. 192), by the action of PC15 or POCl3 
upon nitric acid, or better, its silver salt: 
3N020Ag + P0C13 = PO(OAg)s + 3NOsCI. 
It is a yellowish liquid, boiling at -f- 50. With water it breaks up into 
nitric and hydrochloric acids. 
Nitrosyl Chloride—NOC1—is produced when hydrochloric acid acts 
upon nitric acid, and by the union of NO (2 volumes) with chlorine (1 
volume). It is a reddish-yellow gas which below o° condenses to a liquid. 
It forms nitrous and hydrochloric acids with water: 
Silver nitrate. 
Silver phosphate. 
N0C1 + H*0 = HN02 + HC1. 
It may, therefore, be regarded as the chloranhydride of nitrous acid— 
NO.OH. 
Nitrogen Pentoxide—N205—nitric anhydride, is pro- 
duced when phosphoric anhydride acts on nitric acid : 
2HN03 + P2Os = N205 + 2HPO3; 
further, by conducting nitroxyl chloride over silver nitrate: 
AgO.NOj + NOjCl = S8j}° + AgCL 
It forms colorless, rhombic prisms, fusing at 30° and boiling 
with partial decomposition at 470. It is very unstable, de- 
composing readily with N204 and O, and sometimes exploding 
spontaneously. It yields HN03 with water and evolves much 
heat by the union : 
NO,'}0 + H»° = 2 NO,OH. 
Nitrogen Trioxide, N203, nitrous anhydride, is formed by 
the direct union of nitrogen oxide (4 volumes) with oxygen (1 
volume) at —180 : NITROUS ACID. 
203 
4 NO + O, = 2 N2Os; 
4 vols. 
I vol. 
2 vols. 
by mixing liquid nitrogen tetroxide, N204, with a little water: 
2No}° + H*° = no}0 + 2 NO,OH; 
by the introduction of nitrogen oxide into liquid nitrogen 
tetroxide: 
N204 + 2 NO = 2 N2Os ; 
and by conducting nitric oxide into anhydrous nitric acid:— 
2 N03H , 4 NO = 3 N203 + 11,0. 
It is a dark-blue liquid, boiling at o° with partial decompo- 
sition into NO and N02; both gases combine again to N203 
on cooling. The trioxide mixes with a little cold water, 
forming, probably, nitrous acid (HNG2) ; more water, aided 
by heat, decomposes it into nitric acid and nitric oxide gas : 
Nitrous Acid, HN02, is not known in a free state. Its 
salts are obtained by igniting the nitrates : 
3N02H .= HNOj -f- 2NO -f H20. 
kno3 = kno2 + o. 
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 NO,2 and NO. It may be that 
the nitric acid, at first liberated, is broken up into water and 
the trioxide, which, as we have seen above, decomposes very 
readily into N02 and NO. Similar reddish-brown vapors are 
obtained if nitric acid of specific gravity 1.3-1.35 be per- 
mitted to act upon starch or arsenious oxide (As203). If these 
vapors are cooled, they condense to a liquid, which is green 
at medium temperatures, and consists, probably, of N204 and 
N203. When the green liquid is warmed, vapors escape, which 
condense, on cooling, to a blue liquid, consisting principally 
of N203. 
The nitrous acid that has 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 eg., the acid* 
ified red solution of potassium permanganate is decolorized by 
the addition of nitrites. 204 
INORGANIC CHEMISTRY. 
In very dilute .solution, the action proceeds according to 
the following equation : 
5N02H 4- 2Mn04K + 3S04H2 = 5NO3II + S04K2 -f- 
2 S04Mn -j- 3HjO. 
Nitrogen Tetroxide—N204 or nitrogen dioxide NO, (for- 
merly called hyponitric acid), really constitutes two compounds. 
The former only exists at low temperatures; when heated, it 
suffers a gradual decomposition into the simpler molecules 
N02 which recombine to N204 upon cooling. We here meet 
the interesting case of dissociation, occurring even at the 
ordinary temperature. N204 is colorless, while N02 is colored 
red-brown ; it appears, therefore, that the color gradually 
becomes darker as the temperature rises, and that it corres- 
ponds to the increasing dissociation of the complex molecules 
n2o4. 
At ordinary temperatures, nitrogen tetroxide is a yellow 
liquid of sp. gr. 1.45. When cooled to — 20° it solidifies to 
a colorless crystalline mass, melting at —120. In consequence 
of the dissociation that sets in at o°, the liquid, at first color- 
less, becomes yellow, and the intensity in color grows with 
rising temperature. The liquid begins to boil at about 26°, 
and is converted into a yellowish-brown vapor that becomes 
dark as the temperature is increased. 
The theoretical vapor density of N20< (molecular weight = 91.86) 
equals 45.9, while that of NOa (45.9) — 22.9. The experimental vapor 
density has been found to equal 38 at the point which the liquid compound 
boils (26°); it may be calculated from this that, at this temperature, 34.4 
per cent, of the N204 molecules are decomposed into N02 molecules. 
Hence we conclude that the dissociation of the compound N204 com- 
mences already in the liquid state; this is confirmed by the yellow color- 
ation appearing at o°. Sulphuric acid, as we saw (p. 190), exhibits a sim- 
ilar dissociation in liquid condition. With rising temperature the density 
of the vapor steadily diminishes, becomes constant finally at 150° and 
equals 22.9 Then all the molecules (N204) are decomposed into the 
simpler molecules N02; and the dark coloration of the vapors attains its 
maximum. 
Nitrogen tetroxide is formed by the union of two volumes 
of nitric oxide with one volume of oxygen : 
2 Vols. 
2NO + o2 = n2o4. 
I Vo!. 
We can get it more conveniently by heating dry lead nitrate, 
which decomposes according to the following equation: 
(NOs)2Pb — PbO + O + 2N02. NITROSYLSULPHURIC ACID. 
205 
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 chem- 
ical reactions. We saw that by the action of a little cold water, 
the tetroxide was decomposed into nitrogen trioxide and nitric 
acid (p. 203). With excess of cold water, and also with an 
aqueous solution of alkalies, it yields nitric and nitrous acids, 
that is, their salts : 
no2 
NO + H20 = NOj.OH -f- NO.OH. 
NO/" 
Both reactions plainly indicate that the liquid tetroxide 
represents themixed oxideoinitric and nitrous acids; similarly, 
the compound C1204 constitutes the mixed oxide of chloric 
and chlorous acids (p. 176). 
Warm water converts the tetroxide into the dioxide, N02, 
which in turn yields nitric acid and nitrogen oxide: 
3N02 + H20 = 2HNO3 + NO. 
The tetroxide and dioxide possess strong oxidizing proper- 
ties; many substances burn in their vapors; iodine is set free 
from the soluble metallic iodides by them. 
/O.NO. 
Nitrosylsulphuric Acid, SO.NH — S02<^ 
X)H. 
This compound, termed nitrosulphonic acid, is an intermediate 
product in the manufacture of commercial sulphuric acid (see 
p. 188), and is quite important. It is employed in the analy- 
tical determination of the nitrogen oxides. It is produced by 
conducting nitrogen trioxide and tetroxide into concentrated 
sulphuric acid : 
2SO>\oh + N<°> = 2SO>(oh° + ”>°- 
S0<<0H + 5o>° = SO<OH° +NO,OII. 
Nitrogen monoxide—NO—is not absorbed by pure sulphuric 
acid, but will be if the same contains nitric acid : 
O.NO 
3S04H2 -f N03H -j- 2NO = 3SO./" + 2H20. 
OH 206 
INORGANIC CHEMISTRY. 
Further, the nitrosylsulphuric acid results from the com- 
bined action of sulphurous oxide, nitrogen tetroxide, and a 
little water: 
O.NO 
2S02 + N20* + O + H20 = 2SC\ 
\ 
OH. 
It is obtained most readily by conducting sulphur dioxide into 
strongly cooled anhydrous nitric acid : 
O.NO 
SOj 4- N03H = S02 ; 
\ 
OH 
there results a thick magma, which may be dried upon porous 
earthen plates under the dessicator. 
Nitrosylsulphuric acid forms a leafy or granular crystalline, 
colorless mass (chamber-acid crystals, p. 189), which fuses 
about 730 and decomposes into its anhydride, sulphuric acid, 
and nitrogen trioxide. It deliquesces in moist air, and yields 
sulphuric and nitrous acids with water : 
O.NO OH 
/ / 
SO2 + H20 = S02 + NO.OH; 
\ \ 
OH OH 
the nitrous acid decomposes in heat into nilric acid and ni- 
tric oxide. 
Nitrosylsulphuric acid dissolves in concentrated sulphuric acid without 
any change; the solution, called nitroso acid, is also produced in the sul- 
phuric acid manufacture in the Gay-Lussac tower, is very stable and may 
be distilled without decomposition. When diluted with water it remains 
unaltered at first, but when the specific gravity of the solution reaches 
1.55—1.5° (51-48° IL), then all the nitrogen oxides escape, especially on 
warming. When the nitroso acid is poured into a large quantity of water, 
the nitrososulphuric acid breaks up into (like the pure acid, see above) 
sulphuric and nitrous acids, and the latter in part into NOsH and 2NO. 
Therefore, in titrating nitrous acid with permanganate of potassium 
(Mn04K. See p. 203), we get the results corresponding to nitrososulphuric 
acid, if the latter be poured into the permanganate (Lunge). 
All the nitrogen oxides and acids are separated as nitric oxide (NO) on 
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 nitroso acid by dilution with water 
and application of heat (see above.) This occurs more readily and com- 
pletely (even upon concentration to 58° B. = 1.679 specific gravity) by the 
action of sulphur dioxide : NITRIC OXIDE. 
207 
O.NO OH 
/ / 
2So2 4- 2H2o + so2 = 2So2 4- 2N0. 
,\ \ 
OH OH 
Upon this depends the denitrating action of the Glover tower—see p. 188. 
The anhydride of nitrosulphonic acid—S2N2Os = O j g( )2 C) NO's 
produced upon heating the latter (together with S04H2 and N2Os—see 
p. 206) beyond its point of fusion. It is obtained pure by saturating sul- 
phur trioxide with nitric oxide : 
3S03+ 2NO = O {io2o.NO + S°2- 
It is a crystalline, colorless mass, fusing at 2170, and boiling without 
decomposition about 360°. Much water decomposes it, the same as 
nitrosylsulphuric acid. 
The chloranhydride of nitrosulphonic acid, SN04C1 = S02<( 
is formed by the union of sulphur trioxide with nitrosylchloride: 
so3 + NOC1 = S°2\ciN0 
It forms 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 dilute nitric acid this oxide is formed, inasmuch 
as the hydrogen in statu nascendi reduces another portion of 
the acid. It is most conveniently obtained by pouring dilute 
nitric acid (specific gravity 1.2) upon copper filings : 
3Cu + 8HN0,, = 3Cu(N03)2 + 4H2O + 2NO. 
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 ex- 
pelled from the generating vessel by NO, and the gas collected 
over water after the interior of the apparatus has become 
colorless. 
Nitric oxide is a colorless gas, of specific gravity 14.98 
(H = 1) or 1.038 (air = 1), and can be condensed by cold 
and pressure. It is slightly soluble in water, but dissolves very 
readily in an aqueous solution of ferrous salts, imparting a 208 
INORGANIC CHEMISTRY. 
reddish-brown color to the liquid ; heat expels it from the 
same. 
Nitric oxide is readily soluble in nitric acid. As its' solution becomes 
more concentrated, it assumes a brown, yellow, green or blue color, as 
nitrogen trioxide is formed finally with anhydrous nitric acid: 
Potassium permanganate oxidizes it, like nitrous acid (p. 204), to nitric 
acid: 
2HN03 + 4NO = 3N2Os + H2O. 
ioNO + 6Mn04K -f 9S04H2 = ioNOsH + 3S04K2 + 6S04Mn + 4H20. 
As nitric oxide contains 53.3 per cent, oxygen, it is capable 
of sustaining the combustion of some substances, but to bring 
about the previous separation of oxygen from the nitrogen 
requires energetic reaction. Hence, phosphorus continues to 
burn in this gas, while a sulphur flame, developing only a 
slight heat, is extinguished ; ignited charcoal does the same, 
while a splinter, burning energetically, will continue to do so, 
when introduced in the gas. On shaking a cylinder filled 
with NO with a few drops of readily volatile carbon disulphide, 
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, emitting strong actinic rays ; in this combus- 
tion, 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 discovered 
that the same is greater (about 21,600 calories), than that which is de- 
veloped 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 O, than for the separation of the molecules of combined 
oxygen atoms—an additional proof that the molecules of free oxygen (as 
of other elements) consist of atoms (p. 113.) 
With oxygen, NO at once forms brown vapors of nitrogen 
dioxide: 
2NO + 02 = 2N02. 
With less oxygen, nitrogen trioxide is produced (p. 202). 
NO also combines with chlorine to nitrosylchloride, NOC1 
(p. 202), and the compound NOCl2, which is, as yet, little 
investigated ; with bromine, it yields NOBr3. At a red heat 
NO becomes N02 and N. With hydrogen and moderate heat 
it forms water and nitrogen : NO + H2 = N H20 ; a mix- 
ture of both gases burns with a white flame. On conducting 
2 vols. 
I vol. 
2 vols. NITROUS OXIDE. 
209 
NO and H together over platinum sponge, water and ammo- 
nia‘are produced : 
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 into the same 
a piece of sodium and 
heat the latter with a 
lamp. The sodium com- 
bines with the oxygen, 
and free nitrogen sep- 
arates ; the volume of 
the latter always equals 
half the volume of the nitric oxide gas employed ; this fol- 
lows from the formula NO : 
2NO + 5H2 = 2NH3 + 2H2Q. 
Fig. 85. 
2 vols. 
2no = n2 + o2. 
I vol. 
I vol. 
The molecular formula of the oxide is NO = 29.97, as its vapor density 
is 14.98 (H = 1). NO, N02 and chlorine dioxide, C102 (p. 177), present 
an apparent anomaly as regards the common laws regulating the valence 
of the elements. Ordinarily, the valence changes from an odd number to 
an odd, and from an even to an even number (p. 170). Nitrogen is usually 
pentavalent and trivalent; in the cited compounds it appears di- and 
valent. This abnormal behavior of N finds a partial explanation in the 
position it occupies in the periodic system of the elements. 
Nitrous Oxide — Hyponitrous 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: 
This compound is a colorless gas, of sweetish taste and slight 
odor. Its density is 21.9 (H = 1), or 1.52 (air = 1), corre- 
sponding to the molecular formula N20 = 43-98. In cold 
water it is tolerably soluble (1 volume H20 dissolves at o° 
1.305 volumes N,20) ; therefore it must be collected over warm 
water or mercury. Cooled to —88°, or under a pressure of 
30 atmospheres at o°, it condenses to a colorless liquid of 
specific gravity 0.937. By evaporation of the liquid in the 
air its temperature diminishes to — ioo°, and it solidifies to a 
cfystalline, snowy mass. If the aqueous nitrous oxide be evap- 
orated under an air pump its temperature falls to — 140° ; the 
NH4N03 = N20 + 2H20. 210 
INORGANIC CHEMISTRY. 
lowest which has been attained. Although this oxide contains 
less oxygen than nitric oxide, it supports, the combustion of 
many bodies more readily than the latter, because it is more 
easily decomposed into oxygen and nitrogen. A glimmering 
chip inflames in it, as in oxygen ; phosphorus burns with a 
bright luminous flame, while a sulphur flame is extinguished. 
The liquid nitrous oxide behaves like the gas; an ignited 
coal thrown on its surface burns with a bright light. A mix- 
ture of equal volumes of nitrous oxide and hydrogen explodes 
like detonating gas, only less violently: 
n2o + h2 = N2 + H20. 
It resembles oxygen very much, but may be distinguished 
from it by not producing brown vapors with nitric oxide, as 
does the former. It is not capable of combining with oxygen. 
When it is conducted through a red-hot tube, it has an exhila- 
rating effect when inhaled in slight quantity, and is, therefore, 
termed laughing gas. 
Its volume composition may be determined in the same 
manner as with nitric oxide, viz. : by heating a definite volume 
of the gas with potassium. Then we learn that from a volume 
of N20 an equal volume of nitrogen will be separated—cor- 
responding to the molecular formula: 
I vol. 
I vol. 
I vol. 
N20 + K, = N2 + K20. 
Hyponitrons Acid—NOH. As nitrous and nitric acids correspond 
to nitrogen trioxide and pentoxide, so hyponitrous oxide may be regarded 
as the anhydride of the recently discovered hyponitrous acid: 
1 vol. 
1 vol. 
NjO + HaO = 2NOH,. 
although the latter is not formed by the hydration of N20. On the con- 
trary, it yields nitrous oxide when water is withdrawn from it by sulphuric 
acid. 
An aqueous solution of the acid may be obtained by decomposing the 
silver salt, AgNO, with hydrochloric acid. Dissolved in water it is colorless, 
reacts strongly acid and is tolerably stable. It liberates iodine from 
potassium iodide and reduces a permanganate solution. The silver salt is 
again precipitated by silver nitrate. 
The alkali SijltS of hyponitrous acid result from the action of sodium 
amalgam upon the nitrates or nitrites, or of freshly precipitated ferrous 
hydrate upon sodium nitrite, and by the electrolysis of the nitrites. On 
neutralizing the solution with acetic or nitric acid, and adding silver nitrate, 
the silver salt is obtained as a light-yellow amorphous powder, not altered 
by the light. When heated above no0 it decomposes with explosion. If 
is soluble in dilute sulphuric acid, and is reprecipitated upon neutralization 
with ammonium hydrate. Concentrated sulphuric acid liberates N2O from 
both it and the free acid. OXYGEN COMPOUNDS OF PHOSPHORUS. 
The latest researches make the composition of hyponitrous acid rather 
doubtful, since either the formula N202Ag or N405Ag4 may be ascribed to 
the silver salt. 
The varying deportment of the oxygen compounds of nitrogen, their 
little stability, and oxidizing properties, as well as their mode of formation, 
find their explanation in the thermo-chemical relations. All nitrogen 
oxides are endothermic compounds, i.e., they are produced from their 
elements with heat absorption (compare p. 181) corresponding to the 
symbols; 
(N20) = —18,300 (N,0) = —21,570 (N2 = —23,000 
(N,02) = —2000 (N2,Os — gas) = —12,000. 
From this we observe that the oxides of nitrogen cannot be prepared from 
the elements without addition of energy, and their direct production has 
never been noticed. Proceeding from nitric oxide (NO), we see from the 
above numbers that the formation of the higher oxides from it occurs with 
heat disengagement; 
(2N0,0) = 20,140 (N0,0) = 19,570 (2N02,0) .= 2800, 
whereas heat is absorbed in the conversion of nitrous into nitric oxide i 
(N20,0) =-24,840. Plence nitric oxide combines readily with oxygen 
to yield N203 and N02, while nitrous oxide is not reactive. Conversely, 
the higher oxides are easily reduced to nitric oxide, and this explains their 
great ability to oxidize. The absorption of heat in nitric oxide (N,0 = — 
21,570) indicates, therefore, .that the affinity of N to Ois so much less than 
that of the oxygen atoms in the molecule (p. 208). 
Heat disengagement, on the contrary, occurs in the production of nitric 
acid from its elements; 
(N,03H—liquid) = 41,500 (N,Os,H,Aq.) = 49,090. 
This explains the relative stability of that acid, and the possibility of its 
production under different conditions. 
2. OXYGEN COMPOUNDS OF PHOSPHORUS. 
Hypophosphorous 
acid. 
po2h3 
P2O3 
POjjHg 
Phosphorus 
trioxide. 
Phosphorous 
acid. 
PA 
Phosphorus 
pent*xide. 
Orthophosphoric 
acid. 
po4h3 
The following anhydride acids (compafe p. 198) are derived 
from orthophosphoric acid ; 
HP03—Metaphosphoric acid. 
H4P207—Pyrophosphoric acid. 212 
INORGANIC CHEMISTRY. 
The structure of these compounds is expressed by the fol- 
lowing formulas: 
V 
H2PO — OH 
hpo/h 
v /OH 
PO—OH 
,\OH 
Hypophosphorous 
acid. 
Phosphorous 
acid. 
Phosphoric 
acid. 
In hypophosphorous acid two atoms of hydrogen are directly 
combined with pentavalent phosphorus, while the third atom 
forms an hydroxyl group with oxygen. The latter is easily re- 
placed by the-action of bases, and, therefore, hypophosphorous 
acid is a monobasic acid. 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 results— 
an anhydride, which, at the same time, is a monobasic acid, as 
it contains one hydroxyl group: 
POa — OH — Metaphosphoric acid. 
On removing one molecule of H20 from two molecules of 
HsP04, pyro- or diphosphoric acid is formed, (see p. 191): 
v /OH /OH 
PO—OH PO—OH 
ho — \o 
v /OH — /u 
PO—OH PO—OH 
\OH . \OH 
Pyrophosphoric acid contains 4 hydroxyl groups, hence is 
tetrabasic. 
Finally, if, from two molecules of phosphorous acid or phos- 
phoric acid, all the H atoms be removed, in the form of water, 
two perfect anhydrides are obtained: 
2 Molecules Phosphoric acid. 
i Molecule Diphosphorio acid. 
Ill III 
OP — o — PO 
V V 
02P — O — P02 
Phosphorous • 
anhydride. 
and 
Phosphoric 
anhydride. 
The salts of phosphoric acid are termed phosphates; those 
of phosphorous acid, phosphites, and of hypophosphorous acid, 
hypophosphites. 
Hypophosphorous Acid—*H3P02. Hydrogen phosphide 
escapes when a concentrated solution of sodium or barium PHOSPHOROUS ACID. 
213 
hydrate is warmed with yellow phosphorus, leaving behind in 
solution, a salt of hypophosphorous acid : 
4P + 3 NaOH + 3H.0 = 3H.PO.ONa + PHS. 
The free acid may be separated from the barium sftlt by 
means of sulphuric acid ; the insoluble barium sulphate being 
filtered off from the aqueous solution of the acid, and the lat- 
ter concentrated under the air-pump. Hypophosphorousacid 
is a colorless, thick liquid, with a strong acid reaction. 
Below o° it. sometimes solidifies to large white leaflets, which 
fuse at -f- 17.40. Heat converts it, with much foaming, into 
hydrogen phosphide and phosphoric acid: 
It absorbs oxygen readily, becoming phosphoric acid, hence 
acts as a powerful reducing agent. It reduces sulphuric acid to 
sulphur dioxide, and even to sulphur. It precipitates many 
of the metals from their solutions; from copper sulphate it 
separates the hydride—Cu2H2. 
The acid is monobasic. Its salts dissolve readily in water, 
and absorb oxygen from the air, thus becoming phosphates. 
When heated in. a dry condition, they set free the hydride 
of P, and are converted into pyrophosphates; some also yield 
metallic phosphides. ' 
2po2h3 = ph3 + po4h3. 
Phosphorous Acid—H3P03—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 
conveniently : 
PC13 + 3h2o - po3h +• 3HC1. 
By evaporating this solution under the air-pump the phos- 
phorous 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 PH3, and phosphoric 
acid : 
4FO3H3 = Pll3 + 3P04H3. 
In the air the acid absorbs oxygen, and changes to phos- 
phoric acid. Hence, it is a strong, reducing agent, and pre- 
cipitates the free metals from many of their solutions. In the 
presence of water, the halogens oxidize it to phosphoric acid. 
It is a dibasic acid, forming two series of salts, in which i 214 
INORGANIC CHEMISTRY. 
a,nd 2 atoms* of H are replaced by metals. In the air, the 
phosphites do not oxidize, except under the influence of oxi- 
dizing agents. When heated, they generally decompose into 
hydrogen and pyrophosphates. 
Phosphorus Trioxide—P2Os.—Phosphorous anhydride 
results from passing a slow, dry air current over gently heated 
phosphorus; further, by the'action of the trichloride upon 
phosphorous acid : 
po3hs + PC13 = P203 + 3HCI. 
It forms white, voluminous flakes, which sublime readily 
and* possess a garlic-like odor. Water dissolves the oxide, 
yieldiug phosphorous acid. It extracts oxygen and moisture 
very energetically from the air, and becomes phosphoric acid. 
Phosphoric Acid—P04H,—or Orthophosphoric acid, is 
produced when the pentoxide is dissolved in hot water, and 
by the decomposition of the pent-a- or oxy-chloride (POC1), 
by water (see p. 137). It may be obtained by decomposing 
bone ash (P04)2Ca, with sulphuric acid, or, better, by oxidiz- 
ing yellow phosphorus with nitric acid. The aqueous solution 
is evaporated 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 (P04K,). 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 replaced by metals, as e.g.} monopotassium 
phosphate (H2KP04), potassium (K2HP04) phosphate, and 
tripotassium (K,P04) phosphate. 
The tertiary phosphates, excepting the salts of the alkalies, 
are insoluble in water. With a silver nitrate (AgNO,) solu- 
tion soluble phosphates give a yellow precipitate of tri silver 
phosphate, P04Ag3. 
Pyrophosphoric Acid—H4P207—(structure p. 212) is 
formed by the continuous heating of orthophosphoric acid to 
2oo°-3oo°, until a portion of it dissolved in ammonium hy- 
drate does not yield a yellow but pure white precipitate with 
* Therefore, the structural formula, IiPO(OH)2 is assigned to this 
acid. There appears to exist another phosphorous acid, at least in com- 
pounds, to which the formula P(OH)3 belongs. HYPOPHOSPHORIC ACID—METAPHOSPHORIC ACID. 
21 5 
silver'nitrate. The sodium salt is easily obtained by heating 
di-sodium phosphate: 
2Na2HP04 = Na4P207 + II20. 
The acid presents a white crystalline appearance, and is 
readily soluble in water. When in solution, it slowly takes 
up water at ordinary temperatures, more rapidly when heated, 
and, like all anhydrides, passes into the corresponding acid— 
orthophosphoric 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 yield a white precipitate, Ag4P207, with silver nitrate. 
Hypophosphoric Acid—P206II4.—While pyrophosphoric acid is an 
anhydride acid of phosphoric acid (p. 212), the so-called hypophosphoric 
may be viewed as a mixed anhydride of phosphoric and symmetrical phos- 
phorous acids: 
PO(OH), 
PO(OH)3 _ _ 
P(OH)3 — ■ /u 
P(OH>. 
It is produced together with phosphorous and phosphoric acids, by the 
slow oxidation of moist phosphorus in the air. It is separated from these 
acids by means of its difficultly soluble sodium salt, P206Na2Il2 + 6H20; 
by precipitating the solution of the latter with a soluble lead salt we get 
insoluble P2OfiPb2. Its silver salt is more easily obtained by oxidizing 
phosphorus in the presence of silver nitrate. The free acid from the lead 
or silver salt is rather stable in a dilute solution and below 30°, may be 
concentrated to a syrup. At higher temperatures, more readily in the 
presence of hydrochloric or sulphuric acid, the acid decomposes into phos- 
phoric and phosphorus acids. It is not a reducing agent, but is oxidized 
by potassium permanganate to phosphoric acid. 
Metaphosphoric Acid—HPOa or P020H—results upon 
heating the ortho- or pyro-acid to 400°. It can be more con- 
veniently obtained by dissolving the pentoxide in cold water: 
P2Os -f H20 = 2HPO3. 
It is a glassy, transparent mass (Acidum phosphoricum gla- 
ciate), which fuses on heating, and volatilizes at higher tem- 
peratures, without suffering any change. It deliquesces in 
the air, and dissolves with ease in water. (The commercial 
glassy phosphoric acid contains sodium and .magnesium phos- 
phate, and dissolves with-difficulty in water.) The solution INORGANIC CHEMISTRY. 
coagulates albumen ; this is a characteristic method of distin- 
guishing the meta- from the ortho- and pyro- acids. In aqueous 
solution, the acid changes gradually, by boiling rapidly, into 
the ortho- acid : 
It is a monobasic acid. Its salts, the metaphosphates, are 
readily obtained by the ignition of the primary salts of the 
ortho- acid: 
hpo3 + h2o = h3po4. 
NaH2P04 = NaP03 -f H20. 
When the aqueous solutions of these salts are boiled, they 
are converted into the ortho-primary salts. With silver 
nitrate the soluble metaphosphates give a white precipitate, 
AgPOj. 
In addition to the ordinary salts of metaphosphoric acid, various modifi- 
cations of the same exist; these are derived from the polymeric meta- 
acids, H2P206, H3P30g, H4P4012, etc. These acids arise from the corres- 
ponding polyphosphoric acids, which are obtained by the union of n 
molecules of the ortho-acid, with the separation of n — i molecules of water 
(p. 191), just as the meta-acid is formed'from the ortho-. They are all 
changed to primary ortho-pjrosphates by boiling their solutions. 
Phosphorus Pentoxide—P205, or Phosphoric anhydride, 
is obtained by burning phosphorus in a current of oxygen or 
dry air. 
Fig. 86. 
The following procedure serves for the preparation of it. (Fig. 86.) A 
piece of P, placed in an iron dish attached to the glass tube a b, is burned . 
in the gas balloon A. The necessary amount of air is drawn through the COMPOUNDS OF PHOSPHORUS WITH SULPHUR. 
217 
vessel by means of an aspirator. It is first passed through the bent tube 
containing pieces of pumice-stone, moistened with sulphuric acid, in order 
to dry it perfectly. After the phosphorus has been consumed, fresh pieces 
of it are introduced into the little dish through a b, 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 is a white voluminous, flocculent 
mass. It attracts moisture energetically and deliquesces in 
the air. It dissolves in cold water with hissing and yields 
metaphosphoric acid. Owing to its great affinity for water it 
serves as an agent for drying gases, and also for the withdrawal 
of water from many substances. 
Chlor-Anhydrides of the Acids of Phosphorus.—The halogen 
derivatives of P, considered on page 136, may be viewed as the halogen 
anhydrides of phosphorous and phosphoric acids (p. 192). The compounds 
PCI3, PBr3, and PI3, are derived from phosphorous acid, because they yield 
the latter acid with water: 
The compounds POCl3, POBr,, are the halogen anhydrides of phos- 
phoric acid: 
PC13 + 3H2O = H3PO3 + 3HC1. 
POCl3 + 3H2O = PO (OH), + 3HCI; 
while PC15 and PBr5 correspond to the normal hydroxide, P(OH)5, which 
has not been obtained in a free condition. 
The compound PSCI3 is analogous to the oxychloride POCI3. It is 
obtained by the action of PCI. upon hydrogen sulphide and some metallic 
sulphides : 
PCI. + H2S = PSCI3 + 2HC1. 
This reaction is very similar to that occurring in the formation of phos- 
phorus oxychloride. Phosphorus sulphochloride—PSCI3—is a colorless 
liquid, fuming in the air and boiling at 1240. Water decomposes it into 
phosphoric and hydrochloric acids and hydrogen sulphide. 
COMPOUNDS OF PHOSPHORUS WITH SULPHUR. 
With sulphur phosphorus affords a number of compounds 
which are obtained by direct fusion of P with S. As the union 
of ordinary P with S usually occurs with violent explosion, 
red phosphorus should be employed in preparing these com- 
pounds. 
The compounds P2S3 and P2S5, analogous in constitution to 
P203 and P205, are solid crystalline substances, melting at 
higher temperatures and subliming without decomposition; 
P2S5 boils at 530°. Water changes them to hydrogen sulphide 
and the corresponding acids, phosphorous, and phosphoric. 
They combine with metallic sulphides to PS4K3) INORGANIC CHEMISTRY. 
which possess a constitution analogous to that of the salts of 
phosphoric acid (see sulpho-salts of arsenic). 
At ordinary temperatures, P2S and P4S are liquids, which 
inflame readily in the air. 
Besides the preceding, we have other phosphorus derivatives which 
contain N. These have been little studied, and at present offer little 
interest. Such compounds are PN2H (phospham), PNO, PNCI2. The 
so-called amid derivatives, POCI2NH2, and PO| NH2)s, are 
produced by allowing ammonia to act upon POCl3. In these chlorine is 
replaced by the amido-group NH2. 
3. OXYGEN DERIVATIVES OF ARSENIC. 
Arsenic trioxide. 
AS2O3 
As03II3 
Arsenious acid. 
Arsenic pentoxide. 
As205 
As04H3 
Arsenic acid. 
Arsenic Trioxide, As203, 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. It is*obtained metallurgically 
on a large scale as a by-product in the roasting of ores con- 
taining arsenic. The trioxide thus fprmed volatilizes and is 
collected in walled chambers, in which it condenses, in the 
form of a white- powder (white arsenic, poison flour). To 
render it pure, it is again sublimed in iron cylinders, and ob- 
tained in form of a transparent, amorphous, glassy mass (ar- 
senic glass), the specific gravity of which equals 3.78. Upon 
preservation, this variety gradually becomes non-transparent 
and porcelaneous, acquires a crystalline structure, and its spe- 
cific gravity decreases to 3.69. 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 when the solution of the glassy va- 
riety crystallizes it phosphoresces strongly in the dark, while 
the porcelaneous does not exhibit this property. Arsenic tri- 
oxide crystallizes in similar forms of the regular system, when 
its vapors are rapidly cooled, but upon cooling slowly, it as- 
sumes the shape of rhombic prisms ; therefore, it is dimor- 
phous. When heated in the air, it sublimes above 218°, with- 
out fusing; upon higher pressure, however (in sealed tubes), 
it melts to a liquid which solidifies to a glassy mass. 
The vapors of As203 have the vapor density 198 (H = 1). Correspond- 
ing to formula As203 (=197.7), the vapor density should be — 98.8. ARSENIOUS ACID—ARSENIC ACID. 
219 
The vapor density determined experimentally is just twice as great, there- 
fore the gas molecules of the trioxide possess the double formula As406. 
We have already noticed that the molecule of free arsenic consists of four 
atoms (As4): this complex arsenic group consequently is retained in the 
trioxide ; while in arsine (AsH3) and arsenious chloride (AsC13) the mole- 
cules contain but 1 atom of arsenic. 
The trioxide dissolves with difficulty in water; the solution 
possesses a sweetish, unpleasant metallic taste, exhibits but 
feeble acid reaction, and is extremely poisonous. The oxide 
is Very soluble in acids, and probably forms salts with them ; 
at least, on boiling a solution of 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. 145)- 
Nascent hydrogen converts the trioxide into arsine (AsH:f) ; 
but when 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 the anhydride separates out upon evaporation, 
In its salts (arsenites) it is tribasic and usually affords tertiary 
derivatives: 
AgjAsOj, Mg3(As03)?, 
The alkali salts, soluble in water, absorb oxygen from the 
air and serve as powerful reducing agents, they themselves 
becoming arseniates. 
Other salts exist which are derived from the meta-arsenious 
acid, HAsOr 
Arsenic Acid—H3As04—is obtained by the oxidation of 
arsenic or its trioxide with concentrated nitric acid or by means 
of chlorine. Upon evaporating the solution rhombic crystals 
of the formula H3As04-|- y2 H20 separate out; these deliquesce 
on exposure. They melt at ioo°, lose their water of crystal- 
lization and yield orthoarsenic acid H3As04, which heated to 
140-180° passes into pyroarsenic acid—H4As207: 
2HsAsO = As207H4 -f- H20. 
At 2oo° 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. • INORGANIC CHEMISTRY. 
220 
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 and02; in 
contact with water it gradually changes to arsenic acid. 
Orthoarsenic acid is readily soluble, and is a strong tribasic 
acid. Its salts—the. arseniates—are very similar to the phos- 
phates and are isomorphous with them. With the soluble 
salts silver nitrate gives a reddish-brown precipitate of tri- 
silver-arseniate, Ag3As04. 
COMPOUNDS OF ARSENIC WITH SULPHUR. 
Like phosphorus, arsenic, upon fusion with sulphur, yields 
several compounds. The metallic nature of arsenic is seen in 
these derivatives, because they, according to the common 
method of forming the metallic sulphides, can be obtained 
by the action of hydrogen sulphide upon the oxygen deriva- 
tives of arsenic: 
As203 “I- 3H2S — As2S3 -f" 3H2O. 
Arsenic Trisulphide—As2S3—is precipitated from solu- 
tions of arsenious acid or its salts by hydrogen sulphide, as a 
lemon-yellow amorphous powder. It may also be obtained 
from arsenic acid solutions, but then it contains admixed S, 
as the acid is first reduced to arsenious acid and then pre- 
cipitated : 
This compound is readily prepared by fusing As203 with 
sulphur. It occurs as auripigment in nature in the form of a 
brilliant, leafy, crystalline mass of gold-yellow color, and the 
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, but dissolves readily in ammonium 
hydrate and the alkalies. 
Arsenic Pentasulphide—As2S5—separates as a bright 
yellow powder from the solution of sodium sulpharseniate, 
Na3AsS4 (see below), upon the addition of acids. 
The Arsenic Disulphide—As2S2—also exists. It occurs 
in nature as Realgar, forming beautiful, ruby-red crystals, of 
specific gravity 3.5. It is applied as a pigment. It is pre- 
pared artificially by fusing As with S. 
As205 4" 2HsS = As203 -)- 2H2O + 2S'. OXYGEN COMPOUNDS OF ANTIMONY. 
221 
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, H3AsS+. However, these acids are un- 
known in a free state, although their salts, known as sulphur- or sulpho- 
salts, are found, and they correspond perfectly with the oxygen salts. Just 
as the latter arise by the union of metallic oxides with acid oxides, so the 
sulpho-salts are formed by the combination of alkaline sulphides with those 
sulphur derivatives corresponding to the acid oxides (acid sulphides): 
As2S3 -{- 3K2S = 2K3AsS3 
Tripotassium sulpharsenite. 
As2S5 + 3K2S = 2K3AsS4 
For the preparation of these siilpho-salts, arsenic sulphide is dissolved in 
the aqueous solution of potassium or sodium sulphide, or hydrogen sul- 
phide is conducted through the alkaline solution of the oxygen salts; 
Tripotassium sulpharseniate. 
The sulpho-salts of the alkalies and ammonium are easily soluble in water, 
and when the solution is evaporated they generally separate in crystals. 
Acids decompose them, arsenic sulphide separating out and hydrogen sul- 
phide becoming free: 
K3As04 + 4H2S = KsAsS4 + 4H2O. 
2K3AsS4 + 6HC1 = As2S6 -f 6KC1 -f 3H2S. 
Antimony, carbon, tin, gold, platinum and some other metals form sulpho- 
salts similar to those of arsenic- (and of phosphorus). 
4. OXYGEN COMPOUNDS OF ANTIMONY. 
The oxygen derivatives of antimony are analogous in con- 
stitution to those of arsenic: Sb203 and Sb205. The metallic 
nature of antimony, which we observed appearing in the 
halogen derivatives shows itself quite distinctly here. The 
lowest oxygen compound does not possess acid, but basic 
properties almost solely; it forms salts with acids only, hence 
is'called Antimony oxide. The normal hydrate H3Sb03, cor- 
responding to arsenious acid, H;!As03, is not known. An 
hydrate, Sb02H 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 yields 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, H4Sb207, and meta-arrtimonic acid, 
HSbO.,; these exist in a free condition. 
Antimony Oxide—Sb203—is obtained by burning the 
metal in the air, or by oxidizing it with dilute HNQ3. By 222' 
INORGANIC CHEMISTRY. 
sublimation it may be obtained in two different crystal sys- 
tems, 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 anti- 
mony hydrate or antimonious acid, HSb02, separates out: 
2SbCl3 + 3Na.2C03 -f H20 = 2SbO.OH.-f 6NaCl + 3C02. 
The hydrate is changed to oxide by boiling. The latter and 
the hydrate dissolve in sodium and potassium hydrate, and, 
very probably, form salts (NaSbO,) which decompose upon 
evaporating the solution. In this behavior the acid nature of 
antimony hydrate is also seen; therefore it has received the 
name of antimonious*acid. 
The oxide forms salts with acids, which are derived either 
from the normal hydrate, H3Sb03, or from the hydrate, 
HSb02 = SbO.OH. In the salts of the first kind we have 3 
hvdrogen atoms of the hydrate replaced by acid radicals, or, 
what is the same, a trivalent antimony atom displacing 3 hy- 
drogen atoms of the acids: 
Sb03(N02)3 or (N03)3Sb. 
In the second variety of antimony salts derived from the 
hydrate, SbO.OH, hydrogen is replaced by a monovalent acid 
residue, or the hydrogen of the acid is substituted by the 
monovalent group, SbO, known as antimonyl: 
Antimony nitrate. 
SbO.O.NO* or N03.SbO. 
Antimonyl nitrate." 
Of these salts may be mentioned the following: 
Antimony Sulphate—(S04lsSb2—which separates when a solution of 
the oxide in sulphuric acid is cooled. 
Antimonyl Sulphate—SOp SbO)2—is formed when antimony oxide is 
dissolved in somewhat dilute sulphuric acid, and crystallizes in fine needles 
on cooling. Water decomposes both, forming basic’salts; hence the basic 
nature of antimony oxide is slight. 
Antimonic Acid—HSbO,—or, more correctly, Meta-anti- 
monic acid, is obtained upon warming antimony with concen- 
trated 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 in- 
soluble. 
If antimony pentachloride be mixed with much water Py- 
roantimonic Acid, H4Sb.20., separates as a white powder. COMPOUNDS OF ANTIMONY WITH SULPHUR. 
223 
Its salts are produced by fusing antimonic acid or meta-anti- 
moniates with potassium or sodium hydrates: 
Hydrochloric acid precipitates pyroantimonic acid from the 
solutions of these salts. 
By gentle ignition the meta- and pyro-acids yield Anti- 
mony Pentoxide, Sb205, a yellow, amorphous mass, soluble 
in hydrochloric acid. By heating the oxygen compounds for 
some time with air access they are .converted into the oxide, 
Sb2Oi} which can be viewed as antimonyl antimoniate (Sb03. 
SbO), or as a mixed anhydride g^Q2 j O. It is a white pow- 
der, becoming yellow when heated, and is non-volatile. 
•2KSb03 + 2KOH = K4Sb207 + • H20. 
COMPOUNDS OF ANTIMONY WITH SULPHUR. 
* These are perfectly analogous to the S compounds of ar- 
senic, and form sulphosalts with alkaline.sulphides, correspond- 
ing 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 precipi- 
tating a solution of the oxide with hydrogen sulphide, is an 
amorphous red powder. When 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, 
can be artificially prepared, and serves as a beautiful red color, 
under the name of antwiony cinnabar. Kermes mineral, -em- 
ployed in medicine, is obtained by boiling antimony sulphide 
with a sodium carbonate solution, and is a mixture of Sb2S3 
and Sb203. 
Antimony Pentasulphide—Sb2S5—or gold sulphur (sul- 
phur auratum) is precipitated by H2S from acid solutions of 
antimonic acid; it is more conveniently obtained by the pre- 
cipitation of sodium sulphantimoniate, Na3SbS4, with hydro- 
chloric acid: 
2Na3SbS4 + 6HC1 = Sb2S5 -f 6NaCl + 3H2S. 
It is an orange-red powder, like the trisulphide; it decom- 
poses on being heated into Sb2S3 and S2. It dissolves to anti- 224 
INORGANIC CHEMISTRY. 
mony trichloride in strong hydrochloric acid, with separation 
of sulphur and hydrogen sulphide. 
•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 H20 (SbS4- 
Na3 + 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 for hydrogen dimin- 
ishes with increase of atomic weight, corresponds to the addition of 
metallic character, while the affinity for chlorine, concluded from the ther- 
mo-chemical relations, generally increases (compare p. 145). However, 
the heat disengagement in the formation of AsCU is somewhat less than 
that in the case of PC13, which would afford a partial explanation for the 
non-existence of the compounds AsX5 (see p. 142). The slight affinity of 
arsenic is seen more distinctly in the oxygen compounds, because, as in 
the case of the halogen and oxygen groups (pp. 181 and 197), the arsenic 
corresponding to'bromine and selenium — 
Br = 79.7 Se = 79 As = 75. 
shows a less heat disengagement.in the formation of its compounds: 
(N,04,H,,Aq.) = • 117,400 
(P,04,Hs,Aq.) = 305.300 
(As,04,H3,Aq.) = 215,200 
(N2,05,Aq.) = 29,800 
(P2A) = 363.8oo 
(As2,05) = 219,400 
Phosphoric acid is, therefore, more stable and more energetic than nitric 
and arsenic acids; nitric acid oxidizes phosphorus and arsenic to phos- 
phoric and arsenic acids. The latter acid is readily reduced to arsenious 
acid. 
VANADIUM. NIOBIUM. TANTALUM. 
Vd = 51.2 • Nb = 94 Ta = 182 
The three rare elements, vanadium, niobium and tantalum, are closely 
related to the Phosphorus group. They yield derivatives very much like 
those of P; but possess a more metallic character, and do not combine with 
hydrogen. They exhibit many characteristics similar to those of chromium, 
iron and tungsten, with which they are frequently associated in their nat-. 
urally occurring compounds (compare the Periodic System of the ele- 
ments). 
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 pre- 
pare it in a pure state; it is a grayish-white,-metallic, lustrous powder, 
of specific gravity 5.5. It is difficultly fusible, and does not change in the 
air. When heated, it burns to Vd205. 
Vanadium Trichloride—VdCl3—forms red plates, which readily deli- 
quesce in the air; it is not volatile. OXYGEN DERIVATIVES OF THE CARBON GROUP. 
225 
Vanadium Oxychloride—VdOCl3—results on heating a mixture of 
Vd203 and C in chlorine gas. It is a lemon-yellow liquid, of specific 
gravity 1.84, and boils at 120°. It fumes strongly in the air and decom- 
poses with water (analogous to phosphorus oxychloride) into vanadic acid 
and hydrochloric acid. Its vapor density equals 86 (VdOCl., = 173.2). 
Vanadium Oxide—Vd203—is a black powder obtained by heating 
Vd205 in hydrogen. It combines with O, to form Vd205. 
Vanadium Pentoxide—Vd205—or vanadic anhydride, is a brown 
mass obtained by fusing the naturally occurring vanadates with nitre, etc. 
It is soluble in the alkalies, and forms salts of Vanadic, H3Vd04, and 
Metavanadic acids, HVd03 with the metals. All these compounds are 
similar in constitution to those of P. In addition to these, vanadium forms 
other compounds, constituted like those of sulphur and chromium In this 
class belong VdCl2 (dichloride), the tetrachloride, VdCl4, vanadious oxide, 
VdO, vanadium dioxide, Vd02, and VdOCl2. The tetrachloride, VdCl4, is 
a red-brown liquid, boiling at 1540; its vapor density equals 96 (VdCl4 
= 192.6/. Niobium and tantalum are not known in a free state. The 
chlorides, NbCl5 and TaCl5, are volatile and are decomposed by water. 
The oxides, Nb205 and Ta2Os, form salts of niobic (II3Nb04) and tantalic 
(II3Ta04) acids with bases. 
OXYGEN DERIVATIVES OF THE ELEMENTS OF 
THE CARBON GROUP. 
According to analogy to the hydroxyl derivatives of the 
elements of the three first groups: 
we may conclude the existence, for the tetravalent elements— 
carbon, silicon and'tin, of the following normal hydroxides, 
corresponding to the halogen compounds, CCl4,SiCl4, and 
SnCl4: 
C103,0H S02(0H)2 PO(OH)3, 
C(OH)4, 
Normal 
Carbonic acid. 
IV 
Si(OH)4, 
Normal 
Silicic acid. 
IV 
Sn(OFI)4. 
Normal 
Stannic acid. 
These normal hydrates or acids have but little stability, and 
mostly exist only in some derivatives. By the separation of a 
molecule of water, they pass into 
CO3H2 
SiOgHj 
SnOjH* 
or 
IV 
CO(OH)2 
Carbonic acid. 
IV 
SiO(OH)2 
IV 
SnO(OH)2. 
Stannic acid. 
These hydroxyl derivatives deport themselves toward the 
normal just as the meta-acids of the elements of the N group do' 
to the ortho-acids (see p. 198). They constitute the ordinary 
acids of. the tetravalent elements, carbon, silicon and tin, and 
as they contain 2 hydroxyl groups, are dibasic. 
Silicic acid. 226 
INORGANIC CHEMISTRY. 
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: 
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. The normal carbonic acid (C(OH)4) 
corresponds to this, but it is not capable of existing. Indeed, 
the meta- or ordinary carbonic acid, H2C03, is also very un- 
stable and at once decomposes, when separated from its salts, 
into water and carbon dioxide, C02. Even silicic and stannic 
acids break up readily into water and their oxides: 
C = 11.96, N = 14.01, O = 15.96, FI = 19. 
Carbon dioxide. 
C02 
Si02 
Silicon dioxide. 
Stannic oxide. 
Sn02 
Carbon Dioxide—C02—or carbonic anhydride (generally 
called carbonic acid). It is produced when carbon or its 
compounds are burned in air or oxygen. It is 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 + 2MC1 = CaCl2 + COa + II20. 
Carbon dioxide is a colorless gas, of sweetish odor and taste. 
Its gas density equals 1.527 (air = x), or 21.94(1! = 1), 
corresponding to the -molecular formula, C02 = 43.89. 
Owing to its weight, the gas may be collected by air displace- 
ment, and may be poured from one vessel into another filled 
with air. Under a pressure of 36 atmospheres (at o° C.), car- 
bon dioxide condenses to a mobile, colorless liquid, not mis- 
cible with water, and boiling at —88°. The specific gravity 
of the liquid carbon dioxide is 0.99 at —io°, at o°, 0.94. 
Hence, it expands more strongly than any gases, although 
the coefficient of expansion of liquids is usually less than that 
of gases. Similar deportment is also observed in the case of 
other bodies compressible to liquids by great pressure. 
Calcium carbonate. 
Calcium chloride. OXYGEN COMPOUNDS OF CARBON. 
227 
Above 32.50 carbon dioxide cannot be condensed to a liquid by any 
pressure, 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 (p. 42). Liquid carbon dioxide sometimes occurs 
in minerals like quartz, enclosed in small cavities. It is distinguished 
from other enclosed liquids (e. g., water) in that it is gasified on heating 
the mineral above 32.50. 
When liquid carbon dioxide is 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 painful blister. By 
the evaporation of solid carbon dioxide under the air-pump, 
its temperature is lowered to —130° C. (compare p. 209). 
Water dissolves its volume of the gas at 140 ; at o° it takes up 
1.79 vols. This proportion remains constant for every pres- 
sure, 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 140 and two atmospheres, 2 
volumes, at 3 atmospheres 3 volumes, etc., of carbon dioxide 
—measured at ordinary pressure. The gas absorbed at higher 
pressure escapes with effervescence of the liquid when the 
pressure is diminished ; upon this depends the sparkling of 
soda water and champagne, which are saturated with CO,2 
under high pressure. Every naturally occurring water, espe- 
cially spring water, holds C02 in solution, which imparts to it 
a refreshing taste. 
As the product of a complete combustion carbon dioxide is 
not combustible, and is unable to support the combustion of 
most bodies,.a glimmering chip is immediately extinguished 
in it. In a similar manner it is non-respirable. Although it 
is not poisonous, in the true sense of the word, yet the admix- 
ture 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. 
It is decomposed by continued action of the electric sparks 
into carbonous oxide (CO) and oxygen ; upon heating to 
1300° it suffers a partial decomposition (dissociation) into 228 
INORGANIC CHEMISTRY. 
CO and O. It is 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 forms potassium carbonate (K2C03), with excess of 
C0.2. 
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 C0.2 it follows that in 
one volume it contains an equal volume of O. We may satisfy 
ourselves of this by burning C in a definite volume of O ; 
after cooling, there is obtained an equal volume of carbon 
dioxide: . 
c + o2 = co2. 
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 U tube is closed by means of a glass 
stopper, through which two copper wires 
pass. The one wire bears a combustion 
spoon at its end, upon which lies the carbon 
to be burned, while the other wire termin- 
ates in a thin piece of platinum, which is 
in contact with the carbon. For the per- 
formance of the experiment, the air is ex- 
pelled 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 current from 3-4 Bunsen ele- 
ments, which induces the burning of the 
carbon. As the volume of the enclosed gas 
is greatly expanded by the heat developed, 
it is advisable, in order to avoid the jumping 
out of the stopper, to previously reduce the 
pressure of the gas about two-thirds, by run- 
ning out mercury. A practical modification 
of this apparatus consists in having the usual cylindrical eudi- 
ometer limb provided with two side tubes, in which carbon 
electrodes can be inserted. This same apparatus can also be 
I vol. 
I vol. 
Fig. 87. CARBON MONOXIDE. 
229 
employed for the illustration of the volume relations observed 
in the combustion of sulphur and other bodies. 
The Physiological Importance of Carbon Dioxide.—The gas is present 
in the atmosphere, and is inhaled by the plants. The chlorophyl grains in 
the green parts of the plant decompose carbon dioxide in sunlight, with a 
partial separation of oxygen ; by the mutual action of water and ammonia 
the innumerable carbon compounds peculiar to plants are formed 
fr.oni the residue. The respiration and life process of animals are 
essentially the reverse of the above. These absorb the oxygen of the 
air through the lungs; and, influenced by the blood corpuscles, it 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 O by animals, and its separation 
by plants, as also the reverse course of C02, 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, ex- 
hibits 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. 
We may then regard it as probable that free carbonic acid, 
H2co3, is contained in the aqueous solution, but this readily 
decomposes into the dioxide, C02, and water. The salts of 
carbonic acid are produced by the action of carbon dioxide 
upon the bases: 
2K0H + c°2 = C03K2 -f H20. 
Potassium carbonate. 
Carbon dioxide is, therefore, easily absorbed by potassium 
and sodium hydrate. On conducting it through a solution of 
calcium or barium hydrate, a. white precipitate of barium or 
calcium carbonate, CaCO,, is produced. 
Carbonic acid is dibasic, forming primary (acid) and sec- 
ondary (neutral) salts, C03HK and C03K2, called carbonates. 
As the acidity of carbonic acid is only slight, the secondary 
salts, obtained from strong bases, exhibit a basic reaction. 
Most acids expel the weak carbonic acid from its salts, with 
decomposition into carbon dioxide and water. 
Carbon Monoxide—CO—is produced in the imperfect 
combustion of carbon by insufficient access of air, and when 
carbon dioxide is conducted over red-hot coals: C02+ C = 
2CO. The forms of apparatus described, p. 228, serve for the 
demonstration of this volume relation. 
Zinc dust reacts like carbon: 
C02 -j- Zn = ZnO -j- CO. 230 
INORGANIC CHEMISTRY. 
When carbon dioxide is conducted through a glass tube, containing zinc 
dust heated to a faint red heat, almost pure carbon monoxide escapes. 
A more convenient procedure consists in heating pulverized magnesium 
carbonate and zinc-dust in a glass retort, when CO, containing C02, is 
eliminated ; subsequently the former alone escapes. 
The monoxide is produced, further, by igniting carbon with 
different metallic oxides, e.g., zinc oxide : ZnO + C = Zn 
+ CO. For its preparation, oxalic acid is warmed with sul- 
phuric acid : the latter withdraws water from the former, and 
the residue breaks up into carbon dioxide and monoxide : 
. c2o4h2 = C02 + CO + II20. 
The disengaged mixture of gases is conducted through an 
aqueous solution of sodium hydrate, by which the C02 is ab- 
sorbed, the monoxide passing through unaltered. Pure mon- 
oxide may be prepared by heating yellow prussiate of potas- 
sium (see Iron) with 9 parts H2S04. The resulting gas is 
colorless and odorless, and can only be condensed with diffi- 
culty. Its specific gravity is 13.96 (H = 1), corresponding 
to the molecular formula, CO = 27.93. It is almost insoluble 
in water; but is dissolved readily by an ammoniacal solution 
of cuprous chloride (CuCl). When ignited, it burns in the 
air, with a faintly luminous flame, to carbon dioxide. With 
air or oxygen, it affords an explosive mixture: 
2C0 -f- 02 = 2CO2. 
2 vols. 
i vol. 
2 vols. 
In consequence of its oxidation, it is capable of reducing 
most metallic oxides at a red heat: 
CuO -f~ CO = Cu -f- CO2. 
Burning bodies are extinguished by it. When 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: 
I vol. 
• CO + Cl2 = COCI2. 
I vol. 
I vol. 
This is obtained by bringing together equal volumes of CO 
and Cl2 in direct sunlight, or, better, by conducting CO into 
SbCl5. It is a colorless, suffocating gas, of specific gravity 
49.3 (H = 1), agreeing with the molecular formula COCl2 = CARBON DISULPHIDE. 
231 
98.6. Water decomposes it into hydrogen chloride and car- 
bon dioxide : 
COCli + H2O =• C02 + 2HCI. 
‘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 ignited carbon the escaping disulphide 
vapors are condensed in a cooled receiver. It is a colorless, 
mobile liquid, of characteristic odor, solidifies at —:i 160, and 
refracts light strongly. Its specific gravity equals 1.29 at o°. 
It is very volatile, boils at 470, and burns with a blue flame, 
to carbon dioxide and sulphurous acid. When a mixture of 
carbon disulphide vapors and oxygen is ignited, .a violent ex- 
plosion ensues: 
CS2 + 3O2 = C02 + 2SO2. 
1 vol. 
3 vols. 
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 heat is absorbed by the evaporation, that the residual 
liquid solidifies to a white, snow-like mass. Carbon disulphide 
is insoluble in water; but mixes, in every proportion, with 
alcohol and ether. It dissolves iodine with a violet-red color, 
and serves as an excellent solvent for sulphur, phosphorus, 
caoutchouc, and the iatty oils. On conducting the CS2 vapors 
over heated zinc-dust, all the sulphur unites with the zinc, 
forming zinc sulphide, while the carbon separates as soot, 
Carbon disulphide may be viewed as the anhydride ofi sul- 
phocarbonic acid—H2CS3. The salts of this acid are obtained 
by the solution of CS2 in alkaline sulphides (see Sulpho-salts, 
’221): 
CS2 + K2S = K2CS3. 
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 
* Upon standing in sunlight CS2 is said to break up into S and CS—a 
chestnut-brown powder of specific gravity 1.66. 232 
INORGANIC CHEMISTRY. 
quantity) when a mixture of sulphur vapors and carbon mon- 
oxide gas is passed through red hot tubes and by heating 
carbon disulphide with sulphuric oxide: 
CS2 + 3SO3 = COS + 4SO2. 
It is most readily obtained from potassium sulphocyanide— 
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: 
2 vols. 
2C0S 4- = 2C02 + 2SOr 
3 vols. 
2 vols. 
2 vols. 
It is decomposed at a red heat into CO and sulphur. It is 
soluble in an equal volume of water, decomposing gradually 
into the dioxide and hydrogen sulphide: 
COS + H20 = CO* -f SH*. 
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-ca'lled yellow prussiate of potassium, K4Fe (CN)6. 
All the other cyanogen derivatives may be prepared from these 
two compounds. They all contain the group CN, called 
cyanogen. In it we have a trivalent nitrogen atom combined 
with.a tetravalent carbon atom; the fourth affinity of the latter 
III IV 
is not saturated : N=C—: it is similar, therefore, to the groups 
OH, NH2. CH3, and is a monovalent 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 heated 
with sulphuric acid : 
2KCN + H2S04 = KaS04 + ' 2IICN. 
Hydrogen Cyanide, HCN, is a colorless, mobile liquid, 
of peculiar odor, and boiling at 270. 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. OXYGEN COMPOUNDS OF SILICON. 
233 
If the CN group is separated from its salts it doubles itself, 
yielding dicyanogen or free cyanogen, C2N2 (N=C—C=N), 
because like the ether monovalent- groups (as CH3, see p. 
168), it cannot exist in a free condition. 
The heat occurring in the formation of the simplest carbon compounds 
above cited corresponds with the symbols : 
(0,0) = 28,590. (C0,0) = 68,370. (C,02) = 96,900. 
(C,0,S) = 1400. (C.S,) =—12,600. (C,N,H) =•—28,300. 
(C,H4) = 21,700. (C2,H4) =— 2700. (C2,H2) =—48,300. 
(CO,CL2) = 4300. (C02,Aq.) = 5880. 
If an element combine with another according to multiple proportions, 
there usually occurs, in the union of the first atom, a greater disengage- 
ment of heat than with the following atom (compare nitrogen oxides, p. 
211). The numbers above, on the contrary, show that the union of the 
second atom of oxygen with carbon (CO,Oi sets free 68,370 calories; 
that of the first atom (C,0), however, only 28,590 calories. This can 
only be explained by the fact that, for the vaporization and disaggregation 
of the solid carbon molecules, heat is necessary. If we assume that the 
direct union of the first atom, also disengaged 68,370 calories, it would 
follow from this that, in the dissociation of 12 parts carbon by weight into 
gaseous free atoms, 39,780 (= 68,370 — 28,590) calories were absorbed. 
This would explain the heat absorption in the production of CS2, CNH, 
C II2, while otherwise heat is invariably disengaged in every direct chemi- 
cal union. 
Comparing the elements of the carbon group with each other, we dis- 
c<?ver that the heat disengagement is greatest with the compounds of silicon: 
(C,C14) = 28,360. (Si,Cl4) = 157,600. (Sn,Cl4) = 129,200. 
(C,02) = 96,900. (Si,02) = 219,000. (Sn,02,H20)= 133,500. 
From these numbers we observe that tin dioxide, but not that of silicon 
can be reduced by carbon. 
OXYGEN COMPOUNDS OF SILICON. 
Silicon Dioxide, Si02 (Silica), is widely distributed in 
nature 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 
the silicic acids. It only occurs in nature crystallized 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; but is 
decomposed by hydrofluoric acid with the formation of silicon 
fluoride (SiFl4) and water (p. 161). Strong ignition with 
sodium or potassium reduces it to metallic silicon. The 
dioxide prepared artificially dissolves when boiled with potas- 
sium or sodium hydrate; some of the naturally occurring 234 
INORGANIC CHEMISTRY. 
amorphous varieties are also soluble, but not the crystallized 
dioxide. By fusion with the hydroxides or carbonates of the 
alkalies all varieties of silica yield a glassy mass (water glass) 
soluble in water and containing silicates (K4Si04 or K2SiO,). 
Upon the addition of hydrochloric acid to the aqueous solu- 
tion of the potassium or sodium salt, a very voluminous gela- 
tinous mass separates; this is probably normal silicic acid, 
H4Si04: 
Na4Si04 -f 4HCI = 4NaCl -f H4Si04. 
It becomes a white amorphous powder having the composi- 
tion H2SiOs when washed with water and dried in the air. The 
freshly precipitated acid is somewhat soluble in water, more 
readily in dilute hydrochloric acid. On adding a solution of 
sodium silicate to an excess of dilute hydrochloric acid the 
separated silicic acid remains dissolved. From the hydro- 
Fig. 88. 
chloric acid and sodium chloride solution vve can obtain a 
perfectly pure aqueous solution of silicic acid by dialysis 
by proceeding in the following manner. Pour the hydro- 
chloric acid solution into a wide cylindrical vessel whose 
lower opening is covered with animal bladder or parchment 
paper, and then suspend the vessel (dialyser) in another con- 
taining pure water. (Fig. 88.) Osmosis now sets in. The 
sodium chloride and hydrochloric acid pass through the parch- 
ment paper into the outer water, while on the other hand, 
water passes from the outer vessel into the dialyser; the parch- 
ment paper is not permeable to silicic acid. This alternate 
diffusion of the different particles occurs until the outer and 
inner liquids show the same quantity of diffusible substances. 
Upon introducing the dialyser into a fresh portion of water, 
the .dialysis commences anew. Finally, after repeated renewal 
of the external water, the dialyser will contain a perfectly SILICON DISULPHIDE. 
235 
pure silicic acid solution, free from sodium chloride and 
hydrochloric acid. The solution may be concentrated by 
evaporation; it then readily passes into a gelatinous mass. 
The same occurs instantaneously 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, albu- 
men, starch, glue (*oXXa, hence the name colloid), and especially most of the 
substances which occur chiefly in vegetable and animal organisms. Like 
silicic acid these colloids exist in liquid, soluble, and solid gelatinous 
condition. Many other substances (like ferric and aluminum oxides) 
which ordinarily are insoluble, can be brought into aqueous solution by 
dialysis. 
We have already seen that acids like sulphuric, phos- 
phoric, and arsenic, are capable of forming anhydro- or 
poly-acids by the union of several molecules and the elimina- 
tion of water. Silicic acid is particularly inclined to that 
kind of condensation. It forms a large number poly-silicic 
acids, Si203(0H)2, Si304(0H)4, Si305(0H)2, etc., derived 
from the normal and ordinary acid, according to the com-, 
mon formula: 
wSi(OH)4 — nUtO. 
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 per cent. H20 by 
ignition, represent such poly-acids. The natural silicates are 
the salts 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., chry- 
solite—Mg2Si04. 
Corresponding to CS2 is 
Silicon Disulphide, SiS2, which may be made by heating amorphous 
silicon with sulphur, or by conducting sulphur vapors over an ignited mass 
of silica and carbon. It sublimes in shining, silky needles, which water 
changes to silicic acid and hydrogen sulphide 
Tin also (p. 162) belongs to the group of carbon and sili- 
con. It affords oxygen derivatives perfectly analogous to 
those of these two elements. They, however, distinctly show 
the metallic character of tin and the basic nature of the 
oxides. The tin hydrates, Sn(OH,'4 and SNO(OH)2, are weak 236 
INORGANIC CHEMISTRY. 
acids, which with the alkalies yield basic reacting metallic 
salts that are rather unstable. The basic character is more 
evident in the lower stage of oxidation—stannous oxide, SnO, 
and hydrate, Sn(OH)2. Hence, we will consider tin with the 
metals. 
TITANIUM. 
ZIRCONIUM. 
THORIUM. 
Just as vanadium, niobium, and tantalum attach themselves to the ele- 
ments of the phosphorus group, so the three elements, titanium, zirco- 
nium, and thorium, stand in relation to the silicon group: (See Periodic 
System, p. 244.) 
Ti = 48. 
• • Zr = 90. 
Th = 232. 
P = 30.9 
Sb = 119.6 
As = 74.9 
V = 51 
Nb = 94 
T • s 182 
Si = 28 
Sn == 117.5 
Ti — 48 
Zr = 90 
Th = 232 
In all their deportment, they strongly resemble tin; they possess, how'- 
ever, a more metallic character in their derivatives. They are tetravalent, 
affording compounds of the form MeX4, in which X represents monavalent 
elements and groups; those of the form MeX2, corresponding to the stan-- 
nous derivatives, are unknown. The hydrates, Me(OH)4and MeOi OH 
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,Th. Corresponding to this, the acidity of the hydrates, i.e., 
their capability of exchanging H for metals, gradually diminishes. Tho- 
rium hydroxide, Th(OH)4, is not able to form metallic salts. 
. TITANIUM. 
This metal occurs in nature as titanium dioxide (rutile, anatase, brookite) 
and in titanates (perofskite, Ti03 Ca, Menaccanite, FeTiOg). Free titanium 
is a gray, metallic powder, obtained by heating potassium titan-fluoride 
(K2TiFl6i with potassium. It burns when heated in the air, and decom- 
poses water on boiling. It dissolves in hydrochloric and sulphuric 
acids, with evolution of hydrogen. 
Titanium Chloride—TiCl4—is formed, like silicon chloride, by con- 
ducting chlorine over an ignited mixture of the dioxide and carbon. A 
colorless liquid, of specific gravity i .76, fuming strongly in the air, and 
boiling at 1.36°. The vapor density equals 94.7 (H == 1), corresponding 
to the molecular formula TiCl4 = 189.4. It is not known in a free state. 
It behaves like tin tetrachloride with water. A compound Ti2Cl6, analo- 
gous to C2C16, is known. 
Titanium Fluoride—TiFl4—is not known in a free condition, but 
forms beautifully crystallized double salts, e. g., TiFl4, 2K.Fl, correspond- 
ing to the silico-fluorides (K2SiFl6). 
Titanic Acid—H4Ti04—separates as a w hite, amorphous powder, on 
adding ammonium hydrate to the hydrochloric acid solution of the titan- 
ates. When dried over sulphuric acid it loses 1 molecule H2O and be- 
comes TiO(OH)2. Titanic acid, like silicic and stannic acids, forms 
Ti = 48. ZIRCONIUM THORIUM—BORON. 
237 
poly-acids. The hydrates dissolve in alkalies and strong acids, to form 
salts. On igniting the hydroxides we get 
Titanuim Dioxide—Ti02, which may be procured crystallized as rutile, 
brookite, and anatase. When ignfted in a stream of hydrogen it changes 
to the oxide Ti203. Titanium dioxide is almost insoluble in the acids; 
it is only dissolved by hydrofluoric acid. It forms titanates upon fusion 
with the alkalies. 
The hydroxides, Ti04H4, Ti03FI2, etc., conduct themselves as feeble 
bases with strong.acids and afford salts with them [e.g., TiO.S04l; which 
are decomposed by water. The alkaline titanates (K2Ti03) are very un- 
stable. Other titanates occur in nature, e.g., CaTi03, MgTi03, and the 
so-called Titanic Iron, FeTi03. 
Titanium yields various compounds with nitrogen. When the dioxide is 
heated in ammonia gas, a dark violet powder of the composition TiN2 re- 
sults. The compound Ti5CN4—the so-called cyan-titanium nitride, is 
sometimes found in copper red metallic cubes, in blast-furnace slag, when 
iron ores containing titanium have been fused. 
ZIRCONIUM. 
Zr = go. 
Zirconium is very rare in nature, but is generally found in silicates, 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. r 5. Zirconium tetrachloride 
—ZrCl4, and fluoride, ZrFl4—are very similar to the correspondingtitanium 
compounds. 
Zirconic Acid or Hydrate—Zr(OH)4—is precipitated by ammonium 
hydrate, from acid solutions as a white voluminous precipitate, which be- 
comes Zr02, zirconium dioxide, upon ignition. Zirconic acid, when fused, 
is insoluble in potassium and sodium hydrates; it yields zirconates, 
Na2Zr03 and Na4Zr04, with the alkalies and their carbonates. Water de- 
composes these salts. 
The and hydrate dissolve when warmed with sulphuric acid, 
forming Zr(S04l2 which may be crystallized from water. 
THORIUM. 
Th = 232. 
Occurs very rarely, mostly in silicates (Thorite). Free thorium, sep- 
arated by sodium from the double chloride of potassium and thorium 
(ThCl4.2KCl), is a light-gray, crystalline powder, of specific gravity il.o, 
which burns in the air to the dioxide. 
Thorium Hydrate—Th(OH)4, and Thorium Dioxide, Th02, do not 
form salts with the alkalies. They dissolve in sulphuric acid to the sulphate, 
Th(S04)2,which crystallizes from water with four molecules of water. 
BORON.. 
B — I0.9. 
This element is generally classed with the metalloids, and 
stands isolated among them ; it forms the transition from these 238 
INORGANIC CHEMISTRY. 
to the metals, which is manifest from its position in the 
periodic system. On the one side, especially when free, it 
resembles carbon and silicon ; cm the other, it approaches the 
metals, beryllium, aluminium, and scandium (see the Periodic 
System of the Elements). As recently observed, it affords a 
gaseous hydride, but it is not very stable, and like stibine 
may be easily decomposed into its constituents. Its oxide, 
B203, although really of an acid nature, approaches such un- 
determined metallic oxides as aluminium oxide, A1203. Boron 
is trivalent, and yields only compounds of the form BX3. 
It is found in nature as boracic acid and in the form of 
borates, like borax (sodium salt), boracite (magnesium salt). 
It may be obtained free in an amorphous and crystallized 
state. The first results upon igniting boron trioxide with so- 
dium away from air contact; free boron and sodium borate are 
produced. 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. 
When fused with phosphoric acid it liberates phosphorus. 
Upon boiling with aqueous alkalies it dissolves, like beryllium, 
silicon, and aluminium, with formation of borates: 
The crystalline variety may be obtained by igniting boron 
trioxide with aluminium. Tne boron, separated by the alu- 
minium, dissolves in excess of the lattef, and crystallizes from 
it on cooling; upon dissolving the aluminium in hydrochloric 
acid the boron. remains in shining-, transparent, quadratic 
crystals, which are more or less colored, and have a specific 
gravity of 2.63. The crystals are not pure boron, but contain 
aluminium and carbon. In their lustre, refraction of light, 
and hardness, they resemble the diamond. Crystalline boron 
is more stable than the amorphous; it does not oxidize upon 
ignition, and is only slightly attacked by acids. Fused with 
potassium and sodium hydrate both modifications yield bo- 
rates. 
Boron Hydride, BH3. It has recently been ascertained 
that boron, like the other metalloids, can yield a gaseous pro- 
duct with hydrogen. It results when hydrochloric acid acts 
upon magnesium boride. . The latter is obtained by heating 
boron, boric' anhydride, or boron chloride, with magnesium 
dust. A colorless gas is evolved which, besides much hydrogen, 
contains some boron hydride, and burns with a bright, green 
2B + 2KOH + 2H2Q =• 2BO.OK + 3H2. BORON TRICHLORIDE BORIC ACID. 
239 
flame, with separation of boric anhydride. When the gas is 
conducted through a red-hot tube, or on holding a cold porce- 
lain plate in the flame, boron deposits as a brown sublimate. 
It produces a black precipitate, containing silver and boron, 
when it is conducted into a solution of silver nitrate. 
Boron Trichloride, BC1S, may be prepared by heating 
boron in chlorine, or conducting a stream of the latter over 
an ignited mixture of the trioxide and carbon (see SiCl4 and 
A12C16): 
B2°3 + 3C + 3C12 = 2BCI3 + 3 CO. 
It is a colorless liquid, of specific gravity 1.35, and boiling 
at 18°. Its vapor density equals 58 (H = 1), corresponding 
to the molecular formula BCl3=ii7. The liquid fumes 
strongly in the air and decomposes with water into boric and 
hydrochloric acids: 
BC13 + 3 H20 = B(OH)3 + 3 HC1. 
The trichloride also results from the action of the penta- 
chloride of phosphorus upon the trioxide: 
B2Os + 3 PC15 = 2 BC13 + 3 POCI3. 
Boron Fluoride, BF13, is similar to silicon fluoride, and 
is produced according to the same methods, by the action of 
hydrofluoric acid upon the trioxide, or by warming a mixture 
of the trioxide and calcium fluoride with sulphuric acid: 
B203 + 3 CaFl2 + 3 FI2S04 = 3 CaS04 + 3 H20 + 2 BF1,. 
It is a colorless gas, fuming strongly in the air,, of specific 
gravity 34 (H = 1), and may be condensed to a liquid under 
strong pressure. It dissolves very readily in water (700 vol- 
umes in 1 vol.), producing Hydrogen Boro-fluoride, 
BF14H (= BF13.F1H), which remains in solution: 
The reaction is analogous to the formation of hydrofluo- 
silicic acid from silicon fluoride (see p. 161). Hydrogen boro- 
fluoride is a monobasic acid, only known in solution and in its 
salts. 
Boric Acid—H3B03 = B(OH)3—occurs free in nature and 
in salts. In some volcanic districts, especially in Tuscany, 
steam escapes from the earth (fumeroles, etc.) containing 
small quantities of it. These vapors condense in small natural 
water pools, or are conducted into walled basins. By evapo- 
ration, and concentration of the aqueous solution, boric acid 
separates; the same occurs naturally as sassolite. To prepare 
4BFI3.+ 3H20 = 3HBF14, + H3B03. 240 
INORGANIC CHEMISTRY. 
pure boric acid, precipitate a hot solution of borax with hy- 
drochloric. acid. The acid separates in colorless, shining 
scales; it dissolves in 25 .parts H20 at 140, or in 3 parts at 
ioo°. The solution shows a feeble acid reaction with litmys; 
turmeric paper, moistened with it, is colored red-brown after 
drying. On boiling the solution, boric acid escapes with the 
steam. An alcoholic solution of the acid burns with a green 
flame. These reactions afford a ready means for its detection. 
When heated to ioo°, the acid loses 1 molecule H20, and 
passes into the anhydro- or meta-acid, HB02, which at 140° 
is converted into Tetraboric acid—B407H\2. When ignited, 
boric anhydride 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. 
It is a very weak acid ; and can be expelled from its salts by 
most other acids. By fusion it removes the most acids from 
their salts, in consequence of the difficult volatility of its anhy- 
dride. 
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, produced by the 
union of 4 molecules of trihydric boric acid (compare p. 191) : 
4NaB02 -f- CO2 — B40,Na, -J- C03Na2. 
4B(0H)3-5H20 = B4OtH2. 
On heating amorphous boron in a stream of N or ammonia, 
Or by igniting a mixture of trioxide and carbon in nitrogen 
gas, there is formed boron nitride, BN. This is a white amor- 
phous powder, which gives forth an extremely intense greenish- 
white light when heated in a flame. Boric acid and ammonia 
result'when steam at 200° is conducted over the nitride: 
BN' + 3H2O = B(OH)s + 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. We have repeatedly directed attention to the remark- PERIODIC SYSTEM. 
able relations of the elements of a single group, as well as to 
those of the various groups-to each other, but they appear 
more manifest if viewed in the connection in which they pre- 
sent themselves in the periodic system of elements. The position 
which these elements occupy in this system determines their 
entire physical and chemical character to a marked degree. 
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 rela- 
tions existing between the atomic weights of analogous elements, 
but only recently has the law of periodicity underlying them 
been announced by Mendelejeff and Lothar Meyer, and, accord- 
ing to this, the properties of the elements and their compounds 
present themselves as periodic functions of the atomic weights. 
Arranging thd 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 "O FI 
2. Na Mg A1 Si P • S Cl 
3. K Ca Sc Ti V Cr Mn 
4. Rb Sr Y Zr Nb Mo — 
j 5. CsBa(LaCeDi)— -— 
\ 6. Yb— Ta W — 
7. Th — Ur— 
Fe Co Ni 
Ru Rh Pd 
Os Ir Pt 
Cu Zn Ga — As Se Br 
Ag Cd In Sn Sb Te I 
— — — — — 
Au Hg Tl Pb Bi 
The first two series, lithium (Li) to fluorine (FI), and sodium 
(Na) to chlorine (Cl), present two periods of seven members 
each, in which the corresponding (above and below) members 
exhibit a great, but not complete analogy. Sodium resembles 
lithium; magnesium, beryllium ; chlorine, fluorine, etc. Then 
follow two peridds, 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 probably consti- 
tute a period. In the 7th'series there are as yet but two ele- 
ments: thorium = 232 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 .Bi, etc., are so similar that they remind us of the 
homologous series of the carbon compounds (compare p. 148), 
and, therefore, can be designated as- homologous elements. It 
is only in the third great period (series 5 and 6) that the middle 
members exhibit any variations. 242 
INORGANIC CHEMISTRY. 
Now on comparing the three great periods with the two 
small ones, we discover that the first members are analogous 
to each other : K,Rb,Cs resemble Na and Li ; Ca,Sr,Ba resem- 
ble Mg and Be. Then the similarity gradually lessens, disap- 
pears apparently in the middle members, and only appears 
again 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 func- 
tion of the three great periods is therefore other than that of 
the two small periods. But in all five periods we can detect 
a gradual, regular alteration in the properties of the adjoining 
heterologous elements. This is particularly manifest in the 
measurable physical properties, all of which show a maximum 
or minimum in the middle of the periods (both of the large 
and small), as may be seen, for example, in the specific gravity 
of the elements in solid condrtion (compare further the atomic 
volumes, p. 252). 
Na Mg A1 «i P S Cl 
Specific gravity, 0.67 1.7 2.5 2.5 2.0 1.9 1.3 
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga — As Se Br 
0.86 1.6 — —5.5 6.8 7.2 7.9 8.5 8.8 8.8 7.1 5.9 — 5.6 4.629. 
These relations show themselves very clearly in a graphic 
representation, by making the atomic weights the abscissas, 
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 exhibit? itself even in chemical prop- 
erties, in the two small periods, especially in the valence of 
the elements in their compounds with hydrogen or the hydro- 
carbon groups CH3,C2H-, etc. (compare p. 173 and p. 246). 
The hydrogen valence rises and falls periodically with the 
condensation of the substance (corresponding to the specific 
gravity: , • 
I II III IV ' III II I 
NaR MgR2 A1R3 SiH* PH3 SH2 C1H 
On the other hand, the maximum valence of the elements 
increases successively in the salt-forming oxides : 
I II III IV V VI VII 
Na20 MgO A1203 Si02 P205 S03 C1207 
The chemical valence expresses itself somewhat differently 
in the three great periods. In them we have a double period- 
icity ; thus, e.g.y with the salt-forming oxides: PERIODIC SYSTEM. 
243 
I II III IV V VI VII 
K20, CaO, Sc203 Ti02 V205—Cr03 Mn207 
VI IV 11 
FeO.. CoO, NiO 
I II III IV V VI VII 
Cu20 ZnO Ga2 03 — As205 Se03 Br207. 
In consequence of this double periodicity, the first seven 
and the last seven members of the three great periods, as re- 
gards their valence (and consequently also their compounds), 
resemble the seven members of the two small periods. To 
bring out this double periodicity and analogy, the first seven 
arid last seven members of the great periods are divided into 
two series, and arranged under the corresponding seven mem- 
bers 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. 
241) come to stand apart, as they have no analogues. In this 
manner arises the following table, in which the seven (or ten) 
vertical columns include analogous elements : 
Li 
Be 
B 
C 
N 
O 
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 
Sn 
Sb 
Te 
I 
Cs 
Ba 
La 
(CeDi) 
— 
— • 
— 
— 
— 
Yb 
— 
Ta 
W 
— 
Os Ir Pt 
Au 
Hg 
T1 
Pb 
Bi 
— 
— 
In the adjoining table (p. 244) we have presented the same 
grouping of the elements, together with their atomic weights, 
given in round numbers. 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. 241), and that the decomposi- 
tion of the latter into two series each only corresponds to the 
secondary, double analogy with the small periods. It may 
be further remarked that, in the second small period the last 
three members, P, S and Cl, show a complete homology with 
the corresponding members of the large periods—as is repre- 
sented in the "table. 244 
INORGANIC CHEMISTRY. 
I 
Group. 
II' 
Group. 
Ill 
Group. 
• IV 
Group. 
V 
Group. 
VI 
Group. 
VII 
Group. 
VIII 
Group. 
H-Compounds. 
Highest salt- 
forming oxides. 
m2o 
MO 
M2°3 
MH, 
MOj 
mh3- 
■m205 
mh2 
mo3 
MH 
M.O, 
•MOs 
mo2 
(M2H) 
MO 
Periods. Series. 
ISt ISt 
H i * 
Li 7, 
Be 9 
B 11 
C 12 
N 14 
0 16 
FI 19 
2d 2d 
• Na 23 
Mg 24 
A1 27 
Si 28 
P 31 
S 32 
Cl 35-4. 
' { & 
K39 
Cu 63 
Ca 40 
Zn 65 
Sc 44 
Ga 70 
Ti 48 
— 73 
V51 
As 75 
Cr 52 
Se 79 
Mn 55 
Br 79.7 
Fe 56 
Co 58 
Ni 58 
4* {iS 
Rb 85 
Ag 108 
Sr 87 
Cd 112 
Y 89 
In 113 
Zr 90 
Sn 117 
Nb 94 
Sb 120 
Mo 96 
Te 126 
— 100 
I 126.5 
Ru 104 
Rh 104 
Pd 106 
• . • f 7th 
.5th 9th 
[ loth 
Cs 133 
Ba 137 
I-a 139 
(Ce 140 
Di142) 
— 
— 
— 
— ' 
— 
Au 197 
Hg 200 
Yb 173 
T1 204 
Pb 206 
Ta 182 
Bi 210 
W 184 
— 
Os(i95) 
Ir 193 
Pt 195 
— 
— 
— 
Th 232 
— 
Ur 240 
— 
The Periodic System of the .Elements. PERIODIC SYSTEM, 
245 
When the periodic grouping of the elements was first presented, some 
atomic weights, not sufficiently well established at that time, had to be 
more or less altered. Thus, the atomic weight of indium, formerly 75.6, 
was made 113.4, and that of uranium 240 (before 120). All such altera- 
tions have been proved to be established through recent investigations. 
Further, the atomic weight of tellurium (formerly determined to be 128) 
had to.be less than that of iodine (126.5) > this, also, has been established 
by recent researches, which place it at 126.8 (p. in.) There is, there- 
fore, no doubt that the atomic weight of osmium (found 195) will also 
prove to be somewhat less. This is only the. more likely, because it has 
been shown that the atomic weight of iridium, which was formerly given 
as 197,'is really 192.7. Hence, the periodic system offers a control for 
the numbers of the atomic weight, while formerly they appeared to be ir- 
regular, and, at the same time, accidental. 
Further, upon the basis of the periodic system, the existence of new, not 
yet known, elements may be ascertained, which correspond to yet unoc- 
cupied, free places or gaps in the table. In fact, two such gaps have been 
filled up by the discovery of gallium (Ga =69.8) and scandium (Sc = 44 ; 
their properties have shown themselves to be perfectly accordant with 
those deduced from the periodic system. At present, only two of the ele- 
ments of the first four periods are wanting (see p. 241) : the first homologue 
of manganese (with atomic weight of about 100), and the lowest homo- 
logue of tin (atomic weight about 73). The series 5 and 6 are very in- 
complete ; the elements, terbium, samarium and erbium, little investigated 
as yet, will probably find positions in them. It may be, however, that 
the two series will together form a. single period of somewhat varying 
character. 
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 here con- 
fine ourselves to a notice of some general relations, and the 
connection of atomic weight with the chemical valence of the 
elements and the thermo-chemical phenomena. 
The relation of metalloids to metals is shown with great 
clearness in the periodic system. The first members of all 
•periods (on the left side) consist of electro-positive metals, form- 
ing the strongest bases, the alkalies—Cs, Rb, K, Na, Li, and 
metals of the alkaline earths—Ba, Sr, Ca, Mg, and Be. The 
basic character diminishes successively, in the following het- 
erologous 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 con- 
stantly exceeds the metalloidal. The first period comprises five 
metalloids (B, C, N, O, FI), the second only four (Si, P, S, Cl), 
the fourth and fifth periods each only three (or two) metal- 246 
INORGANIC CHEMISTRY. 
loids (As, Se, Br, and Sb, Te, I), which, at the same time, be- 
come 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 monovalent hydrocarbon groups (as CH3, C2H5, C3H7, 
etc.), which resemble hydrogen in many respects. Such me- 
tallo-organic compounds, in which the elements show the same 
valence as in the hydrogen compounds, are also produced by 
the metals adjacent to the metalloids : 
II III IV III II I 
Mg(CH3)2, A1(CH3)3, Si(CH3)t, P(CH3)„ S(CHs)2, C1CH3. 
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. * 
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 (pp. 
243 and 244) two sub-groups each, with, the seven vertical 
groups; on the left with the more positive, basic, and on the 
right with themore 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 VIII form the gradual transition from the 
latter to the former. 
The fundamental deduction necessarily resulting from the 
law of periodicity is, that the various elementary atoms must 
be aggregations or condensations of one and the same primor- 
dial substance, a necessary correlative postulate of the recog- 
nized unity of all forces. Then only can we comprehend that 
the properties of the elements are functions of the atomic 
weight. It was once believed that this primordial substance 
was hydrogen (hypothesis of Prout), because it seemed that 
the numbers representing the atomic weights were all whole 
numbers (multiples of the hydrogen atom = i). The most 
accurate determinations, made with exceeding care by Stas, 
prove that this is not correct in all instances. It is however, 
noteworthy that of the 18-20 elements whose atomic weights 
have been carefully established, ten (Li, K, Na, C, O, S, N, 
etc., — p. 26), so nearly approach whole numbers, that a 
complete coincidence is not unlikely. It is possible that these 
elements represent multiples of the hydrogen atom. We can PERIODIC SYSTEM. 
24 7 
only expect to. arrive at' the underlying law when the atomic 
numbers of a majority of the elements have been determined 
with equal accuracy. . . 
Periodicity of Chemical Valence.—Group I of the 
table comprises the monovalent metals, group II the divalent. 
In group III is the trivalent metalloid, boron, and the trivalent 
metals Al, Sc, Y, and Ga, In, Tl. In the tetravalent carbon 
group the valence arrives at its maximum; from here 
it gradually decreases with increasing atomic weight; the 
nitrogen group is trivalent, the oxygen group is divalent, 
that of the halogens monovalent. This valence is derived 
from the compounds with hydrogen and hydrocarbons 
(compare p. 246), or where such do not exist, as in the case 
of boron and many metals, from the halogen compounds: 
IV III II 1 
CH4 NH3 OH2 F1H 
1 11 hi iv hi 11 1 
LiCl BeCl2 BC13 CC14 NC13 OCI2 Fl2 
NaCl MgCl2 AfCi3 SiCl4 PC13 . SC12 Cl2. 
The elements of the first 4 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. 167). The higher valence of these elements is more 
manifest in the'more stable oxygen compounds. On bringing 
together the highest oxides of the seven groups capable of 
forming salts (salt-building oxides), we get this series : 
I II III IV V VI VII.. 
Li20 BeO B203 CO2 N205 S03 I207. 
The elements of the first four groups in their oxygen com- 
pounds exhibit consequently the same valence as in the 
compounds with hydrogen (or hydrocarbon radicals) and the 
halogens ; in the last three series, however, there is noticed 
a constant increase of valence for oxygen. 
Besides these highest oxides, remarkable for their greater 
stability, the elements of the last three groups afford lower 
oxides, returning in this manner to the hydrogen valence: 
Ill IV V 
p2o3 so2 i2o5. 
ii hi 
SC12 C1203. 
ci2o. 
PHs SH2 C1H. 248 
INORGANIC CHEMISTRY. 
The hydroxyl compounds of the elements of. the 7 groups 
are analogous to the oxides in constitution. They afford the 
following series, expressing the maximum valence (compare 
p. 171): 
I II III IV V VI VII 
Na(OH) • Mg(OH)2 Al(OH)3 Si(OH)4 P(OH)5 S(OH)6 Cl(OH)7 
The hydroxyl compounds of the elements of the first 4 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, in the succeeding groups, down to the weak acid 
hydrate, Si(OH)4. The hydrates of the last three groups are 
of acid nature, and mostly unstable or not known. By elim- 
ination of 1, 2 and 3 molecules H20, they yield the ordinary 
highest acids : 
PO(OH)3 
IV 
S02(OH)2 
QOjOH 
The non-saturated hydroxides behave in the same way: 
Phosphoric acid. 
Sulphuric acid. 
Perchloric acid. 
Ill IV V 
P(OH)3 S(OH)4 C1(0H)5 
II * III 
S(OH)2 C1(0H)s 
Cl(OH) 
Sulphurous acid, SO(OH)2 (p. 185), is derived from the 
hydrate, S(OH)4; chloric acid, C102,0H, from the hydrate, 
Cl(OH)5; and chlorous acid,’ CIO,OH, from the hydrate, 
Cl(OH)3. The hydrates, P(OH)s, S(OH)2 and ClOH, are 
very unstable, and the first two appear to pass readily into 
H2PO,OH and HSO.OH (compare p. 185). 
It has been shown already in the case of periodic, sulphuric 
and nitric acids, how the so-called hydrates with water of 
crystallization (regarded as molecular compounds) are ex- 
plained by the acceptance of the existence of such hydroxyl 
derivatives. The same may be done for many salts with water 
of crystallization. 
Thus, we see, and in the following pages will find it more 
extensively developed, that the relations of valence of the 
elements have their 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 influenced also by the nature of the 
combining elements: the hydrogen valence is constant, the PERIODIC SYSTEM. 
249 
valence to oxygen and halogens, on the contrary, varies accord- 
ing to definite rules. Valence, therefore, is a relative func- 
tion of the elements (p. 170). 
Periodicity of Thermo-Chemical Phenomena. We observed in 
the case of the elements of the chlorine and sulphur group, that, in their 
hydrogen compounds the heat liberation decreased successively with increas- 
ing atomic weight (pp. 64 and 113), while there is generally an increase 
in their oxygen derivatives (pp. 174 and 182). Similar relations exhibit them- 
selves with the halogen, oxygen, and sulphur compounds of the metals, as 
will be more fully exemplified later with the individual groups. Here it 
is sufficient to call attention to the relations in the heterologous series. 
In the formation of the hydrogen compounds of the elements of the first 
two periods, the following quantities of heat are set free according to 
present data: 
(C,H4) (N,H3) (O.H,) (F1,H) 
21.7 II-9 57-2 — 
(Si,H4) (P,H3) (S,H2) (C1,H) 
33.2 11.6 4.5 22.0* 
In the halogen compounds the heat modulus is more regular : 
(Li,Cl) (Be,Cl2) (B,C13) (C,C14) (N,C13) (0,C12) (FI,Cl) 
93.8 — 104 28.3 —38.1 —18 — 
(Na,Cl) (Mg,Cl2) (A1C13) (SiCl4) (P,C13) (S,C12) (Cl,Cl) 
97.7 151.0 160.9 157.6 75.3 — — 
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. 
However, on calculating the heat modulus, which corresponds to one 
equivalent of the elements (united with one equivalent of chlorine), we 
obtain numbers that diminish successively and correspond to the decrease 
of the basic metallic character of the elements: 
(Na,Cl) j'AlAj Si,Cl (Cl,Cl). 
97-7 75-5 53-6 394 25.1 
Perfectly similar relations are furnished by the oxides: 
(Na20) (Mg,0) (A1203) (Si,02) (P205) (S,03) (C1207). 
(100.2) 145.8 388.8 219 363 104 
Calculated upon one equivalent, the heat modulus is: 
5o.i 72.9 64.6 54-7 36-3 17-3 — 
* 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 
iooo. 250 
INORGANIC CHEMISTRY. 
That the heat disengaged in sodium oxide is less than in 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 the heat 
modulus is also seen with the heterologous elements in their compounds 
of similar type: 
(Mn, ClJ (Fe, Cl,) (Co, Cl,) (Ni, Cl2) (Cu, Cl2) (Zn, Cl2) 
111.9 82 76.4 74.5 51.6 97.2 
(Mn, O) (Fe, O) (Co, O) (Ni, O) (Cu, O) (Zn, O) 
94.7 68.2 63.4 60.8 37.1 85.4 
The following series are also noteworthy : 
(Ag.Cl) (Gd, Cl,) (In, Cl,) (Sn, Cl4) (Sb, Cl,) (Te, Cl,) 
29.3 93.2 127.2 87 
(Ag20) (Cd,0) (In2, O,) (SnfO,) (Sb2> O,) (Te, O) 
5.9 65.6 133.5 THE METALS. 
Although there is no sharp line of demarcation between 
metals and non-metals, yet these two classes of bodies form a 
distinct contradiction in their entire deportment, as may be 
plainly seen in the periodic system of elements. In physical 
respects the character of metals is determined by their ex- 
ternal appearance and by their ability to conduct heat and 
electricity; chemically, it shows itself chiefly in the basicity 
of the oxygen compounds; yet we see that with the increase 
of the number of the oxygen atoms, the basic character gradu- 
ally diminishes and becomes acidic. 
PHYSICAL PROPERTIES OF THE METALS. 
At ordinary temperatures all the metals excepting mercury 
are solid, slightly volatile bodies. They are opaque, and 
only a few, like gold, permit the passage of light to a limited 
extent when beaten into thin leaflets. In compact mass they 
exhibit metallic lustre and mostly possess a whitish-gray color; 
gold and copper are, however, brilliantly colored. In powder 
form almost all the metals are black. Most of them crystal- 
lize in the forms of the.regular system; only a few, showing 
a metalloidal character, are not regular. Thus antimony and 
bismuth crystallize in the hexagonal system, and tin is quad- 
ratic. The specific gravities of the metals vary greatly—from 
0.59 to 22.4 as seen from the following arrangement: 
Lithium, 0.59 
Potassium, 0.86 
Sodium, 0.97 
Rubidium, 1.52 
Calcium, 1.58 
Magnesium, 1.75 
Aluminium, 2.56 
Barium, 3.75 
Arsenic, 5,67 
Antimony, 6.7 
Zinc, 7.1 
Tin, 7.3 
Iron, 7.8 
Cobalt, 8.5 
Cadmium, 8.6 
Copper, 8.8 
Bismuth, 9.8 
Silver, 10.5 
Lead, 11.4 
Palladium, 11.5 
Thallium, 11.8 
Mercury, 13.59 
Gold, 19.3 
Platinum, 21.5 
Iridium, 22.4 
Osmium, 22.4 
In general the specific gravities of the metals, and also those 
of the metalloids, increase with the atomic weights; they stand 
more especially in sharp periodic dependence with reference 252 
INORGANIC CHEMISTRY, 
to the latter. The first members of all periods possess low 
specific gravities; the latter grow gradually until the middle 
of the period, when the maximum is attained, and then they 
again decrease (p. 242). 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 
(o.») — ”-9» that of potassium = 45.4; i.e., the potas- 
sium atom occupies a space 3.8 times larger than that of the 
lithium atom. The periodic alterations of the atomic vol- 
umes are set opposite to those of the specific gravities, as the 
former are obtained by the division of the atomic weights by 
the specific gravities. Therefore, the atomic volumes decrease 
gradually, commencing with the first members of the periods 
(Li, Na, K, Rb),attain a minimum in the middle of the periods, 
and then increase again up to the last members (Cl, Br, I). On 
the other hand, we find that with the homologous elements 
(the vertical series) an increase in the atomic volumes almost 
invariably occurs as the atomic weights increase. 
Since in the three large periods the alterations of the atomic 
volumes (as well as 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 
242) : 
ATOMIC VOLUMES OF THE ELEMENTS. 
Li Be B C* N Of FI 
11.9 5.7 4.1 3.6 — i7 — 
Na Mg A1 Si PS C1+ 
23.7 1.3.8 10.7 11.2 13.5 15.7 25.6 
K Ca Sc Ti V 
Cr Mn Fe Co Ni Cu 
Zn 
Ga 
— As 
Se Br 
45.425.4 9.3 
7.7 6.9 7.2 7.0 6.7 7.2 
9.1 
11.6 
— 13-2 
17.2 26.9 
Rb Sr Y Zr Nb 
Mo — Ru Rh Pd Ag 
Cd 
In 
Sn Sb 
Te I . 
56.1 34.9 — 21.7 15.0 
11.1 —8.4 8.6 9.2 10.2 12.9 15.3 
16.1 18.2 
20.3 25.6 
Cs Ba La Ce Di 
— 36.5 22 21 21 
* As diamond. 
f Liquid. ATOMIC VOLUME OF THE ELEMENTS. 
253 
Yb — Ta W — Os Ir Pt Au Hg T1 Pb Bi — — 
16.9 9.6 — 8.3 8.7 9.I 10.2 I4.7 17.1 l8.1 21.1 
Th — Ur 
21 13 
It is exceedingly noteworthy, that the elements standing at the beginning 
and end of the periods (on 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 an intimate relation exists between chemical 
energy and atomic volume. We can suppose that the specifically light 
elements, with large atomic volumes, execute larger chemical oscillations, 
hence act together more readily and energetically. The fact that greater 
quantities of heat are eliminated in energetic reactions would harmonize 
with this idea. And further, it may be deduced from this, that the expres- 
sions of chemical valence between the elements of greater but opposite 
oscillations (the alkalies and halogens) are the simplest, and that they be- 
have toward each other as monads. 
The metals whose specific gravity is less than 5 are termed 
light, the rest heavy. The former usually possess a greater 
chemical energy, therefore oxidize more easily and form 
strong basic oxides ; their compounds dissolve readily. On 
the whole, the heavy metals possess a varying deportment. 
They are less energetic, less basic, and yield insoluble oxygen 
and sulphur derivatives; their naturally occurring compounds, 
as a general thing, have metallic lustre, and are termed ores. 
Most metals are very malleable and tenacious, hence can be 
beaten into thin plates and leaves and drawn out into wires; 
gold and silver are the most malleable. A few, like antimony, 
bismuth, and tin, 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 Jlame. 
The fusing points of the most important of them are the 
following: 
Mercury, - —390 
•Rubidium, -I-380 
Potassium, 62° 
Sodium, 970 
Tin, 228° 
Bismuth, 270° 
Cadmium, 3150 
Lead, 3340 
Zinc, 423° 
Aluminium, 75°° 
Silver, 954° 
Gold, io35° 
Copper, 1054° 
Cast Iron, 1150° 
Wrought Iron, 1500° 
Palladium, 1500° 
Platinum, 1780° 
Iridium, 1950° 
A greater volatility also corresponds with the greater fusi- 
bility. Mercury boils at 360°; potassium and sodium about 254 
INORGANIC CHEMISTRY. 
440°; cadmium at 86o°; zinc towards iooo0, and the diffi- 
cultly fusible metals may also be volatilized by the galvanic 
current. 
All these physical properties bear a periodic dependence 
upon the atomic weights, as will be more plainly indicated in 
the individual groups. 
Of all physical properties of the elements, from a chemical 
standpoint, their heat capacity is the most important, as it 
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, would require 
very different amounts of heat. This is evident from the fol- 
lowing experiments. If we add to i kilogram of H20 at o° 
1 kilogram of H,20 at ioo°, the temperature of the mixture of 
2 kilograms of water is 50°. The quantity of heat necessary 
to raise 1 part, by weight, of H20, i°, is almost the same for 
all temperatures from o-ioo° ; this is designated the heat unit 
or calorie. On bringing to 1 kilogram H20 at o° 1 kilogram 
Hg at ioo°, the temperature of the water and of the mercury 
after their compensation, will equal only 3.20. Consequently, 
the mercury has cooled about 96.8° (from xoo to 3.20), and 
given off 3.2 calories. The quantities of heat, contained in 
equal parts, by weight, of water and mercury, therefore, are 
as 96.8 to 3.2, i.e., the specific heat of mercury (that of water 
being made = 1) is = 0.0332. 
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 been de- 
termined). W represents the specific heat, A the atomic 
weight, and the product, Wx the atomic heat: 
SPECIFIC HEAT—ATOMIC HEAT. SPECIFIC HEAT. 
255 
Elements. 
W 
A 
W X A 
Hydrogen* 
5,880 
I 
, 5-9 
Lithium 
...Li 
0,941 
7.6 
6.6 
Beryllium 
...Be 
0,408 
9.1 
3-8 
Boron (amorphous).... 
...B 
0,254 
IO.9 
2.8 
Graphite 5 
c 
0,174 
1 12 
2.1 
Diamond j 
OT43 
;12 
i-7 
Sodium 
0,293 
23 
6-7 
Magnesium 
...Mg 
0,245 
23.9 
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 
39-9 
6.8 
Chromium 
...Cr 
0,100 
524 
5-2 
Manganese 
...Mn 
0,122 
54-8 
6-7 
Iron 
...Fe 
o,x 12 
55-9 
6-3 
Cobalt 
...Co 
0,107 
58.6 
‘ 6.3 
Nickel 
...Ni 
0,108 
58.6 
6.4 
Copper 
...Cu 
0,093 
63.2 
5-9 
Zinc 
0,093 
64.9 
6.1 
Gallium 
...Ga 
0,079 
69.8 
5-5 
Arsenic (cryst.) 
0,082 
74-9 
6.2 
Selenium (cryst.) 
0,080 
78.9 
6.4 
Bromine (solid) 
...Br 
0,084 
79-7 
6-7 
Zirconium 
...Zr 
0,066 
90 
6.0 
Molybdenum 
...Mo ; 
0,072 
95-8 
6.9 
Ruthenium 
...Ru 
0,061 
103 
6-3 
Rhodigm 
...Rh 
0,058 
104 
6.0 
Palladium 
...Pd 
0,059 
106.2 
6-3 
Silver 
•••Ag 
0,056 
107.6 
6.0 
Cadmium 
...Cd 
0,054 
xi 1.9 
6.0 
Indium 
0,057 
1134 
6,5 
Tin 
0,054 
II7-5 
6.5 
Antimony 
...Sb 
0,052 
119.6 
6.2 
Tellurium 
...Te 
0,047 
126 
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 
142 
6-5 
Tungsten 
...W 
0,033 
184 
6.1 
Osmium 
...Os 
0,031 
195 
6.2 
Iridium 
...Ir 
0,032 
192.5 
6-3 
Platinum 
..Pt 
0,032 
194-3 
6.4 
Gold 
...Au 
0,032 
196.2 
6.4 
Mercury (solid) 
-Hg 
0,032 
200 
6.4 
Thallium 
...Tl 
0,033 
203.6 
6.8 
Lead 
...Pb 
0,051 
206.4 
6-5 
Bismuth 
..Bi 
0,030 
207 
6-5 
Thorium 
..Th 
0,027 
232 
6.4 
Uranium 
..Ur 
0,027 
240 
6.6 
* As Palladium hydride. 256 
INORGANIC CHEMISTRY. 
From the table, it is evident that the atomic heats of most 
of the elements lie between 5.0 and 6.8, and equal, upon an 
average, 6.4. 
It is only in the case of a few elements that the atomic 
heat is somewhat less (S, P, Si, Al), or considerably less (C, 
B, Be), than the mean. They are such as have low atomic 
weight, a metalloidal character, and occupy the middle of the 
two small periods. These variations bear distinct periodic 
dependence to the atomic weights: 
Li 
Be 
B 
C 
N 
0 
FI 
6.6 
3-8 
2.8 
1.9 
— 
— 
— 
Na 
Mg 
A1 
Si 
P 
s 
Cl 
6.7 
5.9 
5-5 
4.6 
00 
10 
5-7 
— 
The variations from the mean 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. The influence 
of temperature is, however, more important. The figures in 
the table mostly indicate the heat capacities at medium tem- 
peratures. It was known before that these show a slight in- 
crease with the temperature, but it is only recently, that H. 
E. Weber has proved that the increase is very considerable 
for the elements, C, B and Si, which, at medium tempera- 
tures, possess a remarkably low atomic heat; that, beyond a 
definite temperature, the atomic heat becomes tolerably con- 
stant, and then almost agrees with the law of Dulong and 
Petit. According to Nilson, beryllium shows a similar deport- 
ment : 
W 
A ... 
W X A 
Diamond, graphite, above 6oo°... 
0.459 
II.97 
5-4 
Boron, above 6oo° 
0.5 
IO.Q 
5-5 
Silicon, above 200° 
0.203 
28 
5-6 
Beryllium at 2570 
0.58 
9.I 
5-2 
It is probable that there is a definite temperature for all 
elements, at which their heat capacities can be compared with 
accuracy. 
From this close agreement of the found atomic heat of the 
metals with the mean, it follows, without doubt, that there 
does occur a regularity, and we must conclude that the slight SPECIFIC HEAT. 
257 
variations, apart from the inaccuracy of the observations, are 
influenced by secondary causes. Hence, the specific heat may 
serve for the derivation of the atomic weight of the elements ; 
the atomic weight is equal to the constant quantity, 6.4 divided 
by the found specific heat: 
A = -i 
TV 
The atomic weights derived from the specific heat—the so- 
called thermal atomic weights, agree in almost all instances with 
those obtained from the vapor density of the free elements or 
their volatile compounds. Where no volatile compounds of 
an element are known, the specific heat is the only certain 
means of fixing the actual atomic weight. The equivalent 
weight 37.8 (InCl) of indium is fixed with great accuracy by 
analysis ; 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 A* — 
112,5—a number closely approaching the trebled equivalent 
weight of indium 113.4 (= 37.8 X 3). From this it follows 
that the true atomic weight of indium is 113.4 and that indium 
is trivalent (InCl3). 
In their solid compounds the elements retain the specific heat attaching 
to them in their free,solid state; hence the molecular heat is nearly equal 
to the sum of the atomic heats of the elements constituting the molecules 
—law of Neuman and H. Kopp. Hence the atomic heat of elements not 
known in solid condition may be derived from the molecular heat of their 
compounds. In this manner the following atomic heats are found: For 
nitrogen, 5.0; for chlorine, 5.9; for oxygen, 4; for fluorine, 5; for hy- 
drogen, 2.3. 
In the free gaseous state the elements usually have a slighter atomic 
heat, as seen from the following table: 
A 
w* 
A X W 
Oxygen 
15.96 
0.156 
2-5 
Hydrogen 
1 
2.405 
2.4 
Nitrogen 
H 
0.172 
2.4 
Chlorine 
35-37 
0.093 
3-3 
The law that the atoms in solid condition possess the 
same thermal’capacity (A. W= A.' IV.'), finds an interesting 
analogy and exemplification in the results derived from the 
* By constant volume. 258 
INORGANIC CHEMISTRY. 
kinetic gas theory, and in the proposition of Avogadro, that 
the molecules in gas condition, at like temperatures have sim- 
ilar a degree of motion (Tf. v. — M'. 7/.) and that the latter 
experiences like increases. The molecules are the smallest parti- 
cles for gases, and the atoms the smallest parts of the solid, 
which possesses the same heat energy. The velocity of their 
heat motion, both for the molecules and the atoms is, there- 
fore, greater the smaller their masses. 
ISOMORPHISM. 
As indicated in the preceding pages, the atomic weights of 
the elements may be derived directly from the heat capacity 
of solids, while from the gas density of the volatile compounds 
we get the molecular weights, and from the latter, indirectly, 
ascertain the atomic weights (compare p. 77). A third, 
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 Mitscherlich 
(1819), that bodies chemically similar possess the same or 
almost the same crystal form. An essential mark of isomor- 
phous bodies is their ability to crystallize together—to 
form so-called isomorphous mixtures. Conversely from the 
isomorphism of two compounds may be concluded an anal- 
ogous chemical composition, a similar number of atoms in the 
molecule. This would lead us to accept as relative atomic 
weights, those quantities of the elements which replace each 
other in isomorphous compounds. For example, the metals 
calcium, strontium, and barium do not afford volatile 
derivatives. Their atomic weights could not be deduced from 
their thermal capacity, and it was the isomorphism of many 
of their compounds with those of magnesium that determined 
the same; the quantities of these elements, replacing 24 
parts by weight of magnesium (1 atom), were accepted as the 
true atomic weights. 
In the present state of chemistry we attach but secondary 
importance to isomorphism as a method of determining atomic 
weights. The phenomena of pleomorphism, according to 
which one and the same substance frequently possesses several 
crystalline forms, teach.us that the latter are not only depend- 
ent upon the chemical molecules, but that these (according to 
yet unknown laws) may unite to more complicated crystal 
molecules. Hence, isomorphism affords a means for determ- 
ining the molecular value of solid substances. 
On the other hand we know of many cases where com- CHEMICAL PROPERTIES OF THE METALS. 
259 
pounds chemically dissimilar possess a similar isomorphous 
crystalline form. Thus, dimorphous calcium carbonate 
(CaC03), as calcite is isomorphous with sodium nitrate 
(NOsNa), while as aragonite it is isomorphous with potassium 
nitrate (KN03). Consequently, isomorphism is only to be 
applied with care in chemical conclusions. Yet it is generally 
seen that bodies chemically similar have like crystalline forms, 
especially if the similarity of the elements be taken into con- 
sideration according to groups, as expressed in the periodic 
system. 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. 
As a usual thing the metals combine, without difficulty, with 
the metalloids, and with them yield well-characterized com- 
pounds the properties of which are essentially different from 
the elements composing them. The greater the chemical dif- 
ference of two bodies (metals and non-metals, bases and 
acids) the more energetic, in general, is their tendency to unite, 
and the more different and more stable the resulting products. 
As we have seen, the analogous metalloids (the groups of chlo- 
rine, of sulphur) form derivatives with each other that are not 
very characteristic. In the same manner when fused together 
the metals 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 
transition from mechanical mixtures to the real chemical 
compounds. In both instances the constituents possess but a 
slight affinity for each other, and, therefore, unite in almost all 
proportions to the so-called undetermined compounds (see p. 
90). We, however, know that definite compounds frequently 
exist in solutions; thus in an aqueous solution of sulphuric acid 
there is present the hvdrate H2S04,2H20; in aqueous nitric 
acid the hydrate HNO,,H20. And in the solutions of the 
salts crystallizing with water of crystallization there are definite 
compounds with water (e.g., Na2S04, ioH20 ; CoC12,6H20) 
at certain temperatures. Similarly constituted combinations 260 
INORGANIC CHEMISTRY. 
appear to be present, also, in the alloys; they often separate 
in crystalline form after fusion, and represent compouncs with 
definite atomic relations. Antimony and tin form a crystal- 
line compound of the composition Sb2Sn3. The crystalline 
form is always influenced by definite chemical compounds. 
This double character of the alloys manifests itself in their 
properties. In many respects they show the average deport- 
ment 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 a pure condition, 
we can impart a greater hardness by alloying them with cop- 
per ; and the latter, again, may be rendered harder by fusion 
with zinc. The character of a chemical compound is exhib- 
ited in other properties of the alloys. Their temperature of 
fusion is generally not the mean of the metals constituting 
them, but always lies lower. An alloy of 8 parts lead, 15 Bi, 
4 Sn and 3 Cd melts at 65°, although each of the single metals 
fuses above 200°. 
Mercury is able to dissolve almost all metals, forming alloys 
known as amalgams, which can crystallize. In chemical re- 
spects hydrogen is a metal, but most of the metals do not com- 
bine with it, probably because of its volatility. Palladium, 
potassium and sodium furnish the compounds Pd2H, K2H and 
Na2H, which deport themselves as alloys, while copper yields 
a pulverulent compound (CuH). That antimony yields a 
gaseous product (SbHs), is due to 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 ; hydro- 
gen first dissolves and is then evaporated again. 
Halogen Compounds.—The metals unite directly with 
the halogens to form salt-like compounds, which are not decom- 
posed by water at ordinary temperatures, and, in general, are 
very stable ; on the other hand, the halogen compounds of the 
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, hydroxides, 
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 OXIDES AND HYDROXIDES—HYDRATES. 
261 
first, rests upon the simultaneous action of carbon arid chlo- 
rine, or bromine upon the oxides (see Chloride of Aluminium 
and Silicon). 
The following types of halogen derivatives exist and show 
the different valences of the metals : 
I II III IV V VI 
KC1 ZnCl2 InCl3 SnCl4 TaCl5 WC16. 
The higher valence of the elements is more manifest 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 are usually not decomposed 
by ignition. 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 (anhy- 
drides) and water: 
Zn(OH)2 = 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 insol- 
uble oxides and hydroxides of the heavy metals depends upon 
the precipitation of the solutions of their salts by alkaline 
bases: 
CuS04 4- 2K0H = K2S04 + Cu(OH)3. 
The different valence of the metals is most clearly seen in their oxygen 
derivatives, that form salts. We have the following eight forms or types 
of the highest salt producing oxides (see p. 247), corresponding to the 
eight groups of the periodic system of the elements: 
I II III IV . V VI VII VIII 
K2O MgO AI2O3 Sn02 Bi205 CrOs (Mri207) 0s04. 
These correspond to the hydroxides or hydrates :— 262 
INORGANIC CHEMISTRY. 
1 11 in IV v VI VII 
KOFI, Mg(OH)2 Al(OH), Sn(OFI)4 Bi(OH)5 Cr(OH)6 Mn(OII)r 
The oxides and hydrates of the first two 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( ONal3. 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, eg., 
A10,OK; from Sn,OH)4 are derived stannic acid, SnO(OH)2, and its 
salts, as SnOgKj. Finally, the oxides of the last three groups are only of 
an acid nature, and afford salts with bases. Their corresponding highest 
hydrates are very unstable or do not exist; inasmuch as they yield the 
ordinary acids (p. 248) by the elimination of one, two and three molecules 
of water: 
Bi03H 
Cr04H2 
VII 
Mn04H 
Bismuthic 
acid. 
Chromic 
acid. 
Permanganic 
acid. 
hno3 
so4h2 
cio4h. 
Like the metalloids, the metals of the last four series form lower oxides 
and hydrates (p. 247) in which they exhibit a lower valence: 
Nitric acid. 
Sulphuric acid. 
Perchloric acid. 
Sn(OH)2, 
hi 
Bi(OH)3, 
IV 
Mo(om4, 
Mn(OH), 
These lower oxides have a basic character, and it is the more pro- 
nounced the further removed they are from the limiting form. In their 
whole deportment they resemble the corresponding combination forms of 
the metals of the first three groups. 
The metals of the first two groups have higher oxygen com- 
pounds, called peroxides, eg., 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 + 2IIC1 = BaCl2 + H2Q2. 
In consequence of this reaction, it is very probable that in 
the peroxides, the oxygen atoms are arranged in a chain-like 
manner as in hydrogen peroxide: 
Na—0\ 
Na—O/ 
Ba/°\ 
\/ 
When concentrated acid acts upon them, oxygen is evolved, SALTS. 
263 
and salts of the lower oxides result; heated with hydrochloric 
acid, chlorine is generated: 
BaO, + 4HC1 = BaCIj + 2H„0 + CL. 
Ordinarily, all higher oxides which are not able to form salts and which 
evolve chlorine with hydrochloric acid are termed peroxides, e.g., PbOa, 
lead peroxide, and Mn02, manganese peroxide. However, these latter 
compounds do not possess the structure of true peroxides. Lead dfoxide, 
IV 
PbOs, is wholly analogous to tin didxide, SnO?, and is capable of combining 
with bases ; therefore we must grant in it a direct union of the two oxygen 
atoms with tetravalent lead. So manganese is probably tetravalent in 
manganese peroxide. The difference between these oxygen compounds and 
the true peroxides is shown’ by their inability to form hydrogen peroxide. 
Finally, some monovalent metals are capable of forming 
oxides containing four atoms of metal, e.g., K40, Ag40 ; 
these compounds are termed quadrant oxides or suboxides. 
Salts.—By the action of bases upon acids, salts and water 
result. 
These are also produced by the direct union of basic with acid 
oxides : Na20 -f S03 = Na2S04; and by the action of metals 
upon the acids. Hence the salts are usually viewed as acids 
in which hydrogen is replaced by metals. Upon inquiring, 
however, into the composition of salts we discQver them so 
constituted that a divalent oxygen atom connects the metal 
with the acid radical (p. 174): 
NaOH *+ N03H = NOgNa + H20. 
K—O—H 
K—O—N02 
Potassium Hydroxide 
Potassium nitrate 
H—O—N02 
Nitric acid, 
rl he salts, therefore, according as it is more practicable, can 
be regarded as acid derivatives, and also as derived from the 
basic hydroxides by replacement of hydrogen. 
As we have seen, the polybasic acids yield the primary, sec- 
ondary, tertiary, etc., salts by the replacement of one. or several 
hydrogen atoms. In the same manner primary, secondary, etc., 
salts are derived from polyvalent metals (or the polyacid, 
polyhydric bases) : 
hi roH 
Bi \ OH 
(no, 
m/OH 
Bi—NOs 
\no3 
ni r no3 
Bi \ N03 
i no3 
Primary bismuth 
nitrate. 
- Secondary 
bismuth nitrate. 
Tertiary 
bismuth nitrate. 264 
INORGANIC CHEMISTRY. 
Such salts in which not all the hydroxyl groups of the poly- 
acid hydroxide are replaced by acid residues are called basic: 
Pb/0H 
p l no3 
v / OH 
Zn\Cl 
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 anhy- 
dro-acids; similarly, the polyhydric bases form polyhydrates: 
/OH 
Cu/ 
>° 
Cu< 
X)H 
.OH 
Pb< 
P> 
pb)° 
xOH 
from which basic salts are obtained (see copper and lead) by 
replacement of hydroxides by acid residues. 
By the replacement of the hydrogen atoms in the polyhydric 
acids or bases by various radicals we get the so-called mixed 
or double salts: 
/K 
so4/ 
Cu 
so4/ 
/K 
SO/ 
\ 
A1 
so* 
fK 
POJ nh4 
- (H 
Pot. am. phosphate. 
Pot. copper sulphate. 
Pot. aluminfum sulphate. 
/C1 
Pb< 
\co3 
Pb\Cl 
f NOs 
CrJ NOs 
(Clt 
The halogen double salts are usually viewed as molecular 
compounds: 
MgCl2.KCl AuC18.KC1 PtCl4.2KCl. 
If, however, the fluorides of boron and silicon, BF13.KF1, 
SiFl4.2KFI, are derived from peculiarly constituted atomic 
acids, HBF14, H2SiFl6, then a peculiar union of atoms maybe 
regarded as existing in the metallic double chlorides, which 
are often very-similar and isomorphous. ACTION OF METALS UPON SALTS AND ACIDS. 
265 
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 free condition (providing in the 
moment of its formation it does not act upon the acid : 
ACTION OF METALS UPON SALTS AND ACIDS. 
Zn + S04H2 = ZnS04 + H2. 
The metals deport themselves in the same manner with the 
salts. Zinc introduced into a solution of copper sulphate is 
dissolved to sulphate and metallic-copper deposits : 
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 similarity 
of salts and acids shows itself, too, in their acidity. All sol- 
uble 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 calcium, show 
a neutral or basic reaction—providing the base, is stronger 
than the acid. • . 6 
The displacement of metals from their salts by others, was 
formerly regarded as exclusively influenced by their electrical 
deportment. Indeed the more electro-positive, basic metals 
replace the electro-negative, less basic. In the following 
series each metal throws out from solution those preceding it: 
Au, Pt, Ag, Hg, Cu, Pb, Sn (Fe, Zn). Iron and zinc pre- 
cipitate almost all the heavy metals from solutions of their 
salts. The strofigly positive potassium is able to displace all 
other metals. This is very evident from the action of 
molten potassium upon the haloid salts—a reaction which 
frequently serves for the separation of the metals in a free 
condition : 
Zn CuS04 = ZnS04 -(- Cu. 
AlCla + 3K = A1 -f 3KCI. 
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 its vola- 
tility , in fact, we know that hydrogen, under pressure, is ca- 
pable 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 266 
INORGANIC CHEMISTRY. 
oxygen figured as the most negative and potassium as the most positive 
member, () ...... -f- K. 1 he opinion prevailed that the chemical affinity 
of the elements depended upon their electrical differences, and that chemi- 
cal union occurred because the opposite electricities'united—electro-chem- 
ical 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 the affinity, in general, corresponds to 
the electrical difference, yet this does not always occur. Thus, the 
strongly negative chlorine expels bromine and iodine from their hydrogen," 
and nearly all their metallic compounds; however, chlorine and bromine 
are conversely displaced by iodine from their oxygen compounds (C103H 
and C104H) (p. 179). 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 well as 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 33,57° calories greater than that of silver sulphate. Sulphuric 
acid dissolves most metals with liberation of hydrogen, because their heat 
of formation, 
(S,04,H2) = 192,900 (S,04,II2,Aq.) = 210,760 
is less than that of the most sulphates. The heat of formation of lead 
sulphate (Pb,S,04) equals 213,500; therefore, lead would be dissolved by 
dilute sulphuric acid, did not the insolubility of lead sulphate in the-dilute 
acid prevent it from so doing. Concentrated sulphuric acid on the other 
hand does dissolve lead, because lead sulphate is soluble in it. For the same 
reason, potassium displaces almost all the other metals; on the other 
hand, potassium is separated by sodium amalgam, with formation of po- 
tassium amalgam, as the heat of formation of the latter is much greater 
than that of sodium amalgam, and therefore, in the equation, 
the heat modulus upon the right side overbalances. (Berthelot.) . 
Although the affinity relations dependent upon the quantity of heat 
frequently correspond with the electrical differences of the free elements, 
this is so influenced that the electro-motive energy is induced by the heat, 
and is proportional to the same (see p. 269). The heat of formation of 
the compounds constitutes the primary cause of their chemical transpo- 
sition; it varies in the different compounds for the same element, and 
thus explains the opposing deportment of the elements. Chlorine dis- 
places iodine in iodides, not because it is more strongly electro-negative, 
but because the heat of formation of the chlorides is greater than that'of 
the iodides. Conversely, chlprine is eliminated from chloric acid by iodine, 
because the heat of formation of the iodic acid is the greater (compare 
p. 182). H2S and I, are similarly transposed, in the presence of water, 
into HI and sulphur, while iodine is separated from concentrated hydriodic 
acid by boiling with sulphur (p. 65). 
(K,CI) + (Na,Hg) = (Na,Cl) + (K,Hg) ELECTROLYSIS OF SALTS. 
2 67 
ELECTROLYSIS OF SALTS. 
On subjecting a salt in a fused or dissolved condition to the 
action of an electric current, it is decomposed, so that the 
metal separates upon the negative pole and the acid group or 
halogen upon the positive : 
+ — 
NaCl Na + Cl. 
The oxygen salts behave in the same way; the metal upon the 
negative pole, the acid residue upon the positive: 
4- _ 
CuS04 = Cu + S04. 
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 again 
forms the acid: 
so4 + h2o so4h2 + o. 
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 the decomposition, indeed, occurs in the 
manner indicated is confirmed by the fact that free acid arises 
at the positive pole. 
All neutral salts are similarly decomposed. If, however, 
the metal contained in the salt acts upon water when free, 
manifestly a secondary reaction must occur at the negative 
pole. The real electrolytic decomposition of potassium sulphate 
would then take place according to the following equation : 
+ — 
so4K2 = k2 + (SO3 + O). 
The separated potassium decomposes the water with forma- 
tion of potassium hydrate and the disengagement of hydro- 
gen : . ■ 
K + HOH = KOH '+ H. 
Therefore, hydrogen and potassium hydrate occur as definite 
decomposition products, at the negative pole; at the positive, 
however, we have oxygen and sulphuric acid. On coloring 
the liquid exposed to the electrolysis with a little violet syrup, 
that part at the -f- pole will become red, owing to the acid 
formed, while that at the — pole will have a green color from 
the base. That the electrolytic decomposition of potassium 
sulphate and similar salts proceeds in the manner given, may 268 
INORGANIC CHEMISTRY. 
be proved experimentally by using mercury as negative elec- 
trode ; then the separated potassium will combine with the 
mercury and form an amalgam, which will act gradually upon 
the water. • 
It was formerly believed that the alkali salts were directly decomposed 
by electrolysis into metallic and acid oxides, which yielded the hydrates 
(KOH and S04H2) with water; the appearance of H and O was attributed to 
the simultaneous electrolytic decomposition of water (a view which was set 
aside by the behavior of the other salts). With this erroneous idea as a 
basis, all salts were held to be binary compounds of the metallic oxides 
(bases) with acid oxides (acids), e.g., K20.S03 = K2S04, K20.N205 = 
2KNOs—dualistic theory of Berzelius. The acids and bases were also 
thought to be binary compounds of a metallic oxide or acid anhydride 
with water: 
The acid oxides or anhydrides were termed acids and the true acids 
hydrates. 
K20.H20 = 2KOH, so3.h2o = H2SOt. 
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 dis- 
engagement), while at the positive pole water and oxygen ap- 
pear—produced by decomposition-of the hydrogen peroxide 
formed at first: 
(0H)2 = H20 + o. 
It is, therefore, probable that the water is also decomposed 
in an analogous manner : 
2H0H = H2 + 02H2; 
the .peroxide produced at first breaks up, however, for the 
most part, into water and oxygen. 
Considering the quantities which are deposited from various 
compounds, by the same electric current, we will discover that 
a like number of valences are invariably dissolved in like time, i.e., 
equivalent quantities are separated according to the idea of the 
valence theory (p. i68.)(The law of Faraday and Becquerel). 
Thus in the simultaneous decomposition of hydrochloric acid, 
water and ammonia (pp. 71,96,127),equal volumes of hydrogen 
(= 1 part) are liberated, while at the positive pole 1 volume 
of chlorine (=35.5 parts), fz volume of oxygen (= 7.98 
parts) and volume of nitrogen (= 4.67 parts) appear. The 
quantities decomposed by electrolysis, therefore, bear the ratio: 
H„0 H,N 
HQ, -J—. 
> . 2 3 ELECTROLYSIS OF SALTS. 
269 
In the same way, equal quantities of chlorine are set free from 
all metallic chlorides (and other salts, as the chlorine atoms 
are alike in all), while the quantities of the precipitated metals 
agree with their chemical activities. The quantities of the dif- 
ferent salts, decomposed by electrolysis, stand in the following 
relation : 
AgNOs, CllC1»> SbClj Fe2Clg SnCl4 HgCl2 Hg2(NQ3)2 
2 2 3 6 4*2' 2" 
Therefore, 31.6 parts Cu are deposited for the 35.37 parts 
Cl in cupric chloride (Cu"Cl2), but from cuprous chloride 
(Cu'Cl) we obtain 63.2 parts Cu; fronq mercuric chloride 
■( big, Cl2) we obtain 99.9 parts Hg, and from mercurous nitrate 
(Hg'NOs) 199.8 parts Hg, etc. The quantities of the metals 
existing in the different states of oxidation, and which are 
equivalent to each other, vary and correspond to their chemical 
affinity. Tike valences are dissolved in equal periods of time. 
„ As the quantity of heat liberated in the union of like valences (in KC1, 
CuCl, AgN03, etc.), and that necessary for the decomposition are very 
different, and since Faraday’s law calls for the solution of like valences in 
equal periods by the same current—the performance of a different amount 
of work—there must occur, in consequence of the law of the conservation 
of energy (if the electric current does indeed do the electrolytic work), an 
unequal distribution of the energy of the current upon the different electro- 
lytes. The manner of this distribution is not known. Faraday thought 
it probably took place in such a manner that the greater consumption of 
energy by the electrolytes was compensated by their resistance, which 
consequently, is so much less. 
The relation of the electro-motive force of a galvanic cell to the chemical 
transposition occurring within it is equally obscure. Disregarding the elec- 
tric contact theory, the source of the former (with the principle of the 
conservation of energy as basis i can only be found in the loss of chemical 
energy (heat disengagement), which corresponds to the change taking place 
within the galvanic element. Following the experiments of Joule and 
others, W. Thomson assumed that the energy, developed by a cell, was 
proportional or equal to (providing no secondary actions occurred)-the 
heat modulus of the chemical reaction producing it. Thus, the effective- 
ness of a Daniell element (combination of zinc and dilute sulphuric, acid 
or zinc sulphate and copper in a solution of copper sulphate) depended 
on the replacement of Cit in CuS04, replaced by Zn—a reaction in which 
50.1 calories are developed. Further, the action of a Bunsen cell (zinc 
and carbon in a solution of potassium bichromate arid sulphuric acid) was 
supposed to be due to the formation of zinc and chromium sulphates, in 
which instance 99.8 calories are set free. In all such constant batteries J. 
Thomson and others concluded that the electro-motive force is equal or 
proportional to the energy developed in the chemical reaction. Accord- 
ing to a later theory of Helmholtz, and the determinations made by F. 
Braun, chemical energy cannot be completely, but only partly, trans- 
formed into electric energy, just as heat cannot be perfectly changed to 270 
INORGANIC CHEMISTRY. 
mechanical work. Hence, the electro-motive force of an element is 
usually less, but should be greater than the heat energy corresponding to 
the chemical transposition. In the latter cases, we must renounce any 
relation of the electro-motive force to the heat modulus, in case these, as 
is probably true, cannot be explained by secondary chemical reactions which 
have gone unconsidered heretofore. 
When two salts in solution or fusion come together, a chem- 
ical action will frequently occur. Berthollet endeavored (close 
of preceding century) to explain the resulting phenomena by 
referring them to pure physical causes, and excluded every 
special chemical affinity. 
In the opinion of Berthollet, four salts always arise in the 
solution of two. For example, on mixing solutions of copper 
sulphate and sodium chloride, there exist in solution copper 
sulphate, sodium sulphate, copper chloride, and sodium chlo- 
ride : 
2C11SO4 + yield 
CuS04 -f- Na2S04 + CuCl2 -f 2NaCl. 
That copper chloride is really present in the solution 
together with the sulphate, follows, from the fact that the blue 
color of the latter acquires a greenish color, peculiar to the 
copper chloride, by the addition of sodium chloride; other 
phenomena are not noticeable at first. Suppose one of the 
four salts formed in the solution is insoluble or volatile, the 
reaction will occur somewhat differently. Upon adding 
barium chloride to the copper sulphate solution four salts will 
be formed at the beginning just as in the first case. The 
barium sulphate produced separates, however, in consequence 
of its insolubility, the equilibrium of the four salts will be 
disturbed, and new quantities of CuS04 and BaCl2 act upon 
each other until the transposition is complete: 
CuS04 + BaCl2 = BaS04 + CuCla. 
The chemical transposition may, therefore, be explained by 
the insolubility of the barium sulphate. On adding HC1, or 
soluble chlorides to the solution of a sdver salt all the silver 
is precipitated as chloride, because 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 HN()3, are present in the solution. 
Upon warming the latter volatile nitric acid will evaporate, 
and, in proportion to its separation, new quantities of potas- TRANSPOSITION OF SALTS. 
271 
sium nitrate afid hydrogen sulphate will act upon each other 
until the transposition is complete: 
2KNOs + H2S04 = K,S04 + 2HNO3. 
The decomposition of potassium nitrate by sulphuric acid, 
therefore, is the consequence of the volatility of the nitric 
acid. Sulphuric acid decomposes sodium thloride in the 
cold, because hydrogen chloride is volatile. Carbonates are 
even decomposed by very weak acids, because the carbonic 
acid, H2C03, at once separates gaseous carbon dioxide, C02. 
In many instances the chemical transpositions may’be ex- 
plained by such physical causes, and there-is no doubt that an 
important role attaches to them. It is, however, not justifi- 
able to ignore any special chemical affinity between the various 
substances, as did Berthollet. 1 he reactions are determined 
by the chemical affinity and are independent of all physical 
causes. This is seen 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 iron acetate. 
Although an insoluble salt is not produced here, yet the re- 
arrangement of the two salts, evident from the optical proper- 
ties of the solution, is a perfect one; only iron acetate and 
potassium chloride are present in the solution: 
Fe2Cl6 + 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 can 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 four salts are present in the previously 
mentioned solution of copper sulphate and sodium chloride, 
the quantities of copper chloride and sodium sulphate are’ 
however, much greater than those of copper sulphate and 
sodium chloride (proved by the optical properties of the solu- 
tion), because the- affinity of sulphuric acid for sodium is 
greater than the same for copper. 
The phenomena recorded above are the subject of contro- 
versy and many investigations at the present time. Generally 
the chemical transpositions of salts with salts or with acids 
(hydrogen salts) and bases, are determined and governed by the 
tendency (bestreben) toward the greatest heat disengagement. 272 
INORGANIC CHEMISTRY. 
This is true, too, of other chemical reactions (the precipita- 
tion of metals from their salts by other metals, the action of 
metals upon acids and water, the alternating deportment of 
metalloids, etc.). In most cases the reaction corresponding 
to a chemical equation is easier and more complete, the more 
the sum of the heats of formation of the resulting bodies 
exceeds that of those reacting (the chemical energy of the 
first is less than that of the latter). The reverse of the reac- 
tion can only succeed by the consumption of energy and 
requires addition of heat or electricity. This is of practical 
importance (see Magnesium Chloride) in the review of the 
details of chemical reactions. In the action of salts and 
acids these relations are, however, complicated because many 
reactions, especially in aqueous solution, occur accompanied 
by a direct absorption of heat. The aim of thermo-chem- 
istry, in accord with the efforts of Berthelot, is to classify 
such reactions under the principle of the greatest heat disen- 
gagement, and declare its exceptions due to the influence of 
secondary causes. This would bring into-consideration the 
formation of acid and double salts, the influence of the heat of 
solution, the decomposition of salts by solvents, etc., etc. In 
this manner very many of the apparent exceptions have been 
satisfactorily accounted for. It would, however, seem that 
the principle of the greatest heat disengagement is constantly 
conditioned by another, that besides the heat modulus the 
stability of the bodies must be regularly taken into considera- 
tion—in such a manner that the formation of the more stable 
bodies is given the preference even if they possess a less heat 
of formation. The more stable the compounds, the lower the 
temperature, the better do the reactions conform with the 
principle of Berthelot; the disagreements become more evi- 
dent when the compounds approach the state of dissociation 
(by heat or solution). In the latter case considerable influence 
falls to the mass of the reacting bodies, as has already been 
formulated in the theory of mass action by Berthelot and 
Saint Gilles, and in the affinity theory of Guldberg-Waage. 
The principle of the greatest heat disengagement, therefore, 
constitutes a special case of a more general principle, which 
Braun proposes to designate as the principle of the greatest 
work-power. ALKALI METALS. 
273 
GROUP OF THE ALKALI METALS. 
Potassium, 39-°3 
Rubidium, 85.2 
Caesium, 132.7 
Lithium, 7.- 
Sodium, 22.99 
(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 chaiacter of 
both groups is seen, too, in their monovalence : in their com- 
binations with each other, they saturate their affinity by single 
atoms. The more distinct the chemical character of two 
elements and the more unlike they are, the simpler and the 
more definite will the expressions of valence in general be 
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 formation of-strong 
basic hydroxides, which dissolve readily in water and are 
called 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. 252), sodium is • more energetic 
than lithium, potassium more than sodium, and rubidium 
more than potassium. Caesium is not known in a free con- 
dition, but, judging from its compounds, it possesses a more 
basic character than rubidium. We saw in other analogous 
groups (of chlorine, oxygen, phosphorus, carbon), that the 
metalloidal, negative character diminishes, and the basic in- 
creases with the increasing atomic weight. 
The atomic weights increase simultaneously with the specific 
gravities ; but as the increase of the former is greater than 
that of the latter, the atomic volumes (the quotients — — 
p.252), are always the greater. The increasing fusibility and 
volatility correspond to the increase of the atomic volumes ; 
rubidium distils at a red heat, while lithium only volatilizes 
■witfi difficulty: 
Li 
Na 
K 
Rb 
Cs 
Atomic weight’. 
7 
23 
39 
85 
132 
Specific gravity 
°-59 
O.97 
0.86 
I.52 
(2-4) 
Atomic volume 
11 -9 
23-7 
45-4 
56.1 
Fusion temperature 
i8o° 
95.6° 
62 50 
38-5° 274 
INORGANIC CHEMISTRY. 
Although all the alkali metals exhibit a great similarity in their 
chemical deportment, we discover more marked relations be- 
tween potassium, rubidium and caesium upon the one hand, 
and lithium and sodium on the other, which accords with their 
position in the periodic system of the elements. Especially is 
this noticed in the salts. The first three metals form diffi- 
cultly soluble tartrates and chlorplatinates (see Platinum). 
Their carbonates deliquesce in the air, while those of sodium 
and lithium are stable under similar circumstances ; the last 
is, iudeed, tolerably insoluble in water. The phosphates de- 
port themselves similarly ; lithium phosphate is very difficultly 
soluble. It must be remarked that th.e normal carbonates and 
phosphates 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 it in 
many respects, approaches the elements of the second group, 
especially magnesium, just as beryllium approaches aluminium ; 
this is indicated by the position of the elements in the table, 
p. 244. The elements of the two small periods are, indeed, 
similar, but not completely analogous, while the homology of 
the three great periods finds expression in K, Rb, Cs. 
The affinity relations of the alkalies are expressed and explained by their 
thermo-chemical relations. Generally the heat liberation is greater as the 
atomic, weights increase : thus, e.g., in the formation of the chlorides and 
hydrates (the numbers represent large calories, p. 249). 
(Li,Cl) = 93.8 (Li,Cl,Aq.) = 102.2 (Li,0,H,Aq.) = 117.4 
(Na,Cl) = 97.7 (Na,Cl,Aq.) = 91.5 (Na,0,H,Aq.) = 111.8 
(K,C1) = 105.6 (K,Cl,Aq.) = ioi.i (K,0,H,Aq.) = 116.4 
(Na2,0) = 100.2 (Na2,0,Aq.) = 155.2 (Na,0,H) = 102.0 
(K2,0.) = 97.1 (K2,0,Aq.) = 164.5 (K.O.H) , = 103.9 
The position of lithium in the periodic system explains the varying, de- 
portment of its’compounds, which frequently show a greater heat disen- 
gagement than those of sodium. Again, it is very probable that a constant 
increase in the heat modulus occurs with the true homologues of potassium 
—rubidium and caesium. 
On the basis of the principle of the greatest evolution of heat, the num- 
bers above would explain why sodium and lithium are displaced from 
their chlorides, etc., by potassium. It separates most other metals because 
the heat of formation of the potassium compounds is generally much 
greater (see p. 266). On comparing the heat of formation of water 
(H2O = 68,600 calories), we immediately perceive why it is so readily 
decomposed by the alkali metals. All metals, disengaging "more than 
68,600 calories in the formation of their oxides, Me20, or their hydroxides, 
MeOH, decompose water, and the energy will be greater, the greater the 
difference of heat. The insolubility of the oxides constitutes an obstacle POTASSIUM. 
275 
to the action ; this however, may be removed (see Aluminium) by addi- 
tion of neutral solvents. Conversely, all oxides, affording'less heat in 
their formation, are reduced by hydrogen. 
POTASSIUM. 
In nature, potassium is found principally in silicates, viz. : 
leldspar and mica. By the disintegration of these frequently 
occurrmg minerals, potassium passes into the soil, and is ab- 
sorbed by plants; the ashes of the latter consist chiefly of dif- 
ferent potassium salts. The chloride* and sulphate are also 
ound in sea water, and in large deposits in Stassfurt, at Magde- 
burg, and in Galicia, where they were left by the evaporation 
of the water of inclosed seas. Metallic potassium was first ob- 
tained by Davy, in the year 1807, by the decomposition of the 
hydroxide, by means of a-strong; galvanic current. At present 
it is prepared by igniting an intimate mixture of carbon and 
potassium carbonate: 
K = 39-03. 
K2COs -f- 2C = 2K -)- 3CO. 
Such a mixture may be made by the carbonization of organic 
potassium salts, e.g., crude tartar. It is then ignited to white 
heat, in an iron retort, and the escaping potassium vapors col- 
lected in receivers of peculiar construction, filled with rock-oil. 
The latter, an hydrocarbon, serves as the best means of pre- 
serving 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 or- 
dinary temperature it is soft, like wax, and may be easily cut. 
It crystallizes in octahedra, and has a specific gravity = 0.86. 
It melts at 62-5°C., and, when raised to a red heat, is con- 
vened into a greenish vapor. It oxidizes in the air, and becomes 
dull in color; heated, it burns with a violet flame. It decom- 
poses water energetically, with formation of potassium hydrate 
and the liberation of hydrogen. If a piece of the metal be 
thrown upon water, it will swim on the surface with a rotary 
motion ; so much heat is disengaged by the reaction that the 
generated hydrogen and the potassium inflame. Finally a 
slight explosion usually results, whereby pieces of potassium are 
tossed here and there 5 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. 276 
INORGANIC CHEMISTRY. 
On conducting hydrogen over metallic potassium heated to 300°-400°, 
potassium hydride, K2H, results. This is a metallic, shining, brittle com- 
pound, which, upon stronger heating (410°), more readily in vacuo, is again 
decomposed. Exposed to air, it ignites spontaneously. The sodium hydride, 
Na2H, obtained in the same way, does not possess this latter property. 
Potassium forms three oxygen compounds, of which only the 
following yields corresponding salts. 
Potassium Oxide—K20—results from the oxidation of 
thin pieces of metallic potassium in dry air, and by heating 
potassium peroxide with metallic potassium. 
It is a white powder, fusing at a red heat, and evaporating 
at higher temperatures. It unites with water, with evolution 
of much heat, and the formation of potassium hydroxide. 
When heated in a stream of hydrogen it yields the hydrate, 
and metallic potassium is separated : 
This peculiar behavior is explained by the heat of formation 
of KOH (103.9 C.) being greater than that of K20 (97.1) '; 
hence the reaction occurs according to the preceding equation 
and heat is disengaged. Conversely, KOH cannot, therefore, 
be decomposed by potassium with the production of K..O 
(Beketoff). 
K20 + H = • KOII + K. 
Potassium peroxide, KO2 or K204, and potassium suboxide, K40, are 
very unstable, and readily pass 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 Hydroxide, or Caustic Potash—KOH—is 
obtained by the action of potassium or its oxide upon water. 
For its preparation, potassium carbonate is decomposed .by 
calcium hydrate (slaked lime): 
K2COs + Ca(OII)2 = CaC03 -f 2KOH. 
The solution of i part potassium carbonate in 10-12 parts water is 
boiled with t part slaked lime in an iron pot, until a filtered portion does 
not effervesce when hydrochloric acid is added ; i.e:, when there is no longer 
any carbonic acid present. On standing a while, the insoluble calcium 
carbonate subsides, and the liquid becomes clear. The solution of potas- 
sium hydrate is then poured off, evaporated, the residue melted in a silver 
dish (which it does not attack), and poured into moulds. The caustic 
potash, prepared in this way, is not entirely pure, but contains potassium 
chloride.and other salts. To obtain a product that is chemically pure, 
fuse potassium nitrate with copper filings, and treat the fusion with water. POTASSIUM CHLORIDE—POTASSIUM IODIDE. 
277 
Potassium hydroxide forms a white, crystalline mass that fuses 
rather easily, and volatilizes undecomposed at a very high tem- 
perature. Exposed to the air it deliquesces, as it absorbs water 
and carbon dioxide and changes into carbonate. It is very 
soluble in alcohol, and especially in water. The solution pos- 
sesses a strong alkaline reaction, saponifies the fats, and has a 
corrosive action upon the skin and organic tissues; hence it 
cannot be filtered through paper. At low temperatures the 
hydrate KOH + 2H20 crystallizes out from concentrated so- 
lutions. 
The haloid salts of potassium are obtained by the direct 
union of the halogens with potassium, and by the saturation 
of the hydroxide 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 
exists 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 
crystallizes in vitreous cubes, of specific gravity 1.84. It melts 
at 7340, and volatilizes at a strong red heat. 100 parts water 
dissolve 30 parts of the salt at o°, and 59 parts at ioo°. • 
Potassium Bromide—KBr—is generally obtained by 
warming a solution of potassium' hydroxide with bromine, 
when the. bromate is also produced: 
6KOH +. 3Br2 == 5K.Br + ICBrO •+ 3H..O. 
The solution is evaporated to dryness, mixed with charcoal, 
and ignited, which reduces the bromate to bromide: 
KBrOs + 3C = 3CO + KBr. 
It is readily soluble in water and alcohol; forms cubes of sp. 
gr. 2.4, and melts at 699°. 
Potassium Iodide—KI—may be prepared like the pre- 
ceding. It is usually obtained according to the following 
method: Iodine and iron filings are rubbed together under 
water, and potassium carbonate added to the solution of the 
iron iodide ; this will precipitate ferrous-ferric oxide ; carbon 
dioxide escapes, and potassium iodide will be found in the 
solution. It forms large white crystals, fuses at 634°, and is 
tolerably volatile. Its specific gravity equals 2.9. At medium 278 
INORGANIC CHEMISTRY. 
temperatures it dissolves in 0.7 parts water, and 40 parts of al- 
cohol. The aqueous solution dissolves iodine in large quan- 
tity. Many metallic insoluble iodides dissolve in it without 
difficulty, forming double iodides, e.g.,Hg\v 2KI. The iodide 
is employed in medicin'e and photography. 
Potassium Fluoride—KF1, is obtained by dissolving the carbonate in 
aqueous hydrofluoric acid. It crystallizes in cubes at ordinary tempera- 
tures, with 2H20, but above 350 does not contain water of crystalliza- 
tion. It is very soluble in water. The aqueous solution attacks glass. It 
is greatly inclined to combine with other fluorides: KF1. HF1; BF13. KF1. 
On adding hydrofluosilicic acid to the solution of potassium salts, a gelati- 
nous precipitate of potassium silicofluoride is thrown down; this is very 
difficultly soluble in water. • ’ 
Potassium Cyanide—KCN.—This salt can be produced 
by saturating potassium hydrate with hydrocyanic acid, and 
by heating yellow prussiate of potash (see Iron). It forms a 
white, easily fusible mass, which deliquesces in the air. The 
solution may be 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 re- 
duction 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. • The following reaction 
occurs when chlorine gas is conducted through a hot concen- 
trated potassium hydroxide solution: 
6KOH + 3C12 = 5KC1 + KCIO3 + 3HaO. 
When the solution cools, the difficultly soluble potassium 
chlorate separates out. It is generally made, in trade, by the 
action of chlorine upon a mixture of calcium hydrate and po- 
tassium chloride.- The reaction occurs in two phases; first, 
calcium chlorate is formed : 
6Ca(OH), + 6C12 = sCaCla + Ca(C103)2 + 6H20; 
this then reacts with the potassium chloride : 
Ca(C103)2 + 2KC1 = 2KCIO3 + CaCl2. 
Potassium chlorate crystallizes from the hot solution in 
shining tables of the monoclinic system, which are difficultly 
soluble in water (too parts at ordinary temperature dissolve 6 
parts of the salt). Its taste is cooling and astringent. When POTASSIUM HYPOCHLORITE—POTASSIUM SULPHATE. 
279 
heated it melts (at 359°) giving up a portion of its oxygen, and 
changes to the Perchlorate—KC104—which on further heat- 
ing decomposes into oxygen and potassium chloride (see p. 
178). As it gives up oxygen readily, it serves as a strong oxi- 
dizing agent. With hydrochloric acid it liberates chlorine: 
Mixed with sulphur, or certain sulphides, it explodes on 
heating and when struck a sharp blow. The igniting ma- 
terial upon the so-called Swedish (parlor) matches consists 
of antimony sulphide and potassium chlorate ;• when this is 
rubbed upon the friction surface coated with red phosphorus it 
ignites. 
Potassium Hypochlorite — KCIO—is formed when 
chlorine is permitted to act upon a cold solution of potassium 
hydrate: 
KC103 + 6HC1 = KC1 4- 3H20 + 3(4. 
2K0H 4- ' Cl. = KC1 4- KC10 + H..O. 
It only exists in aqueous solutions; when the latter is evap- 
orated the salt is decomposed into chloride and chlorate: 
3CIOK = 2KC1 + C103K. 
The solution has an odor resembling that of chlorine, and 
bleaches strongly, especially upon the addition of acids. The 
bleaching solutions occurring in trade (Eau de Ja'velle) are 
prepared by the action of chlorine upon solutions of sodium 
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—KI03 
—are prepared by the action of bromine or iodine upon hot potassium 
hydroxide ; the second is also produced by the action of iodine upon po- 
tassium chlorate, when the chlorine is directly replaced (p. 182). If chlo- 
rine be passed through a hot solution of potassium iodate in potassium 
hydrate—the periodate of potassium, KI04, arises; it is difficultly soluble 
and when heated decomposes into O and KI03, which then breaks up into 
potassium iodide and oxygen. 
Besides the normal periodates, KI04, NaI04, other salts exist which are 
derived from the highest hydroxyl compound, I(OH)7, and its anhydro-de- 
rivatives (p. 180). These salts are very numerous, and are in part mono- 
periodates, IO(OH)5, and I02(0H)3 and partly polyperiodates, produced 
by the condensation of several molecules of the highest hydrates, eg., 
I203(0H)8 and I205(0H)4. 
Potassium Sulphate—K2S04—is formed in the action 
of sulphuric acid upon potassium chloride, and as a by- 
product in many technical operations. It crystallizes without 
water, in small rhombic prisms,, having a bitter, salty taste, 280 
INORGANIC CHEMISTRY. 
and dissolves in 10 parts H20 of ordinary temperature. It 
is employed principally for the preparation of potassium car- 
bonate, according to the method of Le Blanc. (See Soda.) 
The acid or prittiary salt—-HKS04—crystallizes in large 
rhombic tables, and is very readily soluble in water. It fuses 
about 20Q0, loses water, and is converted into potassium 
pyrosulphate—K,2S207 (p. 192)—which at 6oo° yields K2S04 
and S03. 
The salts of sulphurous acid—the primary, S03KH, and the secondary 
sulphites, S03K2—are produced when SO2 comes in contact with a potas- 
sium carbonate solution ; they are very soluble and crystallize with diffi- 
culty. The first salt shows an acid, the second an alkaline reaction. If 
sulphur dioxide be passed into a solution of pptassium carbonate until 
effervescence ceaseS and then cooled, the pyrosulphite—K2S2Os—corre- 
sponding to the pyrosulphite, will crystallize out. 
Potassium Nitrate, Saltpetre, KN03, does not occur 
anywhere in large quantities, but is widely distributed in the 
upper strata of the earth and is found as an efflorescence on the 
soil in some regions of the hot zone (in Egypt and East India). 
It is produced whenever nitrogenous organic substances decay 
in the presence of potassium carbonate—conditions which are 
present in almost every soil. The intentional introduction 
of these is the basis of 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 the action of the 
air for two or three years, when nitrates are produced from the 
slow oxidation of the nitrogen. The heaps are then treated 
with water and potassium carbonate added to the solution, 
which contains potassium, calcium, and magne'sium nitrates, 
to convert the last two salts into potassium nitrate: 
Ca(N03)2 + K2C03 = CaC03 + 2KN03. 
The precipitate of calcium and magnesium carbonate is 
filtered off and the solution evaporated. The procedure was 
formerly employed, universally in the manufacture of potas- 
sium nitrate. At present, however, almost all of it is ob- 
tained by the decomposition of the sodium salt, occurring 
in large deposits in Chili, by means of potassium carbonate 
or chloride : 
NaNOa + KC1 = • NaCl + KNOa, 
Warm saturated solutions of sodium nitrate, and potassium 
chloride are mixed and boiled, when sodium chloride, the POTASSIUM NITRITE—POTASSIUM CARBONATE. 
281 
least soluble in hot water, separates. On cooling the solution 
potassium nitrate, the least soluble in cold water, crystallizes 
out; sodium chloride is about equally soluble in hot and cold 
water, for which reason the portion not separated by boiling 
remains in solution. 
Potassium saltpetre crystallizes without water of crystalliza- 
tion in large six-sided rhombic prisms. It is far more soluble 
in hot than in cold water; ioo parts of water dissolve 244 
parts at ioo°, but at o° only 13 parts. It possesses a cooling 
taste, fuses at 338°, and decomposes when further heated into 
oxygen and potassium nitrite, KN02. Heated with carbon it 
yields potassium carbonate: 
4KN03 + 5C = 2K2COs + 3co2 + 2N2. 
Its principal use is in the manufacture of gunpowder. This a granular 
mixture of potassium nitrate, sulphur, and charcoal. The relative quanti- 
ties of these constituents are somewhat different in the various kinds of 
powder (sporting, blasting, and cannon). Upon an avenfge, the powder 
consists of 75 per cent. KN03, 12 per cent, sulphur, and 13 per cent, carbon, 
which closely corresponds to the atomic composition 2KN03 —|— S —(— 3C. 
When the powder burns its decomposition is approximately expressed by 
the following equation : 
2kno3 + s + 3c = k2s +■. 3co2 + N2. 
The effectiveness of the powder, therefore, depends upon the disengage- 
ment of carbon dioxide and nitrogen gas, the volume of which is almost 
1000 times as great as that of the decomposed powder. 
Potassium Nitrite, KN02, is obtained by fusing 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 acid: K:i P04, K2HP04, 
and KH2P04, meet with no practical application, they 
are readily soluble in water and crystallize poorly; therefore, 
the sodium salts are generally used. The borates, B02K 
and B407K2 + 5H O (see Borax), crystallize with difficulty. 
Potassium Carbonate—K2COs—ordinarily known as 
potashes, 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. When the 
plants are burned the organic acids are destroyed and potas- 
sium carbonate remains. The ashes are lixiviated with hot 
water, the filtrate evaporated and the residue ignited. The 
crude potashes thus obtained contain, besides the carbonate, 282 
INORGANIC CHEMISTRY. 
also chloride, sulphate, and other salts. To purify them, treat 
with a little water, which will dissolve the easily soluble car- 
bonate, leaving nearly all of the other ingredients behind. 
In this way we obtain pure potashes. This method of getting 
potashes from plant ashes was formerly pursued extensively in 
America, Hungary, and Russia; it is not much used at pres- 
ent, because potassium carbonate is, upon the one hand, 
replaced by the cheaper sodium carbonate in practice ; on the 
other hand, the immense deposits in Stassfurt and Galicia 
afford an inexhaustible supply of potassium salts. Consider- 
able quantities of potassium carbonate, used at present almost 
entirely for the production of Bohemian or crystal glass, have 
been recently obtained from Stassfurt, according to the meth- 
ods employed in the preparation of sodium carbonate from 
the chloride. (See Soda.) Chemically pure potassium car- 
bonate is obtained most conveniently by the ignition of cream 
of tartar or by heating the primary carbonate. 
The commercial carbonate is a white, deliquescent powder 
melting at 830-°, and vaporizing at a red heat. It crystallizes 
from concentrated aqueous solutions with molecules of 
water, in monoclinic prisms; at ioo° it loses 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 potassiutn carbonate is produced : 
C03K2 -f H20 + CO2 = 2KHCO3. 
This salt, ordinarily called bi-carboriate, crystallizes in mono- 
clinic prisms, free from water. It dissolves in 3-4 parts water 
and exhibits a neutral reaction. Heated to 8o°, it decom- 
poses into K2C03,C02 and water. The decomposition of the 
dry salt does not begin until about no°, while the aqueous 
solution decomposes even on evaporation. 
Potassium Silicate, water-glass, does not possess a con- 
stant composition and cannot be obtained crystallized. It is 
produced by solution of silicic acid or amorphous silicon 
dioxide in potassium hydrate, or by the fusion of silica with 
potassium hydrate or carbonate. The concentrated solution 
dries when exposed, to a glassy, afterwards opaque, mass 
which when reduced to a powder will dissolve in boiling 
water. Potassium (and also sodium) water-glass has an ex- 
tended application, especially in cotton printing, for the fixing 
of colors (stereochromy), in rendering combustible material 
fireproof, in soap boiling, etc. SULPHUR COMPOUNDS. 
283 
Potassium Hydrosulphide, KSH, is obtained when 
potassium hydrate is saturated with hydrogen sulphide: 
SULPHUR COMPOUNDS OF POTASSIUM. 
Evaporated in vacuo it crystallizes in colorless rhombohedra, 
of the formula, 2KSH -f- H20, 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 hydroxide, it has an alkaline reaction. 
On adding an equivalent quantity of potassium hydroxide to 
the sulphydrate solution, we get potassium sulphide: 
KOH + H,S = KSH + H,0. 
KSH + KOH = K2S + H20. 
Potassium Sulphide, K2S, is usually obtained by fusing 
potassium sulphate with carbon : 
K2S04 -f 2C = KjS 4- 2CCV 
When fused, it solidifies to a red crystalline mass. It crys- 
tallizes from concentrated aqueous solutions with 5 molecules 
of H2Q, in colorless prisms, which deliquesce in the air. The 
solution absorbs oxygen from the latter, and is decomposed 
into potassium hyposulphite and caustic potash : 
3k2s + H2Q + 20, = k2s2o3 + 2KOH. 
Potassium hydrosulphide and sulphide precipitate insoluble 
sulphides from the solutions of many metallic salts. They are 
decomposed by acids with liberation of hydrogen sulphide. 
When the aqueous solution of the sulphide is boiled with 
sulphur the poly sulphides, K2S3, K2S4 and K2S5, are formed, 
which after fusion solidify to yellowish-brown masses. The 
aqueous solutions of the polysulphides are decomposed by 
acids, with disengagement of H2S and separation of sulphur 
(milk of sulphur). The so-called liver of sulphur (Hepar sul- 
phuris), a liver-brown mass, used in medicine, is obtained by 
the fusion of potassium carbonate with sulphur, and consists 
of a mixture of potassium polysulphides with potassium sul- 
phate. 
The aqueous solution of the potassium, as well as that of the 
sodium sulphide, dissolves some metallic sulphides and forms 
sulpho-salts with them (p. 221). 
When dry ammonia is conducted over heated potassium, 
potassamide (NH2K) results. ' This is a dark-blue liquid which 
solidifies to a yellowish-brown mass. Water decomposes it 284 
INORGANIC. CHEMISTRY. 
into potassium hydroxide and ammonia. When potassamide 
is ignited away from the air, it loses ammonia, and leaves 
•behind potassium triamide, NK3, 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 charac- 
terization 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 potassium solutions, a yellow, crystalline precipitate of 
PtCl4. 2KCI. Potassium compounds introduced into the flame 
of an alcohol or gas lamp impart to the same a violet color- 
ation. The spectrum of the flame is characterized by two 
bright lines, one red and one violet (see Spectrum Analysis). 
RUBIDIUM AND CESIUM. 
Rubidium and Caesium are the perfect analogues of potassium (p. 274). 
They were discovered by means of the spectroscope, by Bunsen and Kirch- 
hoff, in i860. Although only occurring in small quantities they are yet 
very widely distributed, and frequently accompany potassium in mineral 
springs, salt, and plant ashes. The mineral lepidolite contains 0.5 per cent, 
of rubidium; upwards of 30 per cent, of caesium oxide is present in the 
very rare pollucite, a silicate of aluminium and caesium. The spectrum of 
rubidium is marked by two red and two violet lines; caesium by two dis- 
tinct blue lines ; hence, the names of these elements. 
Rubidium and caesium form double chlorides (PtCl4. 2RbCl) with pla- 
tinum chloride, and they are more insoluble than the double platinum salt 
of potassium, hence may answer for the separation of these elements from 
potassium. Rubidium and caesium may be obtained free by decomposing 
their fused chlorides with the electric current. Rubidium is also prepared 
by igniting its carbonate with charcoal. Metallic rubidium has a silver- 
white color, with a somewhat yellowish tinge; its vapor is greenish-blue. 
Caesium has only been obtained alloyed with mercury. 
Rb r= 85.2. Cs =132.7. 
SODIUM. 
Na = 22.9. 
Sodium is widely distributed in nature, especially as chloride 
in sea water and as rock-salt; and is also found in silicates. 
The metal was obtained in 1807, by Davy, by the action of SODIUM OXIDE—SODIUM HYDRATE. 
285 
a strong electric current upon fused sodium hydroxide. At 
present, like potassipm, it is obtained upon a large scale by 
igniting a mixture of sodium carbonate and carbon in an iron 
retort: 
Na2COs -)- 2C = 2Na 3CQ. 
The liberated sodium vapors are condensed on flat iron receiv- 
ers 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°, distils 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 potassium. 
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 somewhat different. 
It is obtained by burning sodium in a stream of oxygen. Its formula is 
Na202. When heated it absorbs iodine vapors, forming the compound, 
Na2OI2 (Na202 + I2 = Na2OI2 + O), soluble in water, but decomposed 
by acids into free iodine and sodium salt. This compound, like some 
others, seems to indicate that sodium has several valences. 
Upon heating sodium oxide it is decomposed by hydrogen, with separa- 
tion of metallic sodium and formation of sodium hydroxide : 
Na20 -f- H = NaOH -f- Na. 
This is explained by the fact that the heat of formation of NaOH is 
greater than that of Na20 (p. 274). Carbon monoxide decomposes sodium 
monoxide in a similar manner, when the latter is heated 10290-310° : 
2Na20 CO = Na2C03 Na2 
This reaction occurs because the heat of formation of sodium carbonate 
(271.2 C.) is greater than that of 2Na20(2 . 100.2 C.) and CO(30.i C.). 
Na2G combines with CO2 about 400° with production of light and yields 
C03Na.2(Na20, C02 = 74.1 C.). 
Sodium Hydroxide, Sodium Hydrate, or Caustic Soda, 
NAOH, like .potassium hydroxide, is formed by boiling a so- 
lution of sodium carbonate with calcium hydrate: 
At present it is directly produced in the soda manufacture 
by adding a little more carbon to the fusion (see Soda), or by 
Na2C03 + Ca(OH)2 = CaC03 + 2NaQH. 286 
INORGANIC CHEMISTRY. 
igniting sodium carbonate (Lowig) with ferric oxide, which 
affords a compound of Fe203,Na20, which warm water decom- 
poses into ferric oxide and sodium hydrate. 
The sodium hydroxide which solidifies after fusion is a 
white, radiating, crystalline mass, and resembles caustic potash 
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, because 
the resulting carbonate is also deliquescent). The aqueous solu- 
tion, called sodium hydrate, resembles that of potassium. 
Crystals of NaOH -f- 3*^H20 separate at o° from the con- 
centrated solution; they melt at 6°. . 
Sodium Chloride—NaCl—is abundant in nature. It is 
found almost everywhere in the earth and in natural waters; 
in sea-water it averages 2.7-3.2 per cent. As rock-salt it forms 
large deposits in many districts—at Stassfurt and Wielizca in 
Galicia. 
In warm climates, on the coasts of the Mediterranean sea, sodium chlo- 
ride 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, there- 
fore, to summer time. After sufficient concentration, pure sodium chloride 
first separates, and this is collected by itself. Later, there crystallizes a 
mixture of sodium chloride and magnesium sulphate; finally potassium 
chloride, 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 Nor- 
way and at the. White Sea, the cold of winter is employed for the produc- 
tion of salt. In the freezing of sea-water, as well as of other solutions, 
almost pure ice separates at first; the enriched sodium chloride solution 
is then concentrated in the usual way. 
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 openings already formed. When 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 them it will be dis- 
tributed and evaporated by the fall; the concentrated brine collects in the 
basin below, and is then evaporated over a free fire. 
Sodium chloride crystallizes from water in transparent 
cubes, which arrange themselves by slow cooling in hollow, 
four-sided pyramids. It melts at 772° and volatilizes at a 
white heat. It is only slightly more soluble in hot than in Glauber’s salt. 
287 
•cold water; 100 parts at o° dissolve 36 parts salt; at ioo°, 39 
parts. The saturated solution, therefore, contains 26 per cent, 
sodium chloride. The specific gravity of the crystals equals 
2.13. If the saturated solution be cooled below io°, large 
monocljnic tables (NaCl + 2H20) separate; these lose water 
at o° and become cubes. 
The ordinary sodium chloride usually contains a slight 
admixture of magnesium salts, in consequence of which it 
gradually deliquesces in the air; the perfectly pure salt is not 
hygroscopic. When heated the crystals crackle, as water that 
has been mechanically enclosed escapes. 
Sodium Bromide and iodide crystallize at ordinary temperatures with 
2 molecules of H20, which they lose again at 30°; above 30° they separate 
in anhydrous cubes. Sodium bromide fuses at 708° and the iodide at 628° ; 
the former is difficultly soluble in alcohol and the latter is very soluble. 
Sodium Chlorate (NaC10s) and perchlorate (NaC104) are consid- 
erably more soluble in water than the corresponding potassium salts. 
Sodium Iodate—NaI03—is obtained the same as the potassium salt, 
and crystallizes at ordinary temperatures with 3 molecules of H20 in silky 
needles. If chlorine gas be conducted through the warmed solution of 
sodium iodate in sodium hydrate, the periodate IO | |qpj^2 (see P- 180) 
crystallizes out on' cooling This becomes the normal salt (NaI04 -j- 
3H2O) when dissolved in nitric acid. 
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 of crystallization, in Spain. It is a by-product 
in the manufacture of sodium chloride from sea-water and 
brine. It is produced in large quantities by heating salt with 
sulphuric acid : 
and is used in making soda (sodium carbonate). Or it may 
be prepared by the method of Hargreaves, by conducting 
S02, air and steam over strongly ignited sodium chloride : 
2NaCl + H2S04 = Na2S04 + 2HCI, 
2NaCl -+-• S02 + O + H20 = S04Na2 -f 2HCI. 
In modern times the sulphate has been obtained by a transpo- 
sition of sodium chloride with magnesium sulphate at a winter 
temperature—a procedure which is prosecuted chiefly in Stass- 
furt, where immense quantities of magnesium sulphate exist: 
aNaCl + S04Mg = MgCl2 -f- S04Na2. 
Sodium sulphate crystallizes at ordinary temperatures with 
io molecules of H20, in large, colorless, monoclinic prisms, 288 
INORGANIC CHEMISTRY. 
which crumble in the air and fall into a white powder. When 
the salt is heated to 330, it fuses in its own water of crystal- 
lization; by further increase of temperature it gradually loses 
this, becomes solid, and again fuses at a red heat. The solu- 
bility of Glauber’s salt (Na2S04 +10 H20) shows the follow- 
ing interesting deportment: 100 parts of water dissolve, at o°, 
12 parts; at 180, 48 parts; at 25°, 100 parts; at 30°, 200 
parts; at 330, 327 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 ioo°, 238 parts of the salt. While, ordi- 
narily, the solubility increases with temperature, Glauber’s, 
salt exhibits a varying deportment. This is explained in that 
the hydrate, Na2S04 + ioH20, in aqueous solution, above the 
temperature of 330, decomposes into water and the salt, Na2S04 
-|- H20, which is less soluble in water. The decomposition 
does not occur at once, but only gradually, with increasing 
temperature, 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 (p. 190). The solution, 
saturated at 330, becomes turbid upon heating, and a portion 
of the dissolved 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 330, is allowed to 
cool to the ordinary temperature, and even lower, not the slightest sepa- 
ration of crystals occurs, although the salt is vastly more insoluble at lower 
temperatures than at 330. 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 crystallization setting in. If, however, a glass rod, or some other 
solid body, be introduced into the solution, it will solidify suddenly to a 
crystalline mass. The particles of dust floating about in the air will have 
a like effect; therefore, to preserve the supersaturated solution, the vessel 
containing it should be kept well corked. By accurately made investiga- 
tions, it has been determined that the crystallization of the supersaturated 
Glauber’s salt solution is only induced by contact with already formed crys- 
tals. These must then be present everywhere in the atmosphere, because 
only solids which have been exposed to the air, and have not been care- 
fully cleansed afterward, bring about the crystallization. Hence, the for- 
mation of a crystal of Glauber’s salt is always dependent upon the pre- 
vious existence of a similar crystal—just as the production of cells is only 
caused by cells. 
In the crystallization of a supersaturated Glauber’s salt solution, consid- 
erable 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 io°, occasionally, and of their own accord, trans- 
parent crystals, Na2So4 -)- 7H.2O, separate from the supersaturated solu- SODIUM HYPOSULPHITE. 
289 
tions. 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, and finds 
extended application in the manufacture of glass and the prep- 
aration 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 : 
NaCl + H2S04 = NaHS04 + HC1. 
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 sul- 
phur dioxide into solutions of sodium hydroxide or carbonate. The sec- 
ondary sulphite, crystallizes with 7 molecules of H20 at ordinary 
temperatures; in the presence of sodium hydroxide, or by warming the 
solution, it separates in the anhydrous state. The primary sulphite— 
NaHSOs—gives up sulphur dioxide in the air, and is oxidized to sodium 
sulphate. 
Sodium Hyposulphite—Na2S203—is prepared by boiling 
the aqueous solution of neutral sodium sulphite with flowers 
of sulphur: 
Na2SOs + S = Na2S!iQ3. 
It crystallizes with five molecules of H20, in large mono- 
clinic prisms, dissolves very readily in water, and is somewhat 
deliquescent in the air. At 56°, it melts in its water of crys- 
tallization loses all water at ioo°, 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 
decompose 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 solutions. 
The hyposulphite is used as a reducing agent; chlorine, bro- 
mine and iodine are converted by it into the corresponding 
halogen salts: 
2S2OsNa2 -f- I2 = S406Na,2 -f 2NaI 
Sodium tetrathionate. 
An iodine solution is instantaneously decolorized by sodium 
hyposulphite. Chlorine behaves differently; sulphuric acid 
and sodium chloride are produced. Upon this reaction rests 
the application of sodium hyposulphite as an antichlor in 
chlorine bleaching, to remove the excess of the chlorine,- 290 
INORGANIC CHEMISTRY. 
which has a destructive action upon the fibre. In consequence 
of its property of dissolving the halogen silver derivatives, it 
is employed in photography. 
Sodium Carbonate (Soda)—Na2COs.—This, technically, 
very important salt occurs frequently in nature. In some dis- 
tricts, as in Hungary and in Africa, it disintegrates from the 
soil, and occurs also in the so-called sodium seas (in Egypt, 
and upon the coast of the Caspian Sea). It is contained in the 
ashes of many sea-plants, chiefly the algae, etc. These assimi- 
late 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, called varec in Normandy, 
kelp in England, formerly served as the principal material for 
the preparation of soda. At present it is, however, almost ex- 
clusively made in large quantities from sodium chloride, ac- 
cording to a method devised in 1808 by Leblanc. 
According to this method, the sodium choride is converted, 
by warming with sulphuric acid, into sodium sulphate (p. 287). 
When the latter is dry, it is mixed with charcoal and chalk, 
and ignited in a reverberatory furnace. Two principal phases 
may be distinguished in this reaction. First, the carbon re- 
duces the sodium sulphate : 
Na2S04 + 2C = Na2S + 2C02. 
The sodium sulphide then acts upon the calcium carbonate 
to form calcium sulphide and sodium carbonate: 
Na2S + CaC03 = CaS + Na2C03. 
At the same time, the high temperature converts a portion 
of the calcium carbonate into calcium oxide and carbon diox- 
ide, which is reduced by the ignited carbon to the monoxide ; 
the appearance of the latter, which burns with a bluish flame, 
indicates the end of the action. The chief products in the 
soda fusion are, then, sodium carbonate, calcium sulphide and 
oxide; in addition, different other sulphur salts are formed in 
smaller quantity. The fusion is lixiviated with cold water; the 
sodium carbonate dissolves, and there remains behind an in- 
soluble compound of calcium sulphide with oxide, CaO, 2CaS, 
the soda residue. By the evaporation of the solution and the 
ignition of the residue, we get the commercial or crude calcined 
soda, containing different admixtures, among them sodium 
hydroxide. The latter is formed by the action of excess of 
carbon upon sodium carbonate : 
Na2C03 + C = Na40 + 2CO. SODIUM CARBONATE. 
291 
By purposely adding more carbon to the fusion, sodium 
hydroxide may be obtained, together with the carbonate. To 
purify the crude soda it is recrystallized from water; large, 
transparent crystals, Na2C03 +10 H20, crystallized soda, sep- 
arate out; the sodium hydroxide remains dissolved. 
Considerable quantities of soda are obtained at present from cryolite, a 
compound of aluminium fluoride with sodium fluoride (A1F13, 3 NaFl), 
which occurs in great deposits in Iceland. The pulverized mineral is 
ignited w ith burned lime; insoluble calcium fluoride and a very soluble 
compound of aluminium oxide with sodium oxide, called sodium aluminate 
(see Aluminium) are produced : 
2(A1F13, 3NaFl)' + 6CaO = 6CaFl2 + A1203, 3Na20. 
The mass is treated with water and carbon dioxide conducted into the 
solution, which causes the precipitation of aluminium oxide, and sodium 
carbonate dissolves: 
A1203, 3Na20 + 3H20 + 3C02 = Al2(OH)6 + 3Na2C03. 
Latterly, a third procedure has appeared. It depends upon the double 
decomposition of a solution of sodium chloride with primary ammonium 
carbonate, by heat, under high pressure : 
NaCl + C03(NII4)H •== NaHCOs + NH4C1. 
This difficultly soluble primary sodium carbonate separates from solution, 
and is converted by heat into the secondary carbonate. The ammonium 
chloride remains dissolved, and afterwards is converted again into car- 
bonate by aid of calcium carbonate. In this way, one and the same quantity 
of ammonium carbonate will suffice for the conversion of an indefinite quan- 
tity of sodium chloride into soda. The technical difficulties which at first 
opposed the extension of this process, so simple in chemical respects, are 
now mostly removed, and we can expect that the so-called ammonia process 
for the soda manufacture will replace that of Leblanc at least partially. 
At ordinary temperatures sodium carbonate crystallizes 
with io molecules of H20,(Na2C03 + ioH20) in large mono- 
clinic crystals, which crumble upon exposure and become a 
white powder. It melts at 50° in its water of crystallization, 
and upon additional application of heat a pulverulent hydrate 
—Na2C03 + 2H20—separates, which in dry air has 1 molecule 
of H20, and at ioo° loses all of this. At 3o°-5o° rhombic 
prisms of the composition C03Na2 -f 7H20, crystallize from 
the aqueous solution. The anhydrous salt absorbs water from 
the air but does not deliquesce. It melts at a red heat and 
volatilizes somewhat at a very high temperature. 100 parts 
H20 dissolve 15 parts at o°, 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. 292 
INORGANIC CHEMISTRY. 
Primary Sodium Carbonate—ordinary Bicarbonate of 
Soda—Natrium Bicarbonicum—NaHC03—is produced by the 
action of carbon dioxide upon the hydrous secondary car- 
bonate : 
It crystallizes without water, in small monoclinic tables; 
it dissolves, however, at ordinary temperatures in io-ii parts 
water, and possesses feeble alkaline reaction. By heating and 
boiling the solution it passes into the secondary carbonate with 
disengagement of ' carb m dioxide. By rapid evaporation 
small monoclinic prisms of the so-called sodium sesquicar- 
bonate—C308Na4 + 3H20 or Na4H2(C03)3 4- 2H20, separate; 
this also deposits in the sodium seas of Hungary and Egypt. 
It is called Trona or Urao. 
Sodium Nitrate—NaNOs—Chili saltpetre, is found in 
immense deposits in Peru. It crystallizes in rhombohedra 
•very similar to cubes, hence designated cubic saltpetre. It 
fuses about 318°. In water it is somewhat more easily soluble 
than potassium saltpetre. In the air it attracts moisture, 
hence it is not adapted to the manufacture of gunpowder. In 
other respects it is'perfectly similar to potassium nitrate. It is 
largely used in the manufacture of nitric acid, and especially 
in preparing potassium saltpetre (p. 280). 
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 has a strong alkaline reaction, absorbs carbon 
dioxide from the air, and is converted into the secondary salt. 
Di-sodium Phosphate—Na2HP04—is the most stable of 
the sodium phosphates, and hence, is generally employed in 
laboratories (.Natrium phosphoricuni). It may be obtained 
by saturating phosphoric acid with sodium hydrate to feeble 
alkaline reaction, or may be prepared on a large scale by de- 
composing bone ashes (tricalcium phosphate) with an equiva- 
lent amount of sulphuric acid. It crystallizes at. ordinary 
temperatures with i2H20 in large monoclinic prisms which 
effloresce rapidly upon exposure. It separates from solutions 
with a temperature above 30° in ndh-efflorescing crystals con- 
taining 7 mols. H20. It is soluble in 4-5 parts water, and 
shows a feeble alkaline reaction. The solution absorbs large 
quantities of carbon dioxide, without suffering any alteration. 
When heated the salt loses water, melts about 300° and be- 
Na2C03 + C02 + H20 = 2NaHC03. SODIUM BORATE. 
293 
comes Sodium Pyrohosphate—Na4P207—which crystallizes with 
io molecules of H20, and upon boiling with nitric acid passes 
into primary sodium phosphate. 
The primary ox monosodium phosphate—NaH2P04, crystallizes 
with i molecule of H20, and exhibits an acid reaction. At 
ioo° it loses its water of crystallization, and at 200° becomes 
Na2H2P207, disodium pyrophosphate, which at 240° forms sodium 
metaphosphate—NaP03: «. 
H2Na2P20T = 2NaP03 + H20. 
We get various modifications of the metaphosphate, accord- 
ing to the conditions of fusing and'cooling; they are probably 
polymerides, corresponding to the formulas Na2P206, Na3P309 
etc. Upon heating sodium metaphosphate with metallic ox- 
ides the latter dissolve, and salts of orthophosphoric acid are 
formed, e. g.: 
NaP03 + CaO •= NaCaP04. 
In this manner, characteristic colored glasses (phosphorus 
beads) are obtained with various metals. In blowpipe analysis 
this behavior serves for the detection of the respective metals. 
The salts of arsenic acid are perfectly analogous.to those of phosphoric 
acid. Of the antimoniates may be mentioned the disodium pyroantimoniate, 
Na2H2Sb207 -f- 6H20, which is insoluble in cold water, and is therefore 
precipitated from the soluble sodium salts on the addition of dipotassium 
pyroantimoniate. • 
Sodium Borate. The normal salts of boric acid, B(OH)s, 
and metaboric acid, BO.OH (see p. 239), 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: 
4B(OH)3 - 5h2o = h2b4o7. 
The most important of the salts is borax, which crystallizes 
at ordinary temperatures with io molecules of H20 in large 
monoclinic prisms, Na2B407 + ioH,20. Borax occurs naturally 
in some lakes of Thibet, whence it was formerly imported 
under the name of tinkal. At present, it is prepared artificially 
by boiling or fusing boric acid with sodium carbonate. At or- 
dinary temperatures, the crystals dissolve in 14 parts water; 
at ioo° in one-half part; and the solution has a feeble alkaline 
reaction. When heated to 70° rhombohedra crystallize from 
the solution, and have the composition Na2B407 + SH,0, for- 
merly known as octahedral borax. Both salts puff up when 294 
INORGANIC CHEMISTRY. 
heated, lose water, and yield a white, porous mass (burned 
borax), which fuses at a red heat to a transparent vitreous mass 
(Na2Bt07). In fusion this dissolves many metallic oxides, form- 
ing 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 
blowpipe tests for the detection of certain metals. Upon this 
property of dissolving metallic oxides depends the application 
of borax for the fusion and soldering of metals. 
Sodium Silicate—sodium water glass—is analogous to the 
potassium salt, and is most readily obtained by fusing quartz 
with sodium sulphate and charcoal. 
The sulphur compounds of sodium are also analagous to 
those of potassium. 
Recognition of Sodium Compounds.—Almost all the 
sodium salts are easily soluble in water, sodium pyroantimoniate 
—H2Na2Sb207—excepted ; tbis is precipitated from solutions 
of sodium salts by potassium pyroantimoniate, and can serve 
for the detection of sodium. Sodium compounds, exposed in 
a colorless flame, impart to the latter an intense yellow. The 
spectrum of the sodium flame is characterized by a very bright 
yellow line, which, when more strongly magnified, splits into 
two lines. 
LITHIUM. 
Lithium only occurs in nature in small quantities, but is 
tolerably widely disseminated, and is found in some mineral 
springs and in the ashes of many plants, notably in that of 
tobacco and the beet. As compound silicate, it occurs in 
lepidolite or lithia mica; as phosphate (with iron and man- 
ganese) in triphylite. 
The metal is separated from the chloride by means of the 
electric current, and is silver-white in color, decomposing 
water at ordinary temperatures. Its specific gravity is 0.59. 
It is the lightest of all the metals, and swims upon naphtha. 
It melts at 180°, and burns with an intense white light. 
The lithium salts are very similar to the salts of sodium, 
but closely approach those of magnesium (p. 274). 
Lithium Chloride—LiCl—crystallizes, at ordinary tem- 
peratures, in anhydrous, regular octahedra; below io°, how- 
ever, it has two molecules of H20, and deliquesces in the air. 
Li = 7- AMMONIUM COMPOUNDS. 
295 
Lithium Phosphate—Li3P04 + H20—and Lithium 
Carbonate—Li2C03—are difficultly soluble in water; there- 
fore they are precipitated from solutions of lithium salts by 
sodium phosphate or carbonate. By strong ignition the carbon- 
ate loses carbon dioxide. So far as these two salts are concerned, 
lithium approaches the metals of the calcium group (p. 274). 
Its compounds color the flame a beautiful red; the spectrum 
shows an intense red line. 
AMMONIUM COMPOUNDS. 
Upon page 12 7 we observed that ammonia combines directly 
with the acids to form salt-like compounds, which are analo- 
gous to the metallic salts, especially those of potassium. The 
monovalent group, NH4, playing the role of metal in these deri- 
vatives, is called ammonium, and the derivatives of ammonia, 
ammonium 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 
prepared by letting the electric 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. Then as in the case of the decomposition 
of potassium or sodium chloride, a metal—ammonium—sepa- 
rates on the negative pole, and combines to an amalgam with 
mercury. The amalgam may also be obtained if sodium amalgam 
be covered with a concentrated solution of ammonium chloride: 
(Hg + Na) and NH4C1 yield (Hg + NH4) and NaCl. 
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 solu- 
tion ; -no proofs, however, exist which would lead us to accept 
the existence of ammonium hydroxide (NH4OH) in the solu- 
tion. On the other hand, there are organic derivatives of 
ammonium hydrate, in which the hydrogen of the ammonium 
is replaced by hydrocarbon residues; e.g., tetramethyl ammo- 
nium hydroxide—N(CH3)4OH. These are thick liquids, of 
strong basic reaction and, in all respects, are very similar, to 
potassium and sodium hydroxides. 
Sodium amalgam. 
Ammonium amalgam. 296 
INORGANIC CHEMISTRY. 
Ammonium Chloride—NH4C1—is sometimes found in 
volcanic districts, and was formerly obtained by the dry dis- 
tillation of camel’s dung (Sal ammonia cum). At present it is 
prepared almost exclusively by saturating the ammonia water 
from gas works with hydrochloric acid. The solution is evap- 
orated to dryness, and the residue heated in iron vessels, 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, feather-like, 
grouped octahedra or cubes, of sharp, salty taste. When 
heated, ammonium chloride sublimes without melting; at the 
same time a dissociation into NH3 and HCl’is sustained, but 
these products recombine again to ammonium chloride, on 
cooling. The dissociation is complete at 350°, and the vapor 
density then equals 13.3 (H = 1) corresponding to that of a 
mixture of similar molecules, of NH3 (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 sat- 
urating the ammonia water from gas works with sulphuric acid. 
It crystallizes without water in rhombic prisms, and is 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 
potassium nitrate and deliquesces in the air. When heated it 
melts, and then decomposes into hyponitrous oxide and water 
(p. 209).. 
. Ammonium Nitrite—NH4N02—is present in minute 
quantities 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 
the decomposition of silver or lead nitrite by ammonium chlo- 
ride. Heat decomposes it into nitrogen and water (p. 115). 
Ammonium Carbonate.—The ?ieutral or secondary salt 
(NH4)2C03, separates as a crystalline powder, when ammo- 
nia gas is conducted through a concentrated solution of the so- 
called sesquicarbonate. It parts with ammonia in the air and 
becomes the primary or acid salt, NH4HCOs, which when 
■heated to 58°, decomposes into carbon dioxide, ammonia, and 
water. 
The common, commercial, so-called sesquicarbonate of am- AMMONIUM HYDROSULPHIDE. 
297 
tnonium, (C02)2(NH3)3H20, which can be regarded as a com- 
pound of primary ammonium carbonate with ammonium carba- 
mate, C03(NH4) H + NH2C02 NH4 (see Organic Chemistry), 
arises in the decay of many nitrogenous carbon compounds, e.g., 
the urine, and was formerly prepared by the dry distillation of 
bones, horn, and other animal substances. At present it is ob- 
tained by heating a mixture of ammonium chloride, or sulphate, 
with calcium carbonate. It then sublimes as a white, transpar- 
ent hard mass, which gives off ammonia and carbon dioxide in 
the air, falling into a white powder of primary ammonium car- 
bonate. The latter, obtained by the efflorescence of the first 
two salts, or by saturating ammonium hydrate with carbon di- 
oxide, is a white, odorless powder, more insoluble in water. 
In aqueous solution it gradually loses carbon dioxide and is 
changed to the secondary carbonate. 
Ammonium Phosphates.—The most important of 
these is the secondary ammonium sodium phosphate, P04( NH4) 
NaH + 4H20, ordinarily termed salt of phosphorus. It is 
found in guano and decaying urine. It can be obtained by 
the crystallization of a mixture of di-sodium' phosphate and 
ammonium chloride: 
Na2HP04 + NH4C1 = NH4NaHP04 + NaCl. 
It consists of large, transparent, monoclinic crystals. When 
heated it fuses, giving up water and ammonia and forms a 
transparent glass of sodium metaphosphate, NaPOs (p. 293). 
It serves in blow-pipe 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 the 
secondary salt, (NH4)2HP04, which changes to the primary salt, 
PG4(NH4)H2, when its solution is boiled. 
Ammonium Sulphide—(NH4)2S—results upon mixing 1 
volume H2S with 2 volumes NH3at —180. It is a white crys- 
talline mass, decomposing, at ordinary temperatures, into 
NH4HS and NH3. It is obtained in aqueous solution by the 
saturation of an ammonium hydrosulphide solution with ammo- 
nia ; and also seems to dissociate into its constituents. 
Ammonium Hydrosulphide — NH4SH—is produced 
upon conducting hydrogen sulphide into an alcoholic ammonia 
solution. It is obtained in aqueous solution by saturating 
aqueous ammonia with hydrogen sulphide. At first the solu- 
tion is colorless, but becomes yellow on standing in contact 298 
INORGANIC CHEMISTRY. 
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 color- 
less hydrosulphide. Both solutions are often employed in 
laboratories for analytical purposes. 
Recognition of Ammonium Compounds.—All am- 
monium salts are volatile and decompose upon heating. The 
alkalies and other bases liberate ammonia from them, which is 
recognized by its odor and the blue color it imparts to red 
litmus paper. Platinum chloride produces a yellow crystalline 
precipitate of ammonio-platinum chloride, PtCl4,2NH4C1, in 
solutions of ammonium chloride. Tartaric acid precipitates 
primary ammonium tartrate. 
The heat of formation of the most important ammonium compounds 
corresponds to the following symbols : 
(N,H3) = 11.9 
(N,H4,C1) = 75-8 
(NH4,€l,Aq.)= — 3.8 
(N2,H8)S,04)= 282.1 
(NHjAq.) = 8.4 
N,H4,Br) = '65.3 
(NH4,Br,Aq.)=— 4.3 
(N,H4,N,Os) = 88.1 
(N,H8,S,H) s= 39.0 
(N,H4,I) = 49.3 
(NH4,l,Aq.) =— 3.5 
(N,H4,N,02) = 64.9 
METALS OF THE SECOND GROUP. 
Ca 39.9. 
Sr 87.3. 
Ba 136.8. 
Be 9.0. 
Mg 23.9. 
Zn 64.9. 
Cd 111.9. 
Hg 199.8. 
The second group of the periodic system (see table, p. 244) 
comprises chiefly the divalent metals, which form compounds 
of the type MeX2, and in their entire deportment exhibit 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 ap- 
proaches aluminium ; just as lithium attaches itself to magne- 
sium (p. 274). The metals, calcium, strontium, and barium, 
constitute the second members of the three great periods, are 
among themselves perfectly homologous (p. 241), and in accord 
with their strong basic character, -attach themselves to the 
alkali metals K, Rb, and Cs. Zinc, cadmium, and mercury, 
which correspond to them and constitute the second sub- 
group, really belong 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 GROUP OF THE ALKALINE EARTHS. 
299 
combination types. In consequence of the double periodicity 
of the three great periods both sub-groups (Ca, Sr, Ba, and 
Zn, Cd, Hg) exhibit many analogies to magnesium and 
beryllium. 
GROUP OF THE ALKALINE EARTHS. 
Calcium. 
Strontium. 
Barium. 
The metals of this group are termed alkaline earth metals, 
because their oxides attach themselves in their properties, on 
the one side to the oxides of the alkalies, upon the other to 
the real earths (alumina, etc.) They show the same gradation 
in properties 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 en- 
ergetically, and oxidizes more readily than strontium and cal- 
cium. In accord with this, we find barium hydroxide a stronger 
base; it dissolves rather readily in water,-does not decompose 
upon ignition, and absorbs carbon dioxide rapidly from the 
air. Barium carbonate is also very stable, fuses at a white 
heat, and only disengages a little carbon dioxide. Calcium 
hydroxide, on the other hand, is much more difficultly soluble, 
in water, and when ignited breaks up into water and calcium 
oxide ; the carbonate also yields carbon dioxide when simi- 
larly treated. In its entire character, strontium stands between 
barium and calcium. All these affinity relations find full ex- 
pression in the heat of formation of the corresponding com- 
pounds. 
While the alkaline earth metals are similar to the alkalies 
in their free condition and in their hydrates, they essentially 
distinguish themselves from them by the insolubility of their 
carbonates and phosphates, and still more by their sulphates. 
Barium sulphate is almost insoluble* in water and acids, while 
that of calcium dissolves in 400 parts water ; strontium sul- 
phate occupies an intermediate position. 
Ca = 39.9. 
Sr = 87.3. 
Ba = 136.8. 
The metals of this group do not form volatile compounds and their spe- 
cific heats have not yet been determined. As the determination of the va- 
por densities of the elements or their volatile compounds, and the ascer- 
tainment 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 make the 
atomic weights of the calcium group equal to their equivalent weights 
I 
(Ca = 19-9, CaCl). But the great analogy of their compounds to those 300 
INORGANIC CHEMISTRY. 
of the metals of the magnesium group, their isomorphism for instance, 
argues with great probability for the divalence of the metals of the group, 
and that the present accepted double atomic weights are correct (compare 
p. 258). For calcium, this conclusion has already been confirmed by the 
experimental determination of its thermal capacity. 
CALCIUM. 
Calcium belongs to the class of elements most widely dis- 
tributed upon the earth’s surface. As calcium carbonate (lime- 
stone, marble, chalk) and the sulphate (gypsum, alabaster), it 
represents immense deposits in all stratified formations. As 
phosphate, it constitutes phosphorite, as fluoride, jluorite, 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 chlo- 
ride; further, by heating calcium iodide with sodium, or cal- 
cium chloride with sodium and zinc. Although the affinity 
of calcium for oxyge’n is less than that of the alkalies, yet the 
oxide (also BaO and SrO) cannot be reduced to metal by ig- 
nition 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. It is prepared on a large 
scale by burning the ordinary limestone or marble (CaC03) in 
lime-kilns. It is a grayish-white mass, which does not fuse 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 the air lime at- 
tracts m'oisture and C02, becoming calcium carbonate; burned 
lime unites with water with evolution of much heat, breaking 
up into a white voluminous powder of calcium hydroxide 
Ca(OH)2—slaked lime. 
Ca = 39.9. 
When limestone contains large quantities of alumina, magnesium car- 
bonate, 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 is 
readily converted into a powder with water. 
Calcium Hydroxide—Ca(OH)2—slaked lime—is a white, CALCIUM PEROXIDE—CALCIUM CHLORIDE. 
301 
porous powder, forming a thick paste, milk of lime, with water. 
It dissolves with difficulty in cold water (1 part in 760 parts), 
but with still more difficulty in warm water; the solution 
saturated in the cold (lime water) becomes cloudy upon warm- 
ing. It has a strong alkaline reaction. 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 mor- 
tar, a mixture of calcium hydroxide, water, and sand. The 
hardening Qf the mortar in the air depends mainly upon the 
fact that the calcium hydroxide combines with the C02 of the 
air to form the carbonate, and at the same time acts upon the 
silicic acid of the sand forming a calcium silicate, which, in 
time, imparts durability to the mortar. 
Hydraulic mortar, or cement, is produced by gently igniting 
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 upwards of 20 per 
cent, clay, yield hydraulic cements, without any admixtures 
after burning. Their composition varies, and also the pro- 
cess of their hardening; it however depends principally upon 
the formation of calcium and aluminium silicates. 
Calcium Peroxide—Ca02—is precipitated as a hydrate 
in crystalline leaflets, if lime water be added to a solution of 
barium peroxide in dilute hydrochloric acid; it is very unstable. 
It contains 8H20, which it gradually loses in dry air. 
Th.e halogen derivatives of calcium, like those of other 
metals, are prepared by the solution of the oxide or carbonate 
in the haloid acids. They are formed by the direct union of 
calcium with the halogens; calcium burns in the vapors of 
chlorine, bromine, and iodine. Technically, calcium chloride 
is often obtained as a by-product, eg., in the preparation of 
ammonia. 
Calcium Chloride—CaCl2—crystallizes from aqueous so- 
lution with 6 molecules of H20, 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 it is only after it has been exposed above 200° 
that it becomes anhydrous; then it is a white, porous mass. 
The dry salt fuses at 7190, and 'solidifies to a crystalline mass, 
which attracts water energetically, and may be employed in 302 
INORGANIC CHEMISTRY. 
the drying of gases and liquids. The dry calcium chloride 
also absorbs ammonia, forming the compound CaCl2.8NH3. 
The crystallized hydrous salt dissolves in water with reduction 
of temperature; by mixing with snow or ice the tempera- 
ture is lowered to —48°. Upon fusing the dry chloride in 
air it will partially decompose into the oxide and hydrogen 
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, or even in compact masses. It is 
often discolored by impurities. It is found, in sparing quanti- 
ties, in the ashes of plants, bones, and the enamel of the teeth. 
A soluble fluoride added to the solution of calcium chloride 
throws down insoluble calcium fluoride as a white voluminous 
precipitate. 
The fluoride is perfectly insoluble in water, and is only de- 
composed by strong'acids. It fuses easily at a red heat, serv- 
ing, therefore, as a flux in the smelting of ores. When heated 
it phosphoresces. 
Calcium Hypochlorite—Ca(C10)2—is not known in a 
pure condi tion. The so-called bleaching lime or chloride of lime, 
obtained by conducting chlorine, at ordinary temperatures, 
over slaked lime, contains calcium hypochlorite as active prin- 
ciple. 
According to analogy to the action of chlorine upon potas- 
sium, or sodium hydrate, the reaction in the case of calcium 
hydrate may be expressed by the following equation : 
This would incline us to regard chloride of lime as a mixture 
of calcium hypochlorite and water. In accordance with the 
equation, the completely chlorinated chloride of lime must 
contain 48.9 per cent, chlorine, which is never the case, 
because a portion of the calcium hydroxide invariably remains 
unaltered. Calcium chloride does not exist free in bleaching 
lime, because it is not withdrawn from the latter by alcohol, and 
nearly all the chlorine of the bleaching lime can be expelled 
by carbon dioxide. It is therefore probable that the compound, 
/Cl ... 
Ca (Odling and Lunge), is present in bleaching lime: 
\OCl 
2Ca(OH)2 + 2C12 = Ca(OCl)2 + CaCl2 + 2H20. 
Ca(OCl)2 + CaCl2 = 2^a\^oCl’ CALCIUM SULPHATE. 
303 
Chloride of lime is a white, porous powder with an odor re- 
sembling that of chlorine. The aqueous solution has a strong 
alkaline reaction, and bleaches. It decomposes in the air as 
the carbon dioxide of the latter liberates hypochlorous acid. 
Even in closed vessels it gradually 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 ot lime \ the quantity liberated is just twice that which 
the hypochlorite eventually found in it contains: 
Ca(C10)2 + 4HC1 = CaCl2 + 2H20 +' 2C12. 
When sulphuric acid acts, the calcium chloride present par- 
ticipates in the reaction : 
Ca(C10)2 -f- CaCl2 -(- 2H2S04 = 2CaS04 -f- 2C12 -f- 2H20 
The application of chloride of lime for the production of 
chlorine in chlorine bleaching and disinfection is based on this 
deportment. 
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 per cent. 
Calcium chlorate and chloride are produced when the aqueous solution 
of chloride of lime is boiled : 
On this is based the application of chloride of lime for the production of 
potassium chlorate (KC103) by a transposition of calcium chlorate with 
potassium chloride. 
When a small quantity of cobaltic oxide is added to the solution of 
bleaching lime, and heat applied, a regular stream of oxygen is disengaged ; 
this is an advantageous method of preparing oxygen. Other oxides, like those 
of manganese, copper, and iron, behave similarly. In this reaction there 
occurs, apparently, a contact action of the oxides. The reaction is explained, 
doubtless, in the same way as the action of hydrogen peroxide upon cer- 
tain oxides (see p. ioo). The feebly combined oxygen atom in cobaltic 
oxide unites with the oxygen of the calcium hypochlorite to form free 
oxygen: 
3Ca(C10)2 = (C103)2Ca -f- 2CaCl2. 
Ca(C10)2 + 2Co2Os = CaCl2 + 202 + 4C0O. 
Cobaltic 
oxide. 
Cobaltous 
oxide. 
The resulting cobaltous oxide is again converted by the chloride of lime 
into cobaltic oxide, which acts upon a fresh quantity of bleaching lime. 
Calcium Sulphate—CaS04—is very abundant in nature. 304 
INORGANIC CHEMISTRY. 
In an anhydrous condition it forms the mineral anhydrite, 
crystallizing in forms of the rhombic system. With two mole- 
cules of water it occurs as gypsum, in large monoclinic crystals 
or in granular, crystalline masses (Alabaster, etc.). It also 
separates as a fine crystalline powder, CaS04 -(-2H20, when 
soluble calcium salts are precipitated with sulphuric acid. 
Calcium sulphate is difficultly soluble in water ; i part at 
average temperatures dissolves in 400 parts H20. When heated 
to no0 gypsum loses all its water, and becomes burnt gypsum, 
which, pulverized and mixed with water, forms a paste which 
hardens to a solid mass in a short time. The hardening is de- 
pendent upon the reunion of anhydrous calcium sulphate with 
2 molecules of H20. On this depends the use of burned gyp- 
sum for the production of moulds, figures, etc. In case gypsum 
has been too intensely heated (dead-burnt gypsum) it wilt no 
longer harden with water; the naturally occurring anhydrite 
behaves in the same manner. 
Calcium Nitrate—Ca (N03)2—is produced by the decay 
of nitrogenous organic substances in the presence of lime, 
therefore, it is frequently found as an efflorescence upon walls 
(in cattle stables). It crystallizes from water 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 
nitrate (p. 280). 
Calcium Phosphates. The tertiary phosphate—Ca3 
(P04)2—is found in slight quantities in most of the mountain 
rocks. * In combination with calcium fluoride, it crystallizes 
as apatite. As phosphorite, it forms compact masses, more or 
less intimately mixed with other constituents, and occurs in 
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 
accumulates chiefly in the seeds and grains. In the animal 
kingdom, it is principally found in the bones, the ashes of 
which contain upwards of 85 per »cent. calcium phosphate. 
Tertiary calcium phosphate is perfectly insoluble in water. If 
disodium phosphate be added to the aqueous solution of a cal- 
cium salt, and then ammonium hydrate, it will separate as a 
gelatinous precipitate, which, after drying, becomes a white 
amorphous powder. It is very readily soluble in acids, even 
acetic. • 
The secondary calcium phosphate—P04CaH -(- 2H20—is 
sometimes present in guano, in the form of small, shining CALCIUM CARBONATE. 
305 
prisms, and separates as an amorphous precipitate, if disodium 
phosphate be added to a solution of calcium chloride mixed 
with some acetic acid. When ignited, it passes into calcium 
pyrophosphate, P207Ca2. 
The priynary phosphate—Ca(H2P04)2—is produced by the 
action of sulphuric or hydrochloric acid upon the first two 
phosphates. It is readily soluble in water, and deliquesces in 
the air. Heated to 200°, it decomposes into pyrophosphate, 
metaphosphcrric acid and water : 
2Ca(H2P04)2 = Ca2P207 + 2HPO3 + 3H20. 
When intensely ignited, pure calcium metaphosphate re- 
mains (p. 131). 
Calcium phosphate is present in all plants. Its presence in the soil is 
therefore, an indispensable condition for its fertility. When there is a 
scarcity of phosphoric acid it must be added. To this end, bone meal and 
pulverized phosphorite were formerly employed. Since, however, the 
phosphoric acid is contained in these substances as tri-calcium phosphate, 
which is not easily absorbed by the plants, the primary phosphate is exten- 
sively employed at present 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 resulting mass. 
Calcium Carbonate—CaC03—is very widely distributed 
in nature. It crystallizes in two crystallographic systems, 
hence is dimorphous. In rhombic crystals, with the specific 
gravity 3.0, it forms aragonite. In hexagonal rhombohedra, 
with specific gravity 2.7, it occurs as calcite. Iceland spar, 
employed for optical purposes, is perfectly pure, transparent 
calcite. The common calcite, which constitutes immense 
mountain chains, is an amorphous or indistinct crystalline 
stratum, and is usually mixed with other constituents, as clay. 
When the limestone is granular and crystalline, it is termed 
marble. Dolpmite also constitutes large layers, and is a com- 
pound of calcium and magnesium carbonates, with generally an 
excess of the former. Chalk is very pure amorphous calcium 
carbonate, consisting of the shells of microscopic sea animals. 
Calcium carbonate is, further, a regular constituent of all plants 
and animals; the shells of eggs, of mussels, even corals and 
pearls, consist chiefly of it. 
A soluble carbonate, added to the aqueous solution of a cal- 
cium salt, precipitates calcium carbonate as a white, amorphous 
powder, which soon becomes crystalline. In the cold, it as- 
sumes the form of calcite ; upon boiling the liquid, it gener- 
ally changes into aragonite crystals. 306 
INORGANIC CHEMISTRY. 
The carbonate is almost insoluble in pure water ; but dis- 
solves somewhat in water containing carbon dioxide, as it 
very probably is changed to the primary carbonate—Ca(HC03)2. 
For this reason, we find calcium carbonate dissolved in all 
natural waters. When the solution stands exposed, or if it be 
heated, carbon dioxide escapes, and the secondary carbonate 
again separates out. The formation of lime scales, thermal 
tufts, stalactites, boiler scales and similar deposits are due to 
this. Calcium carbonate, like all carbonates, is decomposed 
by acids, with evolution of carbon dioxide. At a red heat, it 
decomposes into CaO and C02. 
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. 
Glass.—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 difficulty 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 sodium a mixture of sul- 
phates with charcoal can be employed; the carbon reduces the sul- 
phates to sulphides, which form silicates when fused with silicon di- 
oxide. 
The following are the varieties of glass : 
Soda glass—a mixture of sodium and calcium silicates—fuses readily, 
and is employed for window-panes and ordinary glass vessels. Potash 
or Bohemian Glass, also called Crown Glass,, consists of calcium and po- 
tassium silicates, is not very fusible, is harder, and withstands the action 
of water and acids better than soda glass; it is, therefore, employed in the 
manufacture of chemical glassware. 
Glass Crystal or Flint Glass is composed of potassium and lead silicates. 
It is not as hard, fuses with tolerable readiness, refracts light strongly, 
and when polished, acquires a clear lustre. On this account it is employed 
for optical purposes (for lenses, prisms) and is used in ornamental glass- 
ware. Strass,—a lead glass containing boron trioxide, is used to imitate 
precious stones. The opaque varieties of enamel consist of lead glass and 
in the fused glass are insoluble admixtures, as tin dioxide and calcium 
phosphate. 
Ordinary window glass is obtained by the fusion of rather impure mate- 
rials; in consequence of the presence of ferrous oxide it is ordinarily col- 
ored green. To remove this coloration, manganese peroxide, MnCb 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. These colors, violet and green, almost neutralize each other 
as complementaries. The colored glasses contain silicates of colored me- 
tallic oxides; chromic and cupric oxides color green; cobaltic, blue ; 
cuprous oxide, a ruby red, etc. STRONTIUM. 
307 
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. When it is dissolved in water we get 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 is obtained, 
which consists of calcium polysulphides and sulphate. Milk 
of lime boiled with sulphur yields a deep, yellow solution of cal- 
cium polysulphides. When the solutions of the latter are 
acted upon by acids, finely divided sulphur—milk of sulphur 
—is precipitated and H2S set free. If the reverse occur, viz., 
the addition of a solution of polysulphides to an excess of 
dilute acids, hydrogen persulphide will separate (p. 108). 
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. 
The metal is obtained by the electrolysis of fused strontium 
chloride. It has a brass-yellow color, and a specific gravity, 
2.5. It oxidizes in the air and burns with a bright light when 
heated. It decomposes water at the ordinary temperatures. 
Of the compounds of strontium we may mention the fol- 
lowing : 
Strontium Oxide—SrO—is most readily obtained by 
igniting the nitrate. It unites with water, with strong evolu- 
tion of heat, forming Strontium Hydrate—Sr(OH)2— 
which is more readily soluble in water than is calcium hydrate. 
It crystallizes from aqueous solution 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—SrS04—is much more difficultly 
soluble in water than calcium sulphate, but is not as insoluble 
as barium sulphate. 
Strontium Nitrate—Sr(N03)2—is obtained by dissolving 
the carbonate in nitric acid, and is readily soluble in water. 
It crystallizes from warm solutions in anhydrous octahedra, 
Sr =87.3. 308 
INORGANIC CHEMISTRY. 
but from cold, with 4 molecules H20, in monoclinic prisms. 
Mixed with combustible substances it colors the flame a beau- 
tiful carmine red, and for this reason is employed in pyrotechny. 
Strontium Carbonate—SrC03—is precipitated from 
aqueous solutions of strontium salts by soluble carbonates,-as an 
amorphous, insoluble power. When ignited it breaks up into 
SrO and C02. This decomposition does not, however, occur 
as easily as with calcium carbonate. 
Ba= 136.8. 
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 fiapbs, heavy. In accordance 
with its general character barium is a stronger basic metal than 
either strontium or calcium (p. 299). 
The barium salts are either prepared from the natural with- 
erite, 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 and 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 amalgam 
is kneaded with water, to remove all the sodium, and then 
heated in a stream of hydrogen, to volatilize the mercury. 
Barium is a bright yellow metal, of specific gravity 3.6.' It 
fuses at a red heat, but does not vaporize. It is rapidly oxi- 
dized in the air; like sodiun\ it decomposes water very ener- 
getically, even at ordinary temperatures. 
Barium Oxide—BaO—:is obtained by the ignition of 
barium nitrate. It is a gray, amorphous mass, of specific 
gravity 5.5, and fusible in the oxy-hydrogen flame. With 
water it yields the hydroxide, with evolution of much heat. 
Barium Hydroxide—Ba(OH)2—is precipitated from con- 
centrated solutions of barium salts by potassium or sodium 
hydrate, not, however, by ammonium hydrate. At ordinary BARIUM. 
309 
temperatures it dissolves in 20 parts, upon boiling, in 3 parts 
water. From aqueous solution it crystallizes with 8 molecules 
of H20 in four-sided prisms or leaflets. The solution—called 
Baryta water—is strongly alkalitte and is very similar to the 
alkalies. When exposed to the air it absorbs carbon dioxide 
and becomes turbid, with separation of barium carbonate. At 
a red heat it fuses without decomposition like the caustic 
alkalies, 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 con- 
tains 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. An excess of baryta 
water is added to the solution, containing barium chloride and 
hydrogen peroxide, when hxdrated bariutn peroxide—Ba02 
8H20—separates in shining scales, which, upon warming, 
readily lose water and break up into a white powder consist- 
ing of barium peroxide. The latter is a compact gray mass 
when obtained directly from the oxide. 
The peroxide dissolves in dilute acids, with production of 
hydrogen peroxide. Concentrated sulphuric acid sets free 
ozonized oxygen from it. When strongly ignited (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 wrater. It is employed as a green 
fire in pyroteehny. 
Barium Sulphate—BaS04—is found in nature as heavy 
spar, in rhombic prisms, with a specific gravity of 4.6. It is 
obtained artificially by the precipitation of barium salts with sul- 
phuric acid in the form of a white, amorphous powder, almost 
insoluble in water and acids. Under the name of permanent 
white, it is used as a paint, as a substitute for poisonous white 
lead, from which it is also distinguished by its unalterability. 
Barium Carbonate—BaCOs—as witherite, occurs in 
shining, rhombic crystals, and is precipitated from barium 
solutions by soluble carbonates, as a white, amorphous powder. 
It fuses at a white heat, and loses some carbon dioxide. 
Barium Sulphide—BaS—is obtained by igniting the sul- 310 
INORGANIC CHEMISTRY. 
phate with carbon. It dissolves in water, with decomposition 
into hydroxide and hydrosulphide. 
RECOGNITION OF THE COMPOUNDS OF THE ALKALINE EARTHS. 
The carbonates and phosphates of this group are insoluble 
in water; hence are precipitated from the aqueous solutions of 
their salts upon the addition of soluble carbonates and phos- 
phates (of the alkalies). The sulphates are also insoluble in 
acids (only calcium 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 calcium, 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 concentrated solutions. In solutions of barium salts hydro- 
fluosilicic acid produces a crystalline precipitate of barium silico- 
fluoride, BaSiFl6. 
The flame colorations produced by the volatile compounds 
are very characteristic; calcium salts impart a reddish-yellow 
color; strontium, an intense crimson; barium, a yellowish- 
green. The 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, 
an orange and a blue, which are less distinct but very charac- 
teristic. Finally, the barium spectrum consists of several orange, 
yellow, and green lines, of which a bright green is particularly 
prominent. 
METALS OF THE MAGNESIUM GROUP. 
Jn this group are usually included beryllium, magnesium, 
zinc, and cadmium. However, these metals do not exhibit 
complete analogy, as is clearly seen in the periodic system (p. 
298). Beryllium shows the greatest variations. It approaches 
aluminium, while magnesium resembles not only zinc and cad- 
mium, but also the alkaline earth metals, calcium, strontium, 
and barium. Its similarity to the latter is expressed by the 
basic nature of its oxide, whereas it resembles zinc and cadmium 
mainly in isomorphism of compounds. 
Beryllium and magnesium bear the same relation to Ca, Sr, 
and Ba, as lithium and sodium bear to the metals of the potas- 
sium group. METALS OF THE MAGNESIUM GROUP. 
311 
The alkaline character of the alkaline earth-metals gradually 
diminishes from barium to calcium, and becomes almost nothing 
in magnesium and beryllium, which possess the lowest atomic 
weights (see p. 298). Magnesium and beryllium are scarcely 
capable of decomposing water, even at boiling temperatures. 
Their oxides and hydroxides are almost insoluble in it; the 
hydroxides decompose, on gentle ignition, into oxides and 
water. Their carbonates are very unstable ; their chlorides, 
too, suff a partial decomposition into oxide and hydrogen 
chloride, even on drying. The solubility of the sulphates of 
magnesium and beryllium further distinguishes them from the 
metals of the alkaline earth group. 
The specific properties of beryllium and magnesium are 
maintained in zinc and cadmium, which constitute a natural 
group with the former. Zinc and cadmium do not decompose 
water at a boiling heat; their hydroxides are insoluble in it, and 
are not very stable; their carbonates and chlorides easily un- 
dergo decomposition ; their sulphates are readily soluble in 
water. The similarity is further expressed by the isomorphism 
of most of their compounds. Thus, magnesium and «inc sul- 
phates crystallize with 7 molecules of H20, in perfectly simi- 
lar forms. If the solution of a mixture of both salts be allowed 
to crystallize, we get crystals with variable quantities of zinc 
and magnesium : the formation of such isomorphous mixtures 
in ad libitum proportions, is a characteristic indication of the 
isomorphism of compounds chemically similar. 
The difference between beryllium and magnesium upon the 
one side, and zinc and cadmium on the other, is shown dis- 
tinctly in their specific gravities. While the first two elements 
possess a low specific gravity (Be—2.1, Mg—1.75), zinc and 
cadmium (with specific gravities 7.2 and 8.6) belong to the 
so-called heavy metals (see p. 253). 
The difference in specific gravity determines, also, many 
differences in chemical character. The light metals (especially 
the alkalies and alkaline earths) form rather unstable sulphides, 
readily soluble in water, while the sulphides of zinc and cad- 
mium, like those of all heavy metals, are insoluble in water, 
and, usually, in acids; in these respects, magnesium and be- 
ryllium behave like the alkalies, while zinc and cadmium are 
precipitated by hydrogen sulphide or alkaline sulphides as 
sulphides from solutions of their salts. Further, the oxides of 
the light metals are very stable, and are only reduced by car- 
bon if they are readily fusible (like potassium and sodium 
oxides); the heavy metals, on the other hand, are easily sepa- 312 
INORGANIC CHEMISTRY. 
rated from their oxides by carbon. Zinc and cadmium oxides 
are reduced by carbon, while those of magnesium and beryl- 
lium are not altered. All these affinity relations are more 
clearly expressed and explained in their thermo-chemical re- 
lations (p. 320.). 
MAGNESIUM. 
■ Magnesium is abundant in nature, and almost always accom- 
panies 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. Magnesium is also present in most of 
the natural silicates ; its soluble salts are contained in almost 
all natural waters. 
Metallic magnesium may be obtained by the electrolysis of 
the chloride, or by heating the same with sodium. It is pre- 
pared, on a large scale, by heating the double chloride of 
magnesium and sodium with metallic sodium : 
Mg = 23.9. 
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. It is tenacious and ductile, and when 
heated may be converted into wire, and rolled out into thin 
ribbons. It fuses at a dull, and distils at a bright red heat. At 
ordinary temperatures, it scarcely oxidizes in the air; it burns 
when heated, with an extremely intense white light, owing to 
the glowing non-volatile magnesium oxide. Magnesium light 
is rich in chemically active rays, and, for this reason, it is em- 
ployed for photographing in dark chambers. Its alloy with 
zinc is generally employed as a substitute for pure magnesium, 
as it burns with an almost equally bright light. Boiling water 
is very slowly decomposed by magnesium. It dissolves easily 
in dilute acids, forming salts; the alkalies do not attack it. 
Magnesium Oxide—MgO—or magnesia, formed by the 
combustion of magnesium, is ordinarily obtained by the ig- 
nition of the hydrate or the carbonate (magnesia usta). It is 
a white, very voluminous, amorphous powder, which finds ap- 
plication in medicine. The feebly ignited magnesia combines 
with water, with slight generation of heat, to produce mag- 
nesium hydrate. 
MgCl2.NaCl -f 2Na = 3NaCl -f Mg. MAGNESIUM HYDROXIDE—MAGNESIUM CHLORIDE. 
313 
Magnesium Hydroxide—Mg(HO)2—is precipitated from 
solutions of magnesium salts by potassium or sodium hydrate 
as a gelatinous mass. Dried at ioo° it is a white amorphous 
powder. It is almost insoluble in water and alkalies; moist 
litmus paper is, however, colored blue. Ammonium salts dis- 
solve it quite easily, forming soluble double salts. Magnesium 
hydroxide attracts carbon dioxide from the air and forms mag- 
nesium carbonate. It yields the oxide and water when gently 
ignited. 
Magnesium Chloride—MgCl2—is present in traces in 
many mineral springs. It may be obtained by the solution of 
the carbonate or oxide in hydrochloric acid ; in large quanti- 
ties it is obtained as a by-product in the technical production 
of potassium chloride. When its solution is evaporated the 
salt crystallizes out with six molecules of H20 in deliquescent 
crystals, isomorphous with calcium chloride. When these are 
heated they give up water, and there occurs at the same time 
a partial decomposition of the chloride into oxide and hydro- 
gen chloride: 
MgCl2 + H20 = MgO -f 2HCI. 
As magnesium chloride is produced in large quantities in various tech- 
nical processes, repeated efforts have been made to utilize the above 
reaction for the preparation of hydrochloric acid, by conducting steam over 
heated magnesium chloride. However, the thermal relations of the re- 
action indicate that this could only be accomplished with difficulty. From 
the chemical affinities coming into play, the reaction pursues an opposite 
course, as the magnesium oxide is readily decomposed by hydrochloric 
acid into magnesium chloride and water: 
MgO + 2HC1 = MgCl2 + H20. 
This is because the heat of formation of MgCh (T51.0C.) and steam (58.0) 
is greater than that of MgO (145.8) and 2HCI 144.0). The decomposi- 
tion is even easier in the presence of water, as is evident from the thermal 
numbers. The reverse reaction is, therefore, endothermic, requires the 
addition of chemical energy in the form of heat, and like all similar reac- 
tions is incomplete. 
To get anhydrous magnesium chloride ammonium chloride 
is added to the solution of the former. The double salt, 
MgCl2,NH4Cl -f 6H20 is formed. When this is heated it first 
loses water, and at 460° throws off ammonium chloride, leav- 
ing anhydrous magnesium chloride. This is a leafy, crystal- 
line mass, which fuses at 708°, and distils undecomposed at a 
red heat; it is very deliquescent in the air. 
Double salts, similar to the above, are also formed from po- 
tassium and calcium chloride. The potassium double salt 
—MgCl2,KCl -f 6H20—occurs in considerable deposits as 
carnallite at Stassfurt. INORGANIC CHEMISTRY. 
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 ordinary 
temperatures it crystallizes with 7 molecules H,20—MgS04-f 
7H20—in four-sided rhombic prisms, readily soluble in water 
(at o° in 2 parts water). It has a bitter, salt-like taste, and 
serves as an aperient. It crystallizes with 6 molecules of H20 
from solutions heated to 70° ; at o°, however, it has 12 mole- 
cules. When heated to 150° these hydrates lose all their water 
of crystallization, excepting one molecule, which escapes above 
2000. One molecule of water, in magnesium sulphate is, 
therefore, more closely combined than the rest. Many other 
salts containing water deport themselves similarly. The more 
intimately combined water is termed Water of Constitution. 
Magnesium sulphate forms double salts with potassium and 
ammonium sulphates, which crystallize with 6 molecules of H20 
in monoclinic prisms, e.g. : 
The sulphates of zinc and several other metals, e.g., iron, cobalt, and 
nickel, in their divalent forms, are very similar to magnesium sulphate. Their 
sulphates crystallize with 7 molecules of H20, are isomorphous, and con- 
tain 1 molecule of intimately combined water. They form double salts 
with potassium and ammonium sulphates; these crystallize with 6H20, 
and are isomorphous; eg. : 
MgS04,K2S04 + 6H20. 
ZnS04 + 7H20 ZnS04.K2S04 -f 6H20. 
FeSOt + 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 divalent metal unites two 
molecules of sulphuric acid : 
/K 
so4 
/Mg + 6H20. 
S04 
\k 
Magnesium Phosphates. — The tertiary phosphate 
(P04)2Mgs, accompanies the tertiary calcium phosphates in 
small quantities in bones and in plant ashes. The secondary 
phosphate, MgHP04 +7H.20, is precipitated from the soluble 
magnesium salts, by disodium phosphate (Na2HP04) as a salt 
difficultly soluble in water. If ammonium salts be present, the 
precipitated double salt will be magnesium ammonium phos- 
phate, MgNH4P04 4- 6H20, insoluble in water. The latter is RECOGNITION OF MAGNESIUM COMPOUNDS. 
315 
found in guano, forms in the decay of urine, and is sometimes 
the ’cause of the formation of calculi. The primary salt, 
H4Mg(P04)2, has not been obtained. 
The magnesium salts of arsenic acid, H3As04, are very sim- 
ilar to those of phosphoric acid. Magnesium-ammonium 
.arseniate (MgNH4As04 -f 6H20) is likewise almost insoluble 
in water. 
Magnesium Carbonate, MgCOa, occurs in nature as 
magnesium spar, crystallized in rhombohedra (isomorphous 
with calcite), in compact masses as magnesite. Combined 
with calcium carbonate, it forms dolomite, to which, when 
pure, is ascribed the formula, CaC03,MgC03; however, it 
usually contains an 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 pre- 
cipitate forms, which consists of a mixture of magnesium car- 
bonate and hydroxide. If the precipitate be dried at low 
temperature, we obtain a white, voluminous powder, whose 
composition generally corresponds to the formula Mg(OH)2, 
4C03Mg -f 4H20. This salt is employed under the name 
Magnesia alba in medicine. 
If it be suspended in water, and carbon dioxide passed 
through it, the salt will dissolve, and upon standing exposed to 
the air, crystals of neutral carbonate, MgC03 -(- 3H20, separate. 
When these are boiled with water they give up carbon dioxide 
and are again converted into the basic carbonate. The natu- 
rally occurring magnesite sustains no change when boiled. 
Magnesium carbonate yields isomorphous double salts, with 
potassium and ammonium carbonate; e.g., MgC03,K2C03 -f 
4h2o. 
Of the silicates of magnesium, we may mention olivine 
(Mg2Si04), serpentine (Mg3Si207 -f 2H20), talc (Si5014Mg4), 
sepiolite (Si308Mg2 -f 2H20), or meerschaum. The mixed sili- 
cates of magnesium and calcium are very numerous; to these 
belong asbestos, the augites and hornblendes. 
Recognition of Magnesium Compounds.—The fixed 
alkaline hydrates precipitate magnesium hydroxide from mag- 
nesium salts; the carbonates throw down basic magnesium car- 
bonate. The precipitates are insoluble in pure water and the 
alkalies, but dissolve readily in solutions of ammonium salts. In 
the presence of the latter, neither the alkaline hydrates nor car- 
bonates cause precipitation. In presence of ammonium salts, 316 
INORGANIC CHEMISTRY. 
disodium phosphate precipitates magnesium-ammonium phos- 
phate, MgNH4P04 -f 6H30, insoluble in water. 
BERYLLIUM. 
Be = 9.1. 
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 attaches 
itself in many respects to magnesium, so does beryllium approach alum- 
inium. Like, the latter, it is scarcely at all attacked by nitric acid, but 
dissolves easily in sodium or potassium hydrate, with elimination of hydro- 
gen. Like aluminium oxide, that of beryllium dissolves in the alkalies, 
and is almost invariably accompanied by the former in its natural com- 
pounds. Beryllium sulphate, like that of aluminium, forms a difficultly 
soluble double salt with potassium sulphate. However, beryllium, in most 
of its compounds, stands nearer to magnesium than to aluminium. 
Beryllium is not very abundant in nature and is found principally in beryl, 
a double silicate of aluminium and beryllium—Al2Be3(SiOs)6. Emerald 
has the same composition, and is only colored green by a slight amount 
of chronfium oxide. 
Metallic beryllium is obtained by the ignition of the chloride with 
sodium, and is a white ductile metal, of specific gravity 1.64. Its specific 
heat equals 0.4084; the atomic heat is, therefore, 3.8 (p. 255). It does 
not decompose water, even upon boiling. It does not oxidize in the air 
at ordinary temperatures. When finely divided it will burn in the air with 
a bright light when heated. In a compact mass (like magnesium), it does 
not do this. It is readily dissolved by dilute hydrochloric and sulphuric 
acids; also by potassium and sodium hydrates. 
Beryllium Chloride—BeCl2—is obtained, like aluminium chloride, by 
the ignition of a mixture of beryllium oxide and carbon in a stream of 
chlorine. It sublimes in shining needles, which deliquesce in the air It 
crystallizes from aqueous solution with four molecules of H2O; upon dry- 
ing it suffers a decomposition, similar that of magnesium chloride. 
The salts of beryllium have a sweet" taste, hence it has been called 
glucinum. Ammonium hydrate precipitates a white, gelatinous beryllium 
hydrate, Be(OH)2, from solutions of the soluble salts. T his dissolves 
readily in sodium and potassium hydrate, but on boiling, separates again 
from solution. When heatfcd, the hydroxide breaks up into water and 
beryllium oxide, BeO, which is a white, amorphous powder, of specific 
gravity 3.06. Its specific heat equals 0.2471. 
Beryilium Sulphate—BeS04—crystallizes from water at various 
temperatures, with four or seven molecules of H2O, of which one is rather 
closely combined. It crystallizes with magnesium sulphate in an isomor- 
phoue mixture. The double salt, S04Be, S04K.2, -J- does not dis- 
solve readily in water; in this respect it resembles the alums. ZINC HYDROXIDE ZINC OXIDE. 
317 
Zn = 64.9. 
ZINC. 
The natural compounds of the heavy metals have generally 
a high specific gravity, frequently possess metallic lustre, usu- 
ally occur 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 carbonate or 
sulphide is converted into oxide by roasting in the air; the pro- 
duct is then mixed with carbon and ignited in cylindrical clay 
tubes. In this manner the oxide is reduced: 
ZnO + C = Zn -f GO, 
and the liberated zinc distilled off. The receivers contain the 
fused, compact zinc and 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 re- 
ducing 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 4120 and distils about 
iooo°. 
It becomes coated with a thin layer of basic carbonate in 
moist air. Heated in the air it burns to zinc oxide with a 
very intense, bluish-white light. Compact zinc will only de- 
compose water at a red heat; zinc dust, however, acts at ordi- 
nary temperatures. Zinc is readily soluble in dilute acids, and 
dissolves with liberation of hydrogen in potassium or sodium 
hydrate, as well as in ammonia, when the solutions are boiled. 
Owing to its slight alteration in the air zinc meets with ex- 
tensive application as sheet-zinc for coating statues and in 
architectural adornment, and in galvanizing sheet-iron. It 
also forms an important constituent of many valuable alloys, 
such as brass and argentan (see these). 
Zinc Hydroxide—Zn(OH)2—is precipitated as a white 
amorphous 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 318 
INORGANIC CHEMISTRY. 
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 ac- 
quires a yellow color, which disappears on cooling. 
Zinc oxide occurs in nature as zincite, colored by im- 
purities. 
Zinc Chloride—ZnCl2—anhydrous, is obtained by heat- 
ing zinc in a stream of chlorine, by the evaporation of the 
solution of zinc in hydrochloric acid, and by the distillation 
of zinc sulphate with calcium chloride. It forms a white, 
deliquescent mass, which fuses when heated and distils about 
68o°. When the aqueous solution of zinc chloride is evapor- 
ated it 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 ob- 
tained, which hardens rapidly; a mixture of magnesium chlo- 
ride and oxide does the same. In both instances the harden- 
ing depends upon the formation of basic oxy-chloride, e.g., 
ZnClOH. Zinc chloride forms deliquescent double salts with 
the alkaline chlorides, e.g., ZnCl2,2KCl. With ammonia it 
yields various compounds, of which ZnCl2,NH3 is charac- 
terized by great stability. 
Zinc Sulphate—ZnS04—is obtained 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. It crystallizes at ordinary temperatures from aque • 
ous solutions with 7 molecules of H20 (zinc or white vitriol) 
in rhombic crystals, resembling those of magnesium sulphate. 
It affords double salts with the alkaline sulphates; these con- 
tain 6 molecules of water (p. 304). 
Zinc Carbonate—ZnC03—occurs native as smithsonite 
in hexagonal crystals. Sodium carbonate precipitates basic 
carbonates 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 pre- 
cipitates it as a white compound, from zinc solutions. Although 
fused zinc reacts with difficulty with sulphur, zinc dust combines 
with the latter in powdered form quite readily, and if the mix- 
ture be heated or struck with a hammer the union is accompanied 
by an explosion. Zinc sulphide is insoluble in water, but is 
readily dissolved by dilute acids, excepting acetic ; therefore it 
may be precipitated by hydrogen sulphide from zinc acetate CADMIUM. 
319 
solutions. This reaction serves to separate zinc from other 
metals. 
Zinc Silicate—Zn2Si04 4- H20—occurs in rhombohedral 
crystals as calamine. 
CADMIUM. 
Cadmium very often accompanies zinc in its ores. As 
much as 5 per cent, of this metal is present in the Silesian 
zinc ores; it was first discovered in these in 1819. Being 
more volatile than zinc, in obtaining the latter it distils off 
first, and may be easily separated from the first portions of the 
distillate. It is a white, tenacious, and rather soft metal, of 
specific gravity 8.6. It fuses at 3150, and boils at 770°. It 
does not alter much in the air. Heated, it burns with the 
separation of a brown smoke of cadmium oxide. It is diffi- 
cultly soluble in dilute hydrochloric and sulphuric acids, but 
dissolves readily in nitric. Zinc throws out the metal from 
solutions of the soluble cadmium salts. 
Cd — 111.9. 
St. Claire Deville found the specific gravity of cadmium vapors (at 
1040°) to be3.9 (air = -i) or 56 (H = 1). Therefore, the molecular weight 
of cadmium is 112. Since the atomic weight of cadmium (determined 
from its specific heat) is also 112, it follows that the gas molecule of cad- 
mium consists of but one atom. We know that the molecules of other ele- 
ments in the gaseous state are composed of two or more atoms (02,N2,P4,S6)- 
Cadmium, therefore, forms an exception to this rule. This is also true of 
mercury, and perhaps, too, of other divalent metals, such as zinc. These 
relations remind us of the behavior of the hydrocarbon residues (radicals); 
while the divalent or tetravalent groups, e.g., ethylene C2H4 and acety- 
lene C2H2, exist in free condition, the monavalent groups (as CHS,CN) 
cannot appear free, but double themselves, when separated from their 
compounds. 
Of the cadmium compounds may be mentioned : 
Cadmium Hydroxide—Cd(OH)2—is precipitated as a 
white powder, from the soluble cadmium salts, by the alkalies; 
it is insoluble in sodium and potassium hydrates, but dissolves 
readily in ammonium hydrate. 
Cadmium Oxide—CdO—is prepared by igniting the 
nitrate. It is a brownish-black powder, consisting of micro- 
scopic octahedra. 
Cadmium Chloride—CdCl.2—crystallizes from aqueous 
solution, with two molecules of H20, and may be dried with- 
out decomposition. The anhydrous salt melts at 5410 and 
sublimes in scales. 320 
INORGANIC CHEMISTRY. 
Cadmium Iodide—Cdl2—is obtained by the direct action 
of iodine upon metallic cadmium in the presence of water. 
It crystallizes from the latter in hexagonal tables. It is used in 
photography. 
Cadmium Sulphate—CdS04—crystallizes from water, 
not like the sulphates of zinc and magnesium, with 7 mole- 
cules of H20, but with 8/3H20; the crystals effloresce in the 
air. It, however, forms double salts with the sulphates of the 
alkali metals, e.g., CdS04K2S04 -f 6H20; these are perfectly 
analogous to those of zinc and magnesium, and isomorphous 
with them (p. 314). 
Cadmium Sulphide—CdS—occurs native as greenockite, 
in yellow hexagonal prisms, hydrogen sulphide precipitates 
it from cadmium salt solutions as a yellow powder, insoluble 
in dilute acids. It is employed as a pigment. 
Almost all the alloys of cadmium have a low fusion tem- 
perature. Freshly prepared cadmium amalgam is a white 
plastic mass, which soon becomes hard. It is used in filling 
teeth. 
The chemical energy of cadmium is less than that of zinc; this is evi- 
dent from the fact that the former may be displaced from its salts by the 
latter. We saw that, with the elements of the groups of potassium and 
calcium, the chemical energy increases inversely with the increasing 
atomic weight; caesium is more energetic than rubidium, barium more 
than calcium. It is worthy of remark that nearly all of the more electro- 
negative elements belonging to the second sub-groups of the seven main 
groups of the periodic system exhibit a diminution in chemical energy 
with rising atomic weight similar to that shown by the members of the 
magnesium group; copper displaces silver; phosphorus is more energetic 
than arsenic and antimony; sulphur more energetic than selenium and 
silver; chlorine sets free or displaces bromine and iodine. 
These relations of affinity find full expression in the thermo-chemical 
phenomena in which are clearly shown the double periodicity of the great 
periods and the relations of the two sub-groups, Ca, Sr, Ba and Zn, Cd, Hg, 
to magnesium. The basic character increasing from Mg to Ba corre- 
sponds to the increase in heat developed by the formation of their com- 
pounds, e.g., the chlorides, hydroxides, and sulphydrates. 
(Mg,Cl2) = 151.0. 
(Ca,Cl2) = 170.2. 
(Sr,Cl2) = 184.5. 
(Ba,Cl2) = 194.5. 
fMg,0,H20) — 148.9. 
?Ca,0,Aq.) 2= 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.) = X07.1. 
That the increase with the hydrates is so slight is explained, probably by 
the decreasing solubility of the same from Ba to Mg, inasmuch as an evo- 
lution of heat (heat of precipitation) corresponds to the difficult solubility. MERCURY. 
321 
The heat of formation of the carbonates (from the metallic oxides and car- 
bon dioxide) must also be introduced here: 
(Ca0,C02) = 42.5 (Sr0,C02) = 53.2 (BaO,C02) = 55:9. 
These seem to indicate that calcium carbonate is less stable and more 
easily decomposed than barium carbonate (p. 299). 
The series Mg, Zn, Cd, Hg deports itself differently. In this the heat 
disengagement becomes successively less and corresponds with the dimin- 
ishing basicity. 
(Mg,Cl„) = 151.0. 
(Zn,CL) = 97.2. 
(CdCL) = 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 3. 
(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 towards this acid. All metals 
liberating a greater quantity of heat than 39-3 C. in the formation of their 
chlorides (calculated for 1 equivalent of metal) are in condition to decom- 
pose the dilute acid. Most of the metals belong to this class; mercury, 
copper, silver, gold, lead, thallium, and some others, set free a less amount 
of heat, and hence are not able to decompose dilute hydrochloric acid (see 
p. 265). 
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 precipitated by hydrogen 
sulphide from their chlorides, etc. 
If in the thermo-chemical equation, 
the sum of the heat developed upon the right side is greater*than that 
upon the left, the reaction will occur (precipitation of metallic sulphides) ; 
in the opposite case the sulphide is decomposed by the dilute hydrochloric 
acid. 
(Me, Cl2.Aq.) -f (S,H2Aq.) = (Me, S) + 2(H, Cl, Aq.), 
The magnitude of the atomic weight of Mercury would 
place the latter in the group of zinc and cadmium. The 
relationship of these three heavy metals is observed in many 
similarities of the free elements and of their compounds (p. 322). 
Occupying a similar position in the three great periods (241) 
they are distinguished among the heterologous members in a 
physical point of view by their ready fusibility and volatility, 
which nearly reach a maximum in them. In the homologous 
series, Zn, Cd, Hg, these properties, like the specific gravities, 
increase with rising atomic weight (just as with the metals of 
the potassium group, p. 273). 322 
INORGANIC CHEMISTRY. 
Zn 
Cd 
Hg 
Atomic weight 
64.9 
111.9 
199.8 
Fusing point 
412° 
315° 
—40° 
Boiling point 
940° 
765° 
360° 
Specific gravity 
7-1 
8.6 
13.6 
The gradation in the heat of formation of their compounds 
(p. 321) clearly indicates that mercury must be arranged in a 
group with cadmium and zinc. 
Like zinc and cadmium, it yields compounds of the form, 
HgX2, in which it appears divalent. These derivatives are, in 
many respects, 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 + 6H20), and are isomorphous with 
the double sulphates of the magnesium group (p. 314). The 
similarity, however, limits itself to few compounds. Since the 
properties of each group sustain a slight change in virtue of 
the increasing atomic weight, we are not- surprised to observe 
this to be very evident in the case of Hg (with- the high atomic 
weight 199), especially as the middle (transition) member of the 
third great period is not known (p. 241). Mercury differs es- 
sentially from zinc and cadmium in that, in addition to the 
11 
compounds of the form HgX2 (mercuric compounds), it is also 
capable <jf yielding those of the form HgX (mercurous com- 
pounds), in which it seems to be monovalent. Here we meet 
an instance, frequently observed, in which one and the same 
metal (as with the most metalloids) is capable of forming com- 
pounds of two or more forms, which are to be referred to a 
different valence of the metal; and it often happens that the 
derivatives of a metal, appearing in different forms or types, 
are frequently more essentially distinguished from one another 
than the compounds of different elements having the same 
type. Thus, the mercuric compounds (HgX2) are similar to 
those of zinc and cadmium, after the same form, while the 
i 
mercurous compounds (HgX) exhibit great resemblance to the 
i i 
cuprous (CuX) and silver (AgX) compounds, constituted 
according to a similar type. 
It shows that the similarity of the compounds is not only MERCURY. 
323 
influenced by the nature of the metals, but frequently, to a 
marked degree, by the forms or types according to which they 
are constituted (p. 330). 
As viewed above, mercury in its ic compounds is a dyad, in the ous a 
monad. According to the theory of constant valence, the mercury atom 
in the ous compounds is, however, also divalent. We suppose that the 
molecules of the same are twice as large, and that in them two Hg atoms 
form a divalent group, as seen from the following: 
Hgx 
I >0 
Hg/ 
Hg — Cl 
I 
Hg — Cl 
Hg - N03 
I 
Hg - N03 
Hg\ 
I >s 
HgX 
An experimental decision upon the above has not yet been given (p. 325). 
Mercurous oxide. 
Chloride. 
Nitrate. 
Sulphide. 
MERCURY. 
Hg = 199.8. 
Mercury (Hydrargyrum) occurs in nature principally as Cin- 
nabar, more rarely native in the form of little drops scattered 
through rocks. Its most important localities are Almaden in 
Spain, New Almaden in California, Idria in Illyria, Mexico, 
Peru, China, and Japan. 
The metallurgical separation of mercury is very simple. 
Cinnabar is roasted in reverberatory furnaces, whereby the 
sulphur burns to dioxide, and the mercury vapors are con- 
densed in large chambers. Or, it is distilled with lime or iron 
from iron retorts. Commercial mercury usually contains a 
slight quantity of other metals dissolved-in it. For its purifi- 
cation, it is poured in a thin stream into a deep layer of sul- 
phuric or dilute nitric acid, by which the accompanying tin 
and lead are more easily dissolved than the mercury. The 
metal is finally distilled out of a small glass retort and pressed 
through chamois skin. 
This is the only metal which is liquid at ordinary tempera- 
tures. At o° its specific gravity equals 13.59; it solidifies at 
—40°, and crystallizes in regular octahedra ; it evaporates 
somewhat at medium temperatures, and boils at 360°. Its vapors 
are very poisonous. The specific gravity of the vapor of mer- 
cury is 99.9 (H = 1) or 6.91 (air =1). Therefore, the molec- 
ular weight of the metal is 199.8, and as its atomic weight is 
also 199.8, the molecule, like that of cadmium, is composed of 
only one atom. At ordinary temperatures, mercury is not 
altered by exposure to the air; near the boiling point, however, 324 
INORGANIC CHEMISTRY. 
it gradually oxidizes to red mercuric oxide. Hydrochloric and 
cold sulphuric acids do not act upon it; hot sulphuric acid 
converts it into mercury sulphate, with evolution of sulphur 
dioxide. Even dilute nitric acid will readily dissolve it. It 
combines with the halogens and sulphur at ordinary tempera- 
tures. 
Mercury dissolves almost all metals (not iron) forming amal- 
gams. It unites with potassium and sodium upon gentle 
warming, with production of heat and light. When the quan- 
tity of potassium and sodium exceeds 3 per cent., 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 mer- 
curic. The first are analogous to the cuprous, and have the form, 
HgX. In them mercury appears to be monovalent; we, how- 
ever, do not know, whether their molecules are not to be ex- 
pressed by the double formula Hg2X2 (p. 323). In many re- 
spects 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 divalent, and is 
very much like zinc and cadmium. The ic compounds almost 
always form, if the substance reacting with the mercury 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(NOs)2 4- Hg — Hg/N03)2. Oxidizing agents con- 
vert the ous into the ic compounds; the latter are, on the other 
hand, converted by reducing substances into the first. 
The heat of formation of some of the mercuric compounds corresponds 
to the symbols: 
(Hg,0) = 30.6 (Hg,Cl2) = 63.1 (Hg,I2) = 34-3 (Hg,S) = 16.8. 
That of the corresponding mercurous salts; 
(Hg2,0) = 42-2 (Hg,Cl) 41.2 (Hg,I) = 24.2 (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. It is generally formed by the sublimation of 325 
MERCURY. 
HgCl2 with mercury; or a mixture of HgS04, mercury and 
sodium chloride is sublimed : » 
HgS04 + 2NaCl -f- Hg = Na2S04 + Hg2Cl2. 
It then forms 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. When 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 (hence the name 
calomel, from xakoiieXaq), and reacts according to the equa- 
tion : 
Hg2Cl2 + 2NH3 = NH4C1 -f NH2Hg2Cl. 
The compound NH2Hg2Cl is viewed as ammonium chloride, 
in which 2H are replaced by Hg2, 
The vapor density of calomel vapors at 440° is 117.6 (H = 1), the 
molecular weight, therefore, 235.2, and corresponds to the formula HgCl 
(235.2). It appears, however, that its vapors consist of a mixture of mer- 
cury and mercuric chloride. Such a mixture must have the same density 
as HgCl: 
HgCl + HgCl = Hg + HgCl2. 
x vol. 
I vol. 
I vol. 
I 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 determina- 
tion of its vapor density. 
Mercurous Iodide—Hgl or Hg2I2—is prepared by rub- 
bing together 8 parts of mercury with 5 parts I, or by precipi- 
tating mercurous nitrate with potassium iodide. It is a greenish 
powder, insoluble in water and alcohol. Light changes it to 
Hgl2 and Hg. 
Mercurous Oxide—Hg20—is black in color, and is formed 
by the action of potassium or sodium hydrate upon mercurous 
salts. In the light or at ioo°, it decomposes into HgO and Hg. 
Mercurous Nitrate—HgN03 or Hg2(N03)2—is produced 
by allowing somewhat dilute nitric acid to act upon excess of 
mercury in the cold. It crystallizes with 1 molecule H20 in 
large monoclinic tables. It dissolves readily in water acidu- 
lated with nitric acid ; pure water decomposes it into the acid 
salt which passes into solution, and the basic salt—Hg , 
which separates as a yellow powder. 
The nitric acid solution of mercurous nitrate oxidizes when 326 
INORGANIC CHEMISTRY. 
exposed to the air, and gradually becomes mercuric nitrate; 
this be prevented by adding metallic mercury to the so- 
lution, whereby the resultant ic salt is again changed to the 
ous state : 
Mercurous Sulphate—Hg2(S04)—results when an excess 
of mercury is heated gently with sulphuric acid ; it separates as 
a crystalline precipitate, difficultly soluble in water, if sulphuric 
acid be added to a mercurous nitrate solution. It fuses upon 
application of heat, and decomposes into S20, 02, and Hg. 
Mercurous Sulphide—Hg2S—is precipitated by potas- 
sium hydrosulphide, as a black compound, from the dilute so- 
lution of mercurous nitrate. When gently warmed, it decom- 
poses into HgS and mercury. 
Hg(N03)2 + Hg = Hg2(N03)2. 
MERCURIC COMPOUNDS. 
Mercuric Chloride—HgCl2—Corrosive sublimate—is 
produced when mercuric oxide is dissolved in HC1, or metallic 
mercury in aqua regia. It is obtained on a large scale by the 
sublimation of a mixture of mercuric sulphate with sodium 
chloride: 
It crystallizes from water in fine rhombic prisms, and dissolves 
at medium temperatures in 15 parts, at ioo°, in 2 parts water; 
it is still more soluble in alcohol. Its specific gravity is 5.4. 
It fuses at 288°, and boils about 300°. Its critical pressure 
is about 420 M111, (p. 88). The vapor density is 135.2 
(H= 1), corresponding to the molecular formula HgCl2 (= 
270.5): 
HgSO, + 2NaCl = HgCl2 + Na2S04. 
Hg + Cl, = HgCl2. 
Reducing substances, like S02 and SnCl2, change it to in- 
soluble mercurous chloride: 
I vol. 
I vol. 
I vol. 
2HgCl2 + S02 + 2H20 = Hg2Cl2 + H2SOi + 2HCI. 
Stannous chloride first precipitates mercurous chloride: 
2HgCl2 -(- SnCl2 = Hg2Cl2 -f- SnCl4, which is afterwards re- 
duced, by excess of the first, to metallic mercury: Hg2Cl2 + 
SnCla = 2Hg + SnClr 
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 heavy white pre- MERCURY. 
327 
cipitate, called white precipitate, NH2HgCl, is thrown down. 
This compound is regarded as a derivative of ammonium chlo- 
ride, in which two atoms of H are replaced by a divalent 
mercury atom, and it has been called Mercur-ammonium 
Chloride. It forms the compound NH2HgClNH4Cl with 
ammonium chloride; the structure of this is expressed by the 
formula: 
„ /NH„C1 
HS\NH3C1. 
Similar mercur-ammonium derivatives are numerous. 
Mercuric Iodide—Hgl2—is formed by the direct union 
of mercury with iodine. When potassium iodide is added to 
a solution of mercuric chloride, Hgl2 separates as a yellow pre- 
cipitate, which immediately becomes red. Hgl2 is readily sol- 
uble in HgCl2 and KI solutions ; it crystallizes from alcohol 
in bright red quadratic octahedra. Upon warming Hgl2 to 
150°, it suddenly becomes yellow, fuses and sublimes in yel- 
low, 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. Mercuric iodide is 
therefore dimorphous. 
Mercuric Oxide—HgO—is obtained by the prolonged 
heating of metallic mercury near the boiling point in the air, or 
by the ignition of mercurous or mercuric nitrate. It forms a 
red, crystalline powder, of specific gravity 11.2. When sodium 
hydrate is added to a solution of mercuric chloride, mercuric 
oxide separates as a yellow, amorphous precipitate. Both 
modifications become black when heated, but change to a 
yellowish-red on cooling. Mercuric oxide breaks up into 
mercury and oxygen about 400°. 
Mercuric oxide combines directly with ammonia, to form the 
compound 2HgO.NH3, which explodes with violence when 
heated. 
Mercuric Nitrate—Hg(N03)2. It is difficult to obtain 
this salt pure, because it is inclined to form basic compounds. 
A solution of it may be made by dissolving mercury or mer- 
curic oxide in an excess of hot nitric acid. On diluting the 
solution with water the basic salt, Hg(N03)2,2Hg0,H20, sepa- 
rates, and this may be converted into pure mercuric oxide by 
boiling with water. 
Mercuric Sulphate—HgS04—is produced by digesting 
mercury or its oxide with an excess of concentrated sulphuric 
acid. It forms a white, crystalline insoluble mass, which be- 
comes yellow on heating. It yields the hydrate HgS04 H20 328 
INORGANIC CHEMISTRY. 
with a little water, but much of the latter decomposes it into 
sulphuric acid and the yellow insoluble basic salt, HgS04.2HgO 
( Turpetum miner ale, Turpeth mineral). 
Mercuric sulphate forms double salts with the alkaline sul- 
phates, e.g., HgSQ4, K2S04 -f- 6H20; these are isomorphous 
with the corresponding double salts of the magnesium group 
(P- 3i4)- 
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 flow- 
ers of sulphur with water, or it is produced as a black amorphous 
mass by the precipitation of a solution of a mercuric Salt with 
hydrogen sulphide. If the black sulphide be heated with 
exclusion of air it sublimes as a dark red mass of radiating 
crystalline structure, and is perfectly similar to natural cinna- 
bar. A similar conversion of the black modification into the 
red is effected by continued heating of the same to 50° with 
a solution of potassium or ammonium sulphide. The red 
mercury sulphide thus obtained is employed as artificial cinna- 
bar in painting. 
The mercury compounds can be readily recognized by the 
following reactions. On fusion with dry sodium carbonate, 
mercury escapes, and (if the operation be executed in a small 
tube) condenses upon the side 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, mercury 
is deposited as a gray coating, which on being rubbed acquires 
a metallic lustre. The mercurous compounds are distin- 
guished from the mercuric by their precipitation by hydro- 
chloric acid. 
COPPER, SILVER, AND GOLD. 
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 bear to magnesium: 
Na = 23 
Cu = 63.2 
Ag = 107.6 
Au = 196.2 
Mg = 23.9 
Zn = 64.9 
Cd =111.9 
Hg = 199.8 
They occupy an entirely analogous position in the three great 
periods of the periodic system of the elements (p. 244), and 
constitute the transition from the elements of group VIII, espe- COPPER, SILVER, AND GOLD. 
329 
cially from nickel, palladiurp, and platinum, to the less basic 
elements of group II—zinc, cadmium, and mercury : 
Ni = 58.6 
Pd = 106.2 
Pt = 194.3 
Cu = 63.2 
Ag == 107.6 
Au = 196.2 
Zn = 64.9 
Cd =111.9 
Hg = 199.8 
This intermediate position of the three elements about to be 
discussed is clearly shown in their entire physical deportment. 
While the elements of group VIII, with the last members, Ni, 
Pd, and Pt, fuse with difficulty and do not volatilize, Cu, Ag, 
Au, in point of fusion and volatility, constitute the transition 
to the readily fusible and volatile elements, Zn, Cd, and Hg. 
They take an intermediate position, too, 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. 
Not only are the properties of the free elements determined 
by the position of the latter in the periodic system, but those 
of their derivatives, and especially such as depend upon the 
valence of the elements, are influenced to a marked degree by 
the above relation. In consequence of the double periodicity 
of the great periods, Cu, Ag, and Au attach themselves to 
group I, and especially to sodium, just as the elements imme- 
diately following, Zn, Cd, and Hg, arrange themselves with 
group II and magnesium. Hence we find Cu, Ag, and Au, 
like Na, yielding compounds of the form MeX, in which they 
appear monovalent. Some of these are isomorphous; thus 
NaCl, CuCl, and AgCl crystallize in forms of the regular sys- 
tem. Silver sulphate, Ag3S04, is isomorphous with sodium 
sulphate, S04Na2; and the same is true of other salts of these 
two metals. Cu and Ag, like the alkalies, afford • so-called 
sub- or quadrant oxides, Na40, Cu40, Ag40. 
•But we may say that the similarity of Cu, Ag, and Au to Na 
is confined to these few external properties. Just as the heavy 
metals, Zn, Cd, and Hg differ in many properties from the light 
metal magnesium (p. 311), so do the metals Cu, Ag, and Au, 
possessing a high specific gravity distinguish themselves in a still 
higher degree from the light metal sodium. They possess all the 
properties belonging to the heavy metals, which are mainly char- 
acterized by the insolubility of the oxides, sulphides, and many 
salts. This character which separates them from sodium is ex- 
plained by the fact that they really belong to the three great 
periods, and are classified with the alkali metals in but few 330 
INORGANIC CHEMISTRY. 
properties. Gold, with the high atomic weight, 196.2, corres- 
ponds in this respect to mercury (p. 322), but is very variable. 
I 
In the compounds constituted according to the form, MeX, 
in which Cu, Ag, and Au appear monovalent, they exhibit great 
similarity as regards their physical and chemical properties. 
The chlorides, CuCl, AgCl, and AuCl, are colorless and inso- 
luble in water; soluble, however, in hydrochloric acid, am- 
monia, the alkaline hyposulphites, etc., and furnish perfectly 
similar double compounds. While silver only enters com- 
I 
pounds of the form AgX, copper and gold are capable of 
yielding another form; copper forms, besides cuprous, 
1 11 . 
CuX, also cupric, CuX2, derivatives, in which it appears to be 
divalent. The latter are much more stable than the former, and 
embrace the most usual copper salts. Gold, however, besides 
1 111 
furnishing ous, AuX, compounds, has ic derivatives, AuXs, in 
which it appears trivalent. 
While Cu and Au, in their ous forms, are analogous to silver 
(and in less degree, Na), the cupnV derivatives show a great 
resemblance to the compounds of the metals of the magnesium 
group, and other metals in their divalent combinations. Thus, 
the sulphates of zinc, magnesium, cupric oxide (CuO), ferrous 
oxide (FeO), nickelous oxide (NiO), cobaltous oxide (CoO), 
and manganous oxide (MnO), are - similarly constituted, 
resemble each other, are isomorphous, and form entirely anal- 
ogous double salts (p. 314) with the alkaline sulphates. In 
II 
the same way the carbonates (MeCOs), the chlorates and bro- 
11 
mates (MeCl206 + 6H20) and others, are similarly constituted 
and isomorphous. In its ic derivatives, gold exhibits some 
III 
similarity to the aluminium compounds (A1XS), to those of 
III 
indium (InX3) and other metals, in their trivalent combina- 
tions. Here we see, as already observed with mercury (p. 
322), that the similarity of the cottipounds of the metals is in- 
fluenced by the similarity of forms or types, according to which 
they are composed, i.e., by the valence of the metals. If a 
metal form several series'of compounds of different types, 
each series is usually more or less similar to the compounds of 
other metals of like type. In this manner is shown the re- 
semblance of the compounds of the following types : COPPER, SILVER, AND GOLD. 
331 
Na20 
Sodium oxide. 
Ag20 
Silver oxide. 
Cu20 
Cuprous oxide. 
Au20 
Aurous oxide. 
ti2o 
Thallous oxide. 
MgO 
Magnesium oxide. 
ZnO 
Zinc oxide. 
CuO 
Cupric oxide. 
FeO HgO 
Ferrous oxide. Mercuric oxide. 
ai2o3 
Aluminium oxide. 
Fe203 
Ferric oxide. 
Au203 
Auric oxide. 
Ti2o3 
Thallic oxide. 
The character of their derivatives varying with the degree of 
combination or valence, becomes quite marked with chromium, 
manganese, and iron, as we shall later see. The heavy metals 
also exhibit a strong, positive basic character in their mono- 
valent combinations. 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. The 
metalloidal character of the metals, and the acid nature of their 
oxides begin to appear in their trivalent combinations. Thus 
in the hydroxyl derivatives of aluminium, indium, and gold, 
Al(OH)3, In(OH)s, Au(OH)3, hydrogen may be replaced by 
the alkalies. Their higher forms of oxidation show, like those 
of the metalloids, a pronounced 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). The character of 
the compound is influenced in a less, if not an -unimportant 
degree, by the position of the elements in the periodic system. 
Hence the properties of the metallic compounds are not only 
influenced by the nature of the metals, but to a high degree by 
the combination forms. These forms of the elements, and par- 
ticularly those of the metals, are regulated, however, if not 
entirely, yet to a considerable degree, by the periodic system, 
as previously observed (see p. 247). 
This connection of Cu, Ag, and Au in an analogous group, expresses 
itself too, in the heat resulting from the formation of their compounds of 
the form MeX: 
(Na, Cl) = 97.6 (Na2, O) == 100.2 (Na2. S) = 88.0 
(Cu, Cl) = 32.8 (Cu2, O) = 40.8 (Cu2, S) = 20.2. 
(Ag, Cl) = 29.3 (Ag,,0) = 5.9 (Ag2, S) = 5.3. 
(Au, Cl) = 5.8 (Au2, O) = - (Au2, S) = 
Consequently relations occur here perfectly similar to those of the ele- 
ments of the zinc group (p. 320), and perfectly analogous conclusions as 
regards the affinity relations may be deduced from them. Thus, for 
example, copper is able to decompose concentrated but not dilute hydro- 
chloric acid. The heat of formation of some cupric compounds equals ; 
(Cu, O) = 37.1 (Cu, Cl2) == 51.6 (Cu, Cl2, Aq.) = 62.7 
(Cu, S, OJ = 182. 332 
INORGANIC CHEMISTRY. 
COPPER. 
Native copper is found in large quantities in America, China, 
Japan, also in Sweden and in the Urals. It frequently occurs 
crystallized in cubes and octahedra. The most important and 
most widely distributed of its ores are : cuprite (Cu20), mal- 
achite and azurite (basic carbonates), chalcocite (Cu2S), and 
especially chalcopyrite or speckled copper ore (CuFeS2). 
Cu = 63.2. 
Metallurgy of Copper.—The extraction of copper from its oxygen ores 
is very simple : metallic copper is melted out when the ores are ignited 
along with charcoal. The sulphur ores are more difficult to work. The 
divided material is first roasted in the air, by which means copper sul- 
phide is partly converted into oxide. The mass is afterward ignited with 
sand, silica fluxes, and carbon, when iron sulphide is converted into oxide 
and passes into the slag. By several repetitions of this process we get the 
so-called copper stone—a mixture of cupric sulphide with oxide. This is 
repeatedly roasted and heated, and metallic copper obtained by the action 
of the cupric oxide upon the sulphide: 
2CuO -f- CuS = 3CU -(- S02. 
The copper obtained in this way is fused again with charcoal, to free it 
from the oxide. 
To obtain chemically pure copper, the pure oxide is heated 
in a stream of hydrogen, or the solution of copper sulphate is 
decomposed by electrolysis. 
Metallic copper possesses a characteristic red color, and trans- 
mits a green light in thin leaflets. It is rather soft and ductile, and 
possesses a specific gravity 8.9. It fuses about 1054°, and vapor- 
izes in the oxy-hydrogen flame. It remains unaltered in dry air ; 
in moist, it is gradually coated with a green layer of copper 
carbonate. When heated, it oxidizes to black cupric oxide. 
Copper is not changed by dilute hydrochloric or sulphuric 
acids ; if it be moistened with these, and exposed to the air, it 
absorbs oxygen, and gradually dissolves. It is similarly dis- 
solved by ammonium hydrate. Concentrated sulphuric acid 
converts it into copper sulphate, with evolution of sulphur 
dioxide. It dissolves in dilute nitric acid in the cold, with 
evolution of nitric oxide. Zinc, iron and also phosphorus 
precipitate metallic copper from the aqueous solutions of its 
salts. 
Copper forms two series of compounds, known as cuprous 
and cupnV. In the ic compounds, copper is divalent: 
CuO CuCl2 Cu(OH)a S04Cu. COPPER. 
333 
These are more stable than the ous derivatives; the ordinary 
copper salts belong to them. In many respects they resemble the 
compounds of other dyad metals, especially those of the mag- 
nesium group, and ous compounds of iron (FeO), manganese 
(MnO), cobalt and nickel (see p. 330). 
The cuprous compounds are, on the other hand, very un- 
stable, absorb oxygen from the air, and pass into cupric deriv- 
atives. They show some similarity to the mercurous deriva- 
tives (p. 324), and possess an analogous composition : 
CuCl Cul Cu20 Cu2S. 
Oxygen salts of cuprous oxide are not known. 
From the formulas given above, copper, like silver, is monovalent in its 
ous compounds. It is, however, questionable, whether these formulas ex- 
press the real molecular values. It is ordinarily assumed that the cuprous 
derivatives, like those of mercury in its ous state (p. 325), correspond to 
the doubled formulas, and that the copper atom is divalent, and forms a 
divalent group composed of two copper atoms, as may be seen from the 
following formulas: 
CuCl 
Cu2Cl2 or | 
CuCl 
CuN 
Cu„0 or | No 
Cu/ 
The vapor density of cuprous chloride corresponds to the formula Cu2Cl2 
(compare p. 340), and, therefore, rather favors the above opinion. 
Cuprous Chloride. 
Cuprous Oxide. 
Cuprous Oxide—Cu20—occurs as cuprite crystallized in 
regular octahedra. It is obtained artificially by boiling a so- 
lution of copper sulphate and grape sugar with potassium hy- 
drate, when it separates as a crystalline, bright red powder. 
It does not change in the air, and is readily soluble in ammo- 
nium hydrate. The solution absorbs oxygen, and while form- 
ing cupric oxide acquires a blue color. By the action of sul- 
phuric and other oxygen acids, it forms cupric salts, the half 
of the copper separating as metal: 
CUPROUS COMPOUNDS. 
Cu20 -f SQ4H2 = CuS04 + Cu + H20. 
The hydrate, Cu2(OH)2, is precipitated by the alkalies as a 
yellow powder from hydrochloric acid solutions of Cu2Cl2. 
It oxidizes in the air 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 low, red 334 
INORGANIC CHEMISTRY. 
heat, by boiling the solution of cupric chloride with copper 
(CuCl2 4 Cu = Cu2Cl2), and by the action of many reducing 
substances upon cupric chloride. It is most conveniently 
made by passing sulphur dioxide through a concentrated solu- 
tion of copper sulphate and sodium chloride, when it separates 
as a white, shining powder, consisting of small tetrahedra. It 
fuses at 430°, and distils about iooo° ; its vapor density cor- 
responds to the formula Cu2C12. In the air, it rapidly be- 
comes green, owing to oxygen absorption, and the formation 
/Cl 
of basic cupric chloride, Cu/ Qjq- Cuprous 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. 
By extracting the co-precipitated iodine by means of ether 
it is obtained as a gray powder, insoluble in acids. 
Cuprous Sulphide—Cu2S—occurs as chalcocite crystal- 
lized in rhombic forms. It is produced by burning copper 
in vapor of sulphur, and by heating cupric sulphide in a current 
of hydrogen; after fusion it solidifies in crystals of the regular 
system. Combined with silver sulphide it constitutes the 
Cu ) 
mineral j- S or Cu2S.Ag2S, isomorphous with 
chalcocite. 
Copper Hydride—CuH or Cu,H2—belongs to the deriva- 
tives of monovalent copper. If a solution of copper sulphate 
be digested with hypophosphorous acid, the hydride separates 
as a yellow amorphous precipitate which soon acquires a brown 
color. At 6o° it decomposes into copper and hydrogen. 
With hydrochloric acid it forms cuprous chloride: 
Cull + HC1 = CuCl + Ha. 
Copper suboxide, Cu40, or quadrantoxide, corresponds to potassium 
suboxide. On adding an alkaline stannous chloride solution to one of copper 
sulphate there separates, at first, cupric hydroxide, which is further reduced 
to cuprous hydroxide, and then to suboxide. The latter is an olive-green 
powder, which oxidizes readily and is decomposed by H2S04 into CuS04 
and 3CU. CUPRIC COMPOUNDS. 
335 
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 tempera- 
tures, settles together and acquires a metallic lustre. By heat- 
ing 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 hydroxide dissolve in ammonium hydrate 
with dark blue color. The solution possesses the power of 
dissolving wood fibre (cotton-wool, linen, filter-paper, etc.)— 
Schweizer's reagent. * 
Cupric Chloride—CuCl2—is formed by the solution of 
cupric oxide or carbonate in hydrochloric acid. It crystallizes 
from aqueous solution, 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 an- 
hydrous chloride, which at a red heat is decomposed into 
chlorine and cuprous chloride. It yields beautifully crystallized 
double salts with potassium and ammonium chlorides. Cupric 
bromide is like the chloride; the iodide is not known, since in 
its formation it at once breaks up into cuprous iodide and iodine. 
Copper Sulphate—CuS04+5H20—cupric sulphate, cop- 
per vitriol—may be obtained by the solution of copper in 
concentrated sulphuric acid. It is produced on a large scale 
by roasting chalcocite. It forms large blue crystals of the 
triclinic system, which effloresce somewhat upon exposure. At 
ioo° the salt loses 4 molecules of water; the fifth separates 
above 200°. The anhydrous sulphate is colorless, absorbs 
water very energetically, and returns to the blue hydrous com- 
pound. 
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 ammonium 
sulphates, which crystallize with 6H20, and are isomorphous 
with the double salts of the metals of the magnesium group. 336 
INORGANIC CHEMISTRY. 
Copper sulphate is employed in electro-plating. When its so- 
lution is decomposed by the galvanic current copper separates 
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 produced at 
first, changes the color of the liquid to a dark blue. From 
this solution alcohol precipitates a dark-blue crystalline mass 
with the composition CuS04, 4NH,"-}- H20. Heated to 150° 
this compound loses water and 2 molecules of NH3, and be- 
comes 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. : 
/NH 
SO/ >Cu 
xnh/ 
The other soluble copper salts afford 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. Heat converts it into 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, CuCCh.CufOHT, or 
rn/O.Cu.OH . T. 
CO/q qjj, separates as a green precipitate. It occurs in 
nature as ?nalachite, 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. It was formerly employed as a pigment, 
under the name of Scheele's green, but at present, owing to 
its poisonous character, it has been replaced by other green 
colors (Guignet’s green and aniline green). 
Cupric Sulphide—CuS—is a black compound, precipi- 
tated from copper solutions by hydrogen sulphide. It is in- 
soluble 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. 
Cuprammonium sulphate. SILVER. 
337 
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 it contracts unequally 
upon cooling 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 im- 
portant copper alloys are: 
Brass, consisting of two to three parts copper and one part 
zinc. It has a yellow color, and is considerably harder than 
pure copper. Ordinarily, one to two per cent, of lead are 
added to the brass, which facilitates its working upon the 
turning-lathe. Tombac contains 15 per cent, 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 bronzes. Most of the modern 
bronzes also contain zinc and lead; those from Japan, gold 
and silver. The cannon bronze contains 90 per cent, copper 
and 10 per cent, tin; bell metal has 20-25 per cent, of tin. 
Argentan is an alloy of copper, zinc, and nickel (see latter). 
The so-called Talmi gold consists of 90-95 per cent. Cu and 
5-10 per cent, aluminium. The German copper coins consist 
of 95 per cent. Cu, 4 per cent. Sn, and 1 per cent. 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 in- 
troduced into a copper solution, it becomes covered with a red 
layer of metallic copper. Volatile copper compounds tinge the 
flame blue or green. The spectrum of such a flame is char- 
acterized by several blue and green lines. 
SILVER. 
Silver 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 
Ag = 107.66. 338 
INORGANIC CHEMISTRY. 
America (Chili, Mexico, California), Saxony (Freiberg), Hun- 
gary, the Altai and Nertschinsk. 
Metallurgy of Silver.—The separation of the metal from its ores is 
rather 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, there is ob- 
tained a copper stone consisting of iron, copper and silver sulphides. This 
is then oxidized in a furnace; from the resulting mixture of ferric and 
cupric oxides and silver sulphate (S04Ag2), the latter is extracted by water. 
The silver is precipitated from this aqueous solution by copper. 
Formerly, in Saxony, the separation of the silver was executed accord- 
ing to the so-called amalgamation process. According to this the mix- 
ture of sulphides is roasted with sodium chloride, whereby silver chloride 
is produced. The divided material is then mixed with iron scraps and 
water in rotating vessels. The iron causes the precipitation of the metallic 
silver from its chloride : 
2AgCl -J- Fe = FeClj -{- 2Ag. 
To free the metal from various impurities it is dissolved in mercury and 
the liquid amalgam ignited; mercury distils off and silver remains. 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 podium 
chloride and copper sulphate in the presence of water. In this way cuprous 
chloride is produced, which is transposed, with silver sulphide, into silver 
chloride and cuprous sulphide : 
2CuCl -f" Ag2S — Cu.2S 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 re- 
moved by sieves, while a readily fusible alloy of lead with more silver 
remains behind. This method of Pattison’s is repeated until the residual 
liquid lead contains 1 per cent, of silver. The lead, rich in silver, is sub- 
jected to cupellation : it is fused in a reverberatory furnace with air access. 
The bottom of the oven consists 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 copper and traces of other metals in 
greater or less quantity. To prepare chemically pure metal, 
the work silver is dissolved in nitric acid, and from the solu- 
tion of the nitrates thus obtained, hydrochloric acid precipi- 
tates the silver as chloride : 
AgNOs + HC1 = AgCl + HNOs. SILVER. 
339 
The latter is reduced by various methods; either by fusion 
with sodium carbonate, or by the action of zinc or iron in the 
presence of water: 
2AgCl -(- Zn = ZnCl2 -)- 2Ag. 
Silver is a pure white, brilliant metal, of specific gravity 
10.5. It is tolerably soft and ve*ry ductile, and can be drawn 
out to a fine wire. It crystallizes in regular octahedra. It 
fuses about 9540, and is converted into a greenish vapor in the 
oxy-hydrogen flame. 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 gas 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: 
2Ag 4- 2H2SOi = Ag2S04 + S02 + 2H20. 
The best solvent of silver is nitric acid, which even in a di- 
lute state and unaided by heat, converts it into nitrate. 
As silver is rather soft, it is usually employed in the arts 
alloyed with copper, whereby it acquires a greater hardness. 
Most silver coins consist of 90 per cent, silver and 10 percent, 
copper; the English shillings contain 92.5 per cent, silver. 
Oxygen forms three compounds with silver, but only the 
oxide affords corresponding salts. 
Silver Oxide—Ag20—is thrown out of silver nitrate solu- 
tion by sodium or potassium hydrate as a dark-brown amor- 
phous precipitate. It is somewhat soluble in water, and blues 
red litmus paper. In this, and in the neutral reaction of the 
nitrate, the strong, basic, alkaline nature of silver and its oxide 
exhibits itself; the soluble salts of nearly all of the other heavy 
metals show an acid reaction. When heated to 250°, the oxide 
decomposes into metal and oxygen ; at ioo° it js reduced by 
hydrogen. The hydrated oxide is not known ; the moist oxide 
reacts, however, very much like the hydroxides. 
On dissolving precipitated silver oxide in ammonium hydrate, 
black crystals (Ag20. 2NH3) separate when the solution evapo- 
rates and when dry these explode upon the slightest disturbance. 
(Fulminating silver). 340 
INORGANIC CHEMISTRY. 
Silver Suboxide—Ag40—corresponding to potassium suboxide, is 
produced by heating silver citrate in a current of hydrogen, and is a black, 
very unstable powder, which decomposes readily into silver oxide and 
silver. 
Silver Peroxide—AgO or Ag202—is formed by passing ozone over 
silver or its oxide, or by the decomposition of the nitrate by the electric 
current. It consists of black, shining octahedra, and at ioo° decomposes 
into Ag20 and oxygen. 
The salt-like compounds of silver correspond to the oxide Ag20, and 
are all constituted according to the form AgX, hence are termed argentic. 
They are analogous to the cuprous and mercurous derivatives, and show 
a great resemblance to the former in physical and chemical qualities. It 
would, therefore, be more correct to designate them argentowr. Com- 
pounds of the divalent form AgX2, are not known for silver. If, however, 
the mercurous and cuprous compounds are expressed by double formulas 
(p. 333): 
CuCl 
I 
CuCl 
Cu\ 
I o 
Cu/ 
HgCl 
I 
HgCl 
Hg\ 
I o 
Hg/ 
and 
which view is supported by their chemical deportment, those of silver 
might be represented by analogous formulas: 
AgCl 
I 
AgCl 
Ag\ 
I o 
Ag/ 
AgNOj 
I 
AgNOg. 
Then the silver atom would be divalent and a complete parallelism 
would be established with copper. The chemical formulas of solid bodies 
do not generally designate their true molecular values as in the case of 
gases, but only their simplest atomic composition. It is very probable that 
even the simplest chemical compounds, e.g., KC1 and AgCl, consist in their 
solid condition (as crystal molecules, p. 104) of complex molecules corres- 
ponding to the formulas (KCl)n (AgCl)m. An argument supporting 
this view is afforded by the existence of different modifications of chloride 
and bromide of silver ; these differ from each other in their external prop- 
erties, and in their different susceptibilities to light. The doubling of form- 
ulas, as shown above with Cu2Cl2, Hg2Cl2 etc., is mainly due to the ten- 
dency to deduce all the compounds of an element from a constant value, 
according to the doctrine of constant valence. This is, however, impos- 
sible (p. 172). According to present notions of valence, and as it is pre- 
sented in the periodic system, compounds (MeCl, MeCl2, MeCl3, etc.) are 
constituted according to definite forms or types that may materially determine 
their properties (p. 330). So far as the similarity of metallic compounds is 
concerned, it is of secondary importance whether the quantities correspond- 
ing to the simple formulas, in the solid or gaseous state, do unite to larger, 
complex molecules (compare HgCl, Cu2Cl2 — BC13, A1(CH3)3 and A12C16, 
GaCl3 and Ga2Cl6 — SnCl3 Sn2Cl4, PbCl2, etc.). In case of the sesqui- 
oxides M203 it is also immaterial whether they are derived from supposed 
trivalent elements (as A1203, Ga2Os, ln203), or from those that are 
tetravalent (as Fe203, Cr2Q3, Mn2Os). The same may be remarked of the 
metallic compounds Me304 = (MeOO)2.Me (see Spinels). 
The use of simple or of double formulas for the metallic compounds is 
therefore of no special importance. SILVER. 
341 
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 4510 to a 
yellow liquid, which solidifies to a horn-like mass. The chlo- 
ride is insoluble in dilute acids ; it dissolves somewhat in sodium 
chloride, readily in ammonium hydrate, potassium cyanide, 
and sodium hyposulphite. It crystallizes from ammoniacal so- 
lutions in large, regular octahedra. Dry silver chloride absorbs 
10 per cent, of ammonia gas, forming a white compound—• 
2AgCl3NH,—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 chlo- 
ride in ammonium hydrate; in other respects it is perfectly 
similar to the latter. Heated in chlorine gas it is converted 
into chloride. 
Silver Iodide—Agl—is distinguished from the chloride 
and bromide by its yellow color, its insolubility in ammonia, 
and its hexagonal crystals. It dissolves readily in hydriodic 
acid, to Agl. HI, which, upon evaporation of the solution, 
separates in shining scales. Heated in chlorine or bromine 
gas, it is converted into chloride or bromide; conversely, 
chloride and bromide of silver are converted into silver iodides 
by the action of hydriodic acid. 
These opposite reactions are explained by the principle of the greatest 
evolution of heat. Chlorine and bromine expel iodine from all iodides 
because the heat of formation of the latter is less than that of the bro- 
mides and chlorides (p. 266). Again, hydriodic acid (gaseous or in 
aqueous solution) converts silver chloride into the iodide according to the 
equation: 
AgCl + HI = Agl + HC1, 
because the heat modulus of the reaction is positive (for gaseous 
HI and HC1 -|- 12.5 C., for the solution -)- 10.6 C. See the Table at close 
of book.) 
Sunlight, and also other chemically active ray§ (magnesium 
light, phosphorus light) color silver chloride, bromide, and 
iodide, at first violet, then dark black, whereby they are prob- 
ably 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 photography: 
In photographic work a negative is first prepared. A glass plate is cov- 
ered with collodion (a solution of pyroxylin in an ethereal solution of alco- 
hol) holding in solution halogen salts of potassium or cadmium. After 
the evaporation of the ether the glass plate is covered by a dry collodion 
layer containing the haloid salts. The plate is now immersed in a solution 342 
INORGANIC CHEMISTRY. 
of silver nitrate, whereby haloid salts of silver are precipitated upon the 
surface. The plate thus prepared is exposed to light in the camera obscura, 
and, after the action, dipped into a solution 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. The plate is now introduced 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, by a dark layer of silver, while the 
places corresponding to shadows of the received image are transparent. 
The copying of the glass negative on paper is executed in a similar manner. 
Silver Cyanide—AgCN—is precipitated from silver solu- 
tions by potassium cyanide or aqueous hydrocyanic acid, as a 
white, curdy mass, not affected by light. It dissolves readily 
in ammonium hydrate and potassium cyanide, forming with 
the latter the crystalline compound AgCN,KCN. The solu- 
tion in potassium cyanide is employed in the electro silver-plat- 
ing of metals. 
Silver Nitrate—AgNOa, 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-half part water or in four parts alcohol; the solution having 
a neutral reaction. It fuses at 218°, and solidifies to a crystal- 
line mass. When perfectly pure it is not affected by light, but it 
usually turns black in sunlight with separation of metallic silver. 
Organic substances also reduce it to metal. Silver nitrate is 
employed in the cauterization 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 silver nitrate extracted 
with water. 
Silver Nitrite—AgN02—is precipitated from concentrated 
silver nitrate solutions by potassium nitrite. It crystallizes in 
needles, difficultly soluble in water, and decomposing above 
90°. 
Silver Sulphate—Ag2S04—is obtained by the solution of silver 
in hot sulphuric acid, and crystallizes in small rhombic prisms which are 
difficultly soluble in water. It is isomorphous with anhydrous sodium 
sulphate. 
Silver Sulphite, Ag2S03, is precipitated as a white, curdy mass, if sul- 
phurous acid be added to the solution of the nitrate. It blackens in the 
light and decomposes at ioo°. 
Silver Sulphide, Ag2S, occurs in regular octahedra, as 
argentite. Hydrogen sulphide precipitates it as a black amor- GOLD 
343 
phous sulphide from silver solutions. By careful ignition in 
the air it is oxidized to silver sulphate. It is insoluble in 
water and ammonium hydrate and dissolves with difficulty in 
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 to redness with access of air. 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 mer- 
cury 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 
cyanide 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 ammoniacal silver solution, 
with reducing organic substances like aldehyde, lactic, and tartaric acids. 
Under definite conditions, the reduced silver deposits upon the'glass as a 
regular metallic mirror. 
Recognition of Silver Compounds. — Hydrochloric 
acid throws down a white, curdy precipitate of silver chloride, 
which dissolves readily in ammonium hydrate. Zinc, iron, 
copper, and mercury throw out metallic silver from solutions 
of silver salts, and from insoluble compounds, like the chlo- 
ride. 
GOLD. 
Au = 196.2 
Gold (aururn) usually occurs in the native state, and is found 
disseminated in veins in some of the oldest rocks. Gold sands 
are formed by the breaking and disintegration of these. It is 
found, in slight quantity, in the sand of almost every river. 
Combined with tellurium it forms sylvanite, found in Tran- 
sylvania and California. It is present, 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 leaves the specifically heavier gold. 
The metals usually accompanying gold are silver and cop- 
per. To remove these, the gold is boiled with nitric or con- 344 
INORGANIC CHEMISTRY. 
centrated 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. There- 
fore, to separate pure gold from alloys poor in silver they must 
first be fused with about three-fourths their weight of the latter 
metal. Gold may be separated from copper and lead by cu- 
pellation (p. 338). 
Pure gold is rather 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 wire and beaten into thin 
leaves, which transmit green light. About 1035° it melts to a 
greenish liquid. It is not altered by oxygen, even upon ig- 
nition; acids do not attack it. It is only in a mixture of nitric 
and hydrochloric acids (aqua regia) which yields free chlo- 
rine, that it dissolves to gold chloride, AuCls. Free chlorine 
produces the same. Most metals, and many reducing agents 
(ferrous sulphate, oxalic acid) precipitate gold from its solu- 
tions as a dark-brown powder. 
As gold is very soft it wears away rapidly, and is, therefore, 
in its practical applications, usually alloyed with silver or cop- 
per, 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 per cent, gold and 10 per cent, copper. A 14-karat gold is 
generally employed for ornamental objects; this contains 
about 58.3 per cent, pure gold (24 karats representing 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. Its character is deter- 
mined to a high degree by these double relations (p. 329). Like 
the other elements of high atomic weight, mercury, thallium, 
lead, and bismuth, belonging to the same series of the periodic 
system, it varies considerably in character from its lower ana- 
logues. 
Gold, like silver and copper, yields compounds of the Form 
AuX—aurous, analogous to the cuprous and argentous. Be- 
sides, it has those of the form AuX3, auric derivatives, in which 
it is trivalent. These show the typical character of the tri- 
valent combination form, which expresses itself in the acidity 
of the hydroxides (p. 331) ; auric hydroxide, Au(OH)s, unites 
almost solely with bases. On the other hand, they show many 
similarities to the highest combination forms of the metals GOLD. 
345 
with high atomic weight: platinum (PtX4), mercury (HgX2), 
thallium (T1X3), and lead (PbX4) (p. 358). 
AUROUS COMPOUNDS. 
Aurous Chloride—AuCl—is produced by heating auric 
chloride, AuC13, to 180°, 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. _ . 
Aurous Iodide—Aul—separates as a yellow powder, if 
potassium iodide be added to a solution of auric chloride : 
When heated it breaks up into gold and iodine. 
, When auric oxide or sulphide is dissolved in potassium 
cyanide, large colorless prisms of the double cyanide, 
AuCN. KCN, crystallize out upon evaporation. 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- 343)- 
Aurous Oxide—Au20—is formed by the action of potas- 
sium hydrate upon aurous chloride. It is a dark violet powder 
which at 250° decomposes into gold and oxygen. It is changed 
to AuC13 and gold by the action of hydrochloric acid. 
Only a few double salts of the oxygen derivatives of mono- 
valent gold are known. 
AuC13 + 3KI = Aul + I2 + 3KC]. 
Auric Chloride—AuC13—results by the solution of gold 
in aqua regia, .and by the action of chlorine upon the metal. 
When the solution is evaporated the chloride is obtained as a 
reddish-brown, crystalline mass, which rapidly deliquesces in 
the air. It dissolves readily in alcohol and ether. 
Gold chloride forms beautifully crystallized double salts with 
many metallic chlorides, e.g., AuC13,KC1 + 2^H20 and 
AuC13,NH’C1 4- 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 
which decomposes, near 250°, into gold and oxygen. If the 
AURIC COMPOUNDS. 346 
INORGANIC CHEMISTRY. 
precipitate containing the magnesia be treated, not with con- 
centrated, but with dilute nitric acid, Auric Hydrate— 
Au(OH)3—remains as a yellowish-red powder. Both the oxide 
and hydroxide are insoluble in water and acids; they possess, 
however, acid properties, and dissolve in alkalies. Therefore 
the hydroxide is also called auric acid. Its salts, the aurates, 
are constituted according to the formula MeAu02, and are 
derived from the meta-acid, HAu02 = HO.AuO. 
. Potassium Aurate—KAu02 + 3H20—crystallizes in bright 
yellow needles, from a potassium hydrate solution of auric 
oxide. These are readily soluble in water ; the solution reacts 
alkaline. The corresponding aurates are precipitated from 
this solution by many metallic salts, e.g. : 
KAu02 + AgN03 = AgAu02 + KN03. 
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. 
Auric Sulphide—Au2S3—is precipitated as a blackish- 
brown compound, from gold solutions, by hydrogen sulphide. 
It dissolves in alkaline sulphides with formation of sulpho-salts. 
Stannous chloride (SnCl2) added to an auric chloride solu- 
tion produces, under certain conditions, a purple-brown pre- 
cipitate, purple of Cassius, which is employed in glass and 
porcelain painting. It probably consists of a mixture of 
aurous stagnate and stannous oxide. 
On pouring ammonium hydrate over auric oxide a brown 
compound is produced —fulminatinggold. When this is dried 
and heated or struck a blow, it explodes very violently. 
METALS OF GROUP III. 
The triatomic elements, affording derivatives mainly of the 
form MeX3, belong to group III of the periodic system 
(P- 243): 
Sc = 44 Y = 89 La = 138 Yb = 173 
B = 10.9 A1 = 27 o° 
Ga —69.8 In = 113.4 Tl= 203.6 
These bear the same relations to each other as do the ele- 
ments of group II ‘(p- 298). Boron has the lowest atomic 
weight, and the basic, metallic character in it is reduced very- 
much or does not appear at all. In its exclusively acidic METALS OF GROUP III. 
347 
hydroxide, B(OH)3, it approaches the metalloids, and is there- 
fore treated with them (p. 237). 
Aluminium is a perfect metal; its hydrate, Al(OH)3, exhibits 
a predominating basic character, and yields salts with acids. 
Its relations to boron are like those of silicon to carbon, or 
of magnesium to beryllium. The connection of aluminium 
and boron to the same group plainly shows itself in the entire 
character of the free elements, and in their compounds. Thus 
aluminium and boron are not dissolved by nitric acid, but by 
boiling alkalies: 
\\ + 3KOH = Al(OK), + 3H 
There is only a gradual difference between th'eir hydrates. 
Boron hydroxide, B(OH)3, not only acts as a feeble acid, but 
we also find that aluminium hydroxide manifests an acidic 
character, inasmuch as it is capable (p. 33 £) of forming metallic 
salts with strong bases (chiefly the alkalies); but owing to the 
higher atomic weight of aluminium the basic character exceeds 
the acidic. The similarity is also shown by the existence of 
perfectly analogous compounds; thus, e.g., the chlorides BC13 
and AICI3 can unite with PC15 and POCl3. 
lanthanum, and ytterbium attach them- 
selves to aluminium as the first sub-group. These constitute 
the third members of the great periods, and hence exhibit a 
pronounced basic character. As light metals, they are very 
similar to aluminium in their compounds, so that they all are 
embraced in one group, which (corresponding to the earthy 
nature of their oxides) is designated the Group of Earth 
Metals. Cerium and didymium bear a peculiar relation to 
lanthanum; their atomic weights are nearly alike and their 
properties very similar. Their apparently abnormal existence 
is explained by the fact that the 5th period (series 7 and 8), 
which is very incomplete, shows a somewhat varying function 
in its intermediate members (p. 245). The metals, erbium, 
terbium, thulium, and decipium, of recent discovery and but 
little characterized, may probably also be included in the 
same period. 
The second sub-group, -is more distinctly characterized and 
accurately investigated; it consists of the heavy metals, gallium, 
indium, and thallium. These belong to the right side of the 
great periods, possess, therefore, a less basic character, and 
bear the same relation to each other as Zn, Cd, and Hg. 
Aluminium was formerly classed together with chromium, iron, manga- 
nese, cobalt, and nickel, in one group, because they all afford sesquioxides, 348 
INORGANIC CHEMISTRY. 
Me203, whose salts are very much alike. Another fact which was thought 
to give weight to this classification was the existence of the similarly con- 
stituted alums: 
(S04)3A12S04K2 + 24H20 (S04)3Fe2S04K2 +. 24H20. 
In its entire behavior, aluminium is, however, very essentially distin- 
guished from the other metals here mentioned—by the acid nature of its 
hydrate, AhOH)s—and by its inability to form higher or lower combina- 
tion forms, while the others yield basic monoxides, Met), and acid-forming 
trioxides (Cr03, Fe03, Mn03l. Here, again, the similarity of the sesqui- 
oxide compounds, Me203, like those of the monoxide derivatives, is to 
be regarded as mainly influenced by the similarity of the combination 
forms (p. 330)., 
Potassium aluminium sulphate. 
Potassium iron alum. 
At present aluminium is assumed to be trivalentand this fact apparently 
contradicts the circumstance that not the simple formulas, A1C13, AlBr3, 
but the double ones, A12C16. Al2Br6, fall to its halogen derivatives (the result 
of vapor density determinations). On the other hand, however, the so-called 
metallo-organic compounds of aluminium exist, whose molecules are con- 
stituted according to the formulas, A1(CH3)3, A’l(C2H5)3; these undoubt- 
edly prove the trivalence of aluminium, because the compounds with hydro- 
carbon groups (like those with hydrogen) afford the surest guide for the 
deduction of the valence (p. 246). The existence of the molecules, A12Q16, 
Al2,Br6, etc., does not prove anything against its being a triad, but must be 
explained by a polymerization of the simple chemical molecules, A1C13, 
AlBr3. We find the same to be the case with arsenious oxide, As203, and 
antimony trioxide, Sb203, whose molecules in vapor form correspond with 
the doubled formulas, As4Oe (= As203,As,203) and Sb406; and with 
stannous chloride, whose molecule in vapor form at low temperature is 
Sn2Cl4, but higher it becomes SnCl2, or with gallium chloride that possesses 
the formulas, GaCl3 and Ga2Cl6 (p. 340). 
GROUP OF THE EARTH METALS. 
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 feld- 
spar, mica, and also in most crystalline rocks. It occurs, too, 
united with fluorine and sodium, as cryolite, in large deposits, 
in Iceland. 
Metallic aluminium is obtained by igniting the chloride, or 
better, the double chloride of sodium-and aluminium with me- 
tallic sodium: 
A1 = 27.0. 
AlCl3.NaCl + 3Na = A1 + 4NaCl. ALUMINIUM. 
349 
It is a silver white metal of strong lustre, is very ductile, and 
may be drawn out into fine wire and beaten into thin leaflets. 
Its specific gravity is 2.583; it belongs, consequently, to the 
light metals and possesses, therefore, all the properties opposed 
to those of the heavy metals (see p. 311). It fuses at a red 
heat but will not vaporize. It changes very little in the air at 
ordinary temperatures, and even when heated. If, however, 
thin leaves be heated in a stream of oxygen they will burn 
with a bright light. Nitric acid does not affect aluminium ; 
sulphuric acid only dissolves it on boiling, while it is readily 
soluble, even in the cold, in hydrochloric acid. It dissolves 
in potassium and sodium hydrate, with evolution of H, and 
forms aluminates: 
M + 3KOH = K3AIO3 + 3H 
Owing to its stability in ai.r 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, as watches, spoons, etc. 
Aluminium affords compounds of the form A1X3, or A1.2X6 (p. 
348) exclusively. Its salts, soluble in water, have an acid re- 
action, and a sweet, astringent taste. 
The heat of formation of some of the aluminium compounds equals : 
(A1„C16) =321.8. (A1„Br#) == 239.3. (A12,T6) = 140.6. 
(Al2,Cl6,Aq.)= 475.5. (Al2,Br6,Aq.) == 4099. (Al2,I6,Aq.) =318.6. 
(A12A,3H20) = 388.8. 
The heat evolved in the formation of a quantity of aluminium hydrate, 
corresponding to one atom of oxygen, is 129.6; since that of water is' far 
less (H2,0 = 69.0), it must be decomposed by aluminium, with liberation 
of hydrogen (p.274). If this does not transpire under ordinary conditions, 
the reason must be sought for in the insolubility of aluminium hydrate. 
Indeed the reaction occurs if aluminium chloride, or another salt, in which 
the aluminium oxide is soluble, be added to the water. Conversely, the 
high heat of formation of aluminium oxide explains why it is not reduced 
by carbon. 
Aluminium Chloride, A1C13, or A12C16, is produced by 
the action of chlorine upon heated aluminium : also by heating 
a mixture of aluminium oxide and carbon, in a current of 
chlorine : 
A1A + 3C + 6C1. = A12C1„ + 3CO. 350 
INORGANIC CHEMISTRY. 
Chlorine and carbon do not act separately upon the oxide ; 
by their mutual action, however, the reaction occurs in conse- 
quence of the affinity of carbon for oxygen, and of chlorine 
for aluminium. The oxides of boron and silicon show a simi- 
lar deportment. 
Aluminium chloride may be obtained in white, hexagonal 
leaflets by sublimation. It sublimes readily, but will only fuse 
when subjected to high pressure. Its vapor density is 133.2 
(H = i),from which we derive the molecular formula, A12C16 
= 266 (p. 348). The’specific heat of free aluminium indicates 
that its atomic weight is very nearly 27.0. 
Aluminium chloride absorbs moisture from the air, and deli- 
quesces. It crystallizes from concentrated hydrochloric acid 
solution, with 6 molecules of water. On evaporating the aque- 
ous solution, the chloride decomposes into aluminium oxide 
and hydrogen chloride: 
It forms double chlorides with many metallic chlorides, viz. : 
AlClg.NaCl, AICI3.KCI. The solutions of these may be evapo- 
rated to dryness without decomposition. It also unites with 
many halogen derivatives of the metalloids : 
A12C16 + 3H20 = A1203 + 6HC1. 
Aluminium Bromide—Al2Br6—is obtained like the chloride, and con- 
sists of shining leaflets which fuse at 90° and boil at 265-270°. Its vapor 
density is 267 4 (H = 1), corresponding to the formula Al2Br6. It behaves 
like the chloride. 
Aluminium Iodide—A12I6—is formed on heating aluminium filings 
with iodine. It is a. white, crystalline mass, fusing at 185°, and boiling 
about 400°. It is best prepared by covering sheet aluminium with carbon 
disulphide, and then adding the calculated amount of iodine gradually, 
letting the whole stand' for some time, and then distilling off the CS2. The 
reaction occurring between aluminium iodide and oxygen is interesting. 
If the vapor of the former be mixed with the latter, and then brought in 
contact with a flame, of if acted upon by an electric spark, a violent 
detonation will ensue; aluminium oxide and iodine result: 
AlClg.PClg, A1C13.P0C13, aici3.sci4. 
This deportment is due to the great difference in the heats of formation of 
the aluminium oxide (about 380 C.), and the iodide (140.6 C.). The chlo 
ride and bromide are similarly decomposed, but with less violence. 
Aluminium Fluoride—A1P’13 or A12F16—obtained by conducting hy- 
drogen fluoride over heated aluminium oxide or hydroxide, sublimes at a 
red heat in colorless rhombohedVa. It is insoluble in water, unaltered by 
acids, and is very stable. It yields insoluble double fluorides with'alkaline 
fluorides. The'compound—A1F1S. 3NaFl—occurs in Greenland, in large 
deposits, as Cryolite, and is employed m the soda manufacture (p. 291). 
A12I6 + 30 — Al20g -(- 61. ALUMINIUM. 
351 
Aluminium Oxide—A1203—is found crystallized in hex- 
agonal prisms in nature, as ruby, sapphire, and corundum, 
colored by other admixtures. Impure corundum, containing 
aluminium and iron oxides, is called emery, and serves for 
polishing glass. -The specific gravity of these minerals is 3.9 ; 
their hardness is only a little below that of the diamond. Arti- 
ficial aluminium oxide may be obtained by igniting the hydrate, 
and is a white amorphous powder, which fuses to a transparent 
glass in the oxy-hydrogen flame. A mixture of aluminium 
fluoride, and boron trioxide, heated to a white heat has the 
boron fluoride volatilized, and crystallized aluminium oxide 
remains: 
A12F16 -j- B203 — A1203 -T 2BF13. 
The crystallized or strongly ignited aluminium oxide is almost 
insoluble in acids;-to. decompose it, it is fused with caustic 
alkalies or with primary, potassium sulphate—HKSOr 
Aluminium Hydrates—The normal hydrate, Al(OH)3 or 
Al2(OH)6, occurs in nature as- hydrargillite. The hydrate, 
Al202(OH)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 an alkaline carbonate (in latter case 
carbon dioxide escapes, p. 353) to a soluble aluminium salt. 
Freshly precipitated, it dissolves in acids and in potassium and 
sodium hydrates. By long standing under water, or after dry- 
ing, it is, without any alteration in composition, difficultly 
soluble in acids. When carefully heated, the. normal hydrate 
first passes into AlO.OH. 
The freshly precipitated hydroxide dissolves readily in a solution of 
aluminium chloride or acetate. On dialyzing (p. 234) this solution the 
aluminium salt or crystalloid diffuses, and in the dialyzer remains the pure 
aqueous solution of the hydrate. This has a faint alkaline reaction and 
is coagulated by slight quantities of acid, alkalies, and many salts; the 
soluble hydrate passes into the insoluble gelatinous modification. 
Gelatinous aluminium hydroxide possesses the property of precipitating 
many dyestuffs from their solutions, forming colored insoluble compounds 
(lakes) with them. On this is based the application of aluminium hydrox- 
ide as a mordant in dyeing. The acetate is generally used for this purpose. 
Goods saturated with this salt are heated with steam, which causes the de- 
composition of the weak acetate; acetic acid escapes, while the separated 
aluminium hydroxide sets itself upon the fibre of the material. If the 
latter now be introduced into the solution of coloring matter the latter is 
fixed by the aluminium hydroxide 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 352 
INORGANIC CHEMISTRY. 
evaporating its solution in sodium or potassium hydrate,'or 
upon addition of alcohol, white amorphous compounds of 
KA102, NaA102 and (NaO)3Al are obtained. The potassium 
compound can be obtained in crystalline form. These deriva- 
tives, known as aluminates, are not very stable, and are even 
decomposed by carbon dioxide, with elimination of aluminium 
hydroxide: 
The aluminium hydrate obtained in this manner, in distinction 
from that precipitated from acid aluminium solutions by the 
alkalies, is not gelatinous, and is more difficultly soluble in 
acids, especially acetic. It comprises the ordinary alumina of 
commerce. 
On adding calcium chloride, strontium chloride, or barium 
chloride to the solution of potassium or sodium aluminate, 
white insoluble aluminates are precipitated : 
2A102Na -)- C02 = A12Os C03Na2. 
Similar aluminates frequently occur as crystallized minerals, in 
nature. Thus the spinels consist chiefly of magnesium Ru- 
minate, chrysoberyl is beryllium aluminate, 
AlO.OX-n ’ . . . . . A10.0\7 
AlO O// 6 ’ ga‘inite 1S zmc aluminate, Q^>Zn. 
Nearly all these minerals, commonly called spinels, crystal- 
lize in regular octahedrons, like the corresponding chromium 
compounds (see these) ; the exceptions are chrysoberyl, crys- 
tallizing in the rhombic system, and hausmannite, Mn304, in 
the' quadratic system. 
2A102Na + CaCl2 = (A102)2Ca + 2NaCl. 
Technically, alumina is obtained from cryolite, bauxite and other min- 
erals containing aluminium. The pulverized bauxite is heated with dry 
sodium carbonate in furnaces, and the resulting sodium aluminate extracted 
with water. From the clear solution carbon dioxide precipitates the hy- 
drate, while sodium carbonate remains dissolved, and is afterward recov- 
ered. The dried aluminium hydrate occurs as a white powder in trade. 
The gelatinous, readily soluble (colloidal) aluminium hydrate (see 
above) precipitated from acid-solutions by alkalies, has lately been pre- 
pared upon a large scale, according to the method of Lowig, by treating 
the sodium aluminate solution with milk of lime ; calcium aluminate precip- 
itates, while sodium hydrate remains in solution ; 
The calcium aluminate is dissolved in hydrochloric acid : 
2A102Na 4- Ca(OH)2 = (A102)2Ca + aNaOH. 
(A102)2Ca + 8HC1 = 2AICI3 + CaCl2 -f 4H20, 
and .to the solution now containing the alumina as chloride the correspond- ALUMINIUM. 
353 
ing amount of calcium aluminate added, and aluminium hydrate is precip- 
itated : 
According to this procedure, the sodium hydrate formed in the first re- 
action is obtained together with the alumina. On conducting carbon 
dioxide into a solution of alkaline, carbonates, and adding a solution of 
an alkaline aluminate at the same time, white aluminium-alkali carbonates 
are precipitated: 
2A1C13 + 3(A102)2Ca = 4A1203 + 3CaCl2. 
The caustic alkali that is formed in this way is converted again into bicar- 
bonate by carbon dioxide. In a dry state the precipitates are white, chalk- 
like masses which .at 90° contain five molecules of water: A1203,K20,2C02 
-(- 5H20. Their may be expressed by the formula: 
A1203,K20 + C03NaH = A1203,K20,2C02 + aNaOH. 
Al/ \C0. 
\Vk 
X0K 
Ala(OK)2(COs)2 
or 
They dissolve readily in dilute acids, even acetic, with evolution of carbon 
dioxide, and are suitable for the preparation of pure alumina mordants 
and antiseptic solutions (Lowig). 
The basic character of aluminium hydroxide exceeds the 
acid; but it is so feeble that it is not capable of forming salts 
with weak acids, as carbon dioxide, sulphurous acid, and 
hydrogen sulphide. When sodium carbonate is added to 
solutions of aluminium salts, aluminium hydroxide is precipi- 
tated, while carbon dioxide is set free: 
A12C16 + 3Na2C03 + 3H20 = Al2(OH)6 + 6NaCl + 3C02. 
The alkaline sulphide^behave similarly: 
A12C16 + 3(NH4)2S + 6H20 = Al2(OH)6 + 6NH4C1 + 3H2S. 
Aluminium Sulphate—A12(S04)3—crystallizes from aque- 
ous solution with 16 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 sul- 
phate is obtained by dissolving the hydroxide 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 
hydrate, insufficient for complete precipitation, is added to the 
sulphate, basic, sulphates separate out. Salts similar to the latter 
are also found in nature; thus,aluminite, used to prepare alum, 
has the composition : 
A1»\Sf}i + or (A1<*0)2S02 + 9H20. 354 
INORGANIC CHEMISTRY. 
Aluminium sulphate can combine with the alkaline sulphates 
and affords double salts, termed alums, e.g., potassium alum: 
(S04)3A12.S0;K2 -f- 24H20 or (S04)2A1K -f- I2H20. 
Their constitution is expressed by the following formula: 
O S-\ VI /S04K • 
n4c5>AL< + 24H20 
°*S// 2\SOjk 
or 
04S^m 
Qn/A1 
SO*\ + I2H„0. 
K 
In this compound the potassium may be replaced by sodium, 
ammonium, rubidium, caesium, and also by thallium. Iron, 
chromium, and manganese afford like derivatives: 
Fe2(S04)s.K2S04 + 24H20 M»2(S04)3. Na2S04 + 24H20. 
All these alums crystallize in regular octahedra or cubes, and 
can form ismorphous mixtures. 
Potassium iron alum. 
Sodium manganese alum. 
The most important of them is Potassium Aluminium Sul- 
phate, or ordinary alum, A1K(S04)2 + i2H20. It crystal- 
lizes from water in large, transparent octahedra, soluble in 8 
parts water of ordinary temperature, or in 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. When heated it melts in 
its water of crystallization, loses all the latter and becomes a 
white, voluminous mass—burnt alum. Upon adding a little 
sodium or potassium carbonate to a hot alum solution the 
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 quan- 
tities near Rome and in Hungary, has a similar composition 
(S04)2(A10)3K + 3h2o. 
Commercial alum is obtained according to various methods: I. From 
alunite, by heating and extracting with hot water. In this way alum dissolves 
while the 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 after moistening with water is ex- 
posed 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 treated directly with sulphuric acid, and to the solution of aluminium 
sulphate potassium or ammonium sulphate is added. 4. Bauxite and 
cryolite are admirable material for the preparation of alum. The work- ALUMINIUM. 
355 
ing of cryolite for alumina and soda is described on p. 291, and that of 
bauxite, p. 352. 
Ammonium Alum—(S04)2A1NH4 + i2H20—crystallizes 
like potassium alum, in large crystals, and at present, owing to 
its cheapness, is applied almost exclusively for technical pur- 
poses. Sodium alum is much more soluble, and crystallizes 
with difficulty. As the alum employed in dyeing must con- 
tain no iron, we understand why this salt is not applicable. 
At present the alum is being more and more supplanted by 
aluminium sulphate and sodium aluminate in all practical 
operations, because these chemicals can be procured perfectly 
free from iron. 
Aluminium Phosphate—A1P04 + 4H20—is thrown out 
of aluminium solutions by sodium phosphate, as a white gelati- 
nous precipitate; this is readily soluble in acids, acetic ex- 
cepted. 
Aluminium Silicates.—The most important of the alu- 
minium double silicates, so widely distributed in nature, are: 
lecuite, (SiOs)2AlK, albite or soda feldspar, Si308AlNa, ordi- 
nary feldspar—orthoclase—AlKSisOg—and the various micas, 
which, with quartz, compose granite. When these disintegrate 
under the influence of water and the carbon dioxideof the air, 
alkaline silicates are dissolved and .carried away by water, 
while the insoluble aluminium silicate, clay, remains. Perfectly 
pure clay is white, and is called kaolin, or porcelain clay; 
its composition mostly corresponds to the formula, Al2(Si03)3 
A1205H4, or Si209Al2H4. When clay is mixed with water a tough 
kneadable mass is obtained. By drying and burning, it be- 
comes compact and hard, and is the more fire-proof, the purer 
the clay. On this depends the use of clay for the manufacture 
of earthenware, from the red brick to porcelain. 
To produce porcelain a very fine mixture of kaolin, feldspar and quartz 
is employed. On strong ignition, the feldspar fuses, fills the pores of the 
clay and thus furnishes a fused transparent mass—porcelain. When it is 
not so strongly ignited, 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 ves- 
sels are constructed from impure clay, and they are usually glazed by 
throwing salt into the ovens at the time of burning. The hot steam de- 
composes the salt into hydrochloric acid and sodium hydroxide, which 
forms an easily fusible silicate on the surface of the clay. 
Ultramarine.—The rare mineral Lapis lazuli, which was 
formerly employed as a very valuable blue color under the 
name of Ultramarine, is a compound of aluminium sodium 356 
INORGANIC CHEMISTRY. 
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 
away from air. Green ultramarine is the product. 
This is then washed with water, dried, mixed with powdered 
sulphur and gently heated with air contact until the desired 
blue color has appeared—blue ultramarine. The 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 hy- 
drogen sulphide—this would point to the existence of a poly- 
sulphide. Violet and red ultramarines are prepared at present 
by conducting dry hydrogen chloride gas and air over com- 
mon ultramarine at 100-150°. 
RARE METALS. 
Li some very rare minerals, like cerite, gadolinite, euxenite and orthite, 
occurring principally in Sweden and Greenland, is found a series of metals 
which, in their entire deportment, closely resemble aluminium (p. 347). 
These are yttrium, cerium, lanthanum, didymium, and the more recent 
scandium, ytterbium,-erbium, terbium, thulium, and samarium (decipium). 
These generally form difficultly soluble oxalates, and are, therefore, pre- 
cipitated from solution by oxalic acid. They also afford difficultly soluble 
sulphates and double sulphates, of which 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 excel- 
lent means for their isolation and separation. 
Lanthanum, cerium and didymium have been most accurately investi- 
gated. Their atomic weights have deduced from the specific heats of the free 
metals, and are consequently approximately correct. That of yttrium (89) 
appears, from the isomorphism of its sulphate with that of didymium, to 
be positively established. Erbia, formerly regarded as an elementary sub- 
stance, consists, according to Nilson and Cleve, of the earths of six dif- 
ferent metals : scandium, ytterbium, thulium, erbium, terbium, and a metal 
designated X. Only the first two are as yet accurately characterized, and 
scandia is especially interesting, because the metal scandium contained in 
it', with the atomic weight 44, fills out a gap existing in the first large 
period. It coincides in all its properties with those deduced theoretically 
from the periodic system for the element ekaboron (compare Gallium). 
Scandium—Sc = 44—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 mag- 
nesia and oxide of beryllium). Its specific gravity equals 3.86; the spe- 
cific heat 0.1530. The hydrate, Sc(OH)3, is precipitated as a gelatinous 
mass front its salts by the alkalies, and is insoluble in an excess of the 
latter. The nitrate crystallizes in little prisms, and is decomposed with RARE METALS. 
357 
difficulty by heat. The potassium double sulphate ', Sc2(S04)3,3K2S04, is 
soluble in warm water, b.ut 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 a pure condition. Its chief source is gadolinite 
(upwards of 35 per cent.). Its potassium double sulphate is soluble in a 
potassium sulphate solution, and in this manner it can be readily separated 
from cerium, lanthanum, and didymium. Its nitrate is much more diffi- 
cult to decompose than those of scandium and ytterbium. The chloride, 
YClj -f- 7H2O, forms large prisms, and gives a spark spectrum. 
Lanthanum—La=i38.2—separated- from its chloride by electrolysis, 
resembles iron as regards color and lustre, oxidizes in the air, and burns 
in a flame with a bright light. Its specific gravity equals 6.16, the spe- 
cific heat 0.0448. The hydrate, La(OH)a, is precipitated as a gelatinous 
mass, and reacts alkaline. 
Cerium—Ce == 141-—occurs in cerite per cent.), and is also ob- 
tained by the electrolysis of the chloride. It is 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, 
Ce02. The first are colorless, while the latter are colored yellow or brown ; 
red ceric hydrate, Ce(OH)4, is precipitated from the first, on addition of 
hypochlorites. A little aqueous hydrofluoric acid will convert the ceric 
hydrate into cerium tetrafluofide. These compounds indicate that cerium 
is tetravalent and that it probably belongs to the fourth group of the pe- 
riodic system (p. 244). 
Didymium—Di = 142 or 145.2—in free condition resembles lan- 
thanum, but shows a somewhat yellowish.color ; it oxidizes in the air and 
burns in the flaifte with a brilliant light. Its specific gravity is 6.54, the 
specific heat 0.Q456. Didymium oxide, Di203, is a white powder. Its 
salts are rose-red, or violet in color. Besides these sesquioxide deriva- 
tives there is a didymium pentoxide, Di2Os -|- 3ll20, which in all proba- 
bility places didymium in Group V of the periodic system. 
The 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, to absorb definite light rays, yielding a very 
characteristic absorption spectrum. 
Samarium—Sm— 150 (Sm2Os)—occurs in orthite, thorite, and samar- 
skite. It is very similar to didymium and is characterized by an absorp- 
tion and spark spectrum. 
Ytterbium—Yb = 173.—Its oxide, Yb203, is obtained from the so- 
called erbium earth (from euxenite and gadolinite) by repeated partial 
heating of the mixed nitrates, whereby the scandium nitrate is the first to 
decompose. It is a white infusible powder, of specific gravity 9.17; its 
specific heat is 0.0646. The salts of ytterbium are colorless, and show no 
absorption spectrum. 
Erbium—Er= 166.—Erbium oxide, Er203, is red in color. Its salts, 
like those of didymium, yield an absorption spectrum. 
Terbium—Tr = 150.— Terbium oxide, Tr203, occurs in large amount 
in samarskite. It has an orange-yellow color, resembles the oxide of 
erbium, but does not show an absorption spectrum. 358 
INORGANIC CHEMISTRY. 
GALLIUM GROUP. 
The three heavy metals, gallium, indium, and thallium, bear 
the same relations to aluminium that we see exhibited by Cu, 
Ag and Au to sodium, Na and Zn, Cd and Hg to magnesium. 
Cu 
63.2 
Zn 
64.9 
Ga 
69.8 
As 
75-o 
Ag 
107.6 
Cd 
111.9 
In 
"3 4 
Sn 
117.0 
Sb 
119.0 
Au 
196 2 
Hg 
199.8 
T1 
203.6 
Pb 
206.0 
Bi 
207.0 
They constitute the corresponding members of the three 
great periods; and as second sub-group attach themselves to 
aluminium, while cerite metals form the first, more basic group 
(p. 347)- The entire character of the three elements under 
consideration is influenced by this position in the periodic 
system, because regular relations appear in all directions, as 
may be observed, for example, in the specific gravities, fusing 
points, and other physical properties in the free metals. 
Ga 
In 
T1 
Atomic weight 
69.8 
”34 
203.6 
Specific gravity 
5-9 
74 
11.8 
Fusing point 
3°° 
176° 
290° 
Belonging to group III of the periodic system, Ga, In, and 
T1 yield compounds of the trivalent form, and these are anal- 
ogous to those of aluminium in many respects. 
Thallium, like other elements with high atomic weight (Au, 
Hg, Pb), exhibits great variations from the group properties 
(p. 322). It yields, for example, not only derivatives of the 
form T1X3, but also those of T1X. If we include thallium 
as a member of the last great period (Pt, Au, Hg, Tl, Pb, Bi) 
we will discover that, as in case of the other metals of this 
series, a remarkable regularity underlies all its forms of com- 
binations—the highest as well as the lowest. 
PtCl2 AuCl HgCl T1C1 PbCl2 BiCl3 
PtCl4 AuC13 HgCl2 TIClj PbCl4 BiX5. 
1. GALLIUM. 
Ga = 69.8. 
Gallium was discovered in zinc blende from Pierrefitte, in 1875, by Lecoq 
de Boisbaudran, by means of the spectroscope. As early as the year 1870, 
Mendelejeff, taking the table of the periodic system devised by him as INDIUM. 
359 
basis, predicted the existence of a metal (standing between aluminium 
and indium, with an atomic weight of nearly 69), which he named Eka- 
aluminium. Its properties were necessarily deduced from its position in 
the periodic system. All the properties of gallium known at that time 
agreed with those of eka-aluminium, and it seemed very probable that this 
element, which had been theoretically established, was in reality gallium. 
This is now confirmed by the fact that the atomic weight, determined by 
experiment, agrees with that deduced theoretically. 
As yet gallium has only been found in very small quantity, and is but 
imperfectly investigated. It is characterized by a spectrum consisting of 
two violet lines. Separated by electrolysis from ammoniacal solution of 
its sulphate, it is a white, hard metal, of specific gravity 5.9, with a fusing 
point 30°. It is only superficially oxidized in the air, not altered by water, 
and is not volatile up to a red heat. Like aluminium, it is scarcely at- 
tacked by nitric acid, but dissolves readily in hydrochloric acid. 
Gallium Oxide—Ga203—is obtained by igniting the nitrate. It is a 
white mass which sublimes when heated in a current of hydrogen. The 
hydroxide—Ga(OH)3—is thrown out of solutions of its salts by the 
alkalies as a white flocculent precipitate, readily soluble-in an excess of the 
precipitant, but rather difficultly soluble in ammonium hydrate. 
Gallium Chloride—GaCl3—is produced on heating gallium in a cur- 
rent of chlorine ga$; it forms CQlorless crystals that fuse at 750, sublime 
about 6o° and boil at 215-220°. Its vapor density at 440°-corresponds to 
the formula GaClg, at 270° very closely to Ga2Cl6. The chloride fumes 
in the air, like aluminium chloride, deliquesces and decomposes in the 
evaporation of its aqueous solution. 
Gallium Nitrate—Ga (NOs)3, and Gallium Sulphate, Ga2(S04)3— 
are crystalline and very deliquescent. 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. 
2. INDIUM. 
Owing to its resemblance to zinc, indium was regarded as a" divalent 
metal, and its compounds composed according to the formula, In X2; this 
fixed its atomic weight at 75.6. The specific heat, however, made the 
atomic weight one and a half times as large (p. 257). Hence it is trivalent 
and its derivatives are constituted according to the form, InX3. It belongs 
to the group of aluminium, and, in its derivatives, manifests some similar- 
ity to this metal. 
It was discovered, in 1863, by Reich and Richter, by the aid of spec- 
trum analysis. Its spectrum is characterized by a very bright indigo blue 
line, hence its name. It only occurs in very minute quantities in some 
zinc blendes from Freiberg and the Hartz. 
It is a silver-white, soft and tenacious metal, of specific gravity, 7.42. 
It fuses at 176° and distils at a white heat. At ordinary temperatures it 
is not altered in the air; heated, it burns with a blue flame to indium 
oxide. It is difficultly soluble in hydrochloric and sulphuric acids, but 
dissolves readily in nitric.acid. 
In = 113.4. 360 
INORGANIC CHEMISTRY. 
Indium Chloride—InCl3—results from the action of chlorine on 
metallic indium, or upon an ignited mixture of indium oxide, and carbon. 
It sublimes in white, shining leaflets, which deliquesce in the air. Its 
vapor density corresponds to the formula InCl3. It does, not decompose 
when its aqueous solution is evaporated. • 
Indium Oxide—ln203—is a yellow powder resulting from the ignition 
of the hydroxide. 
Indium Hydroxide—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(N03l3—crystallizes with three molecules of wa- 
ter, in white deliquescent needles. 
Indium Sulphate— In2(S04)3—remains on evaporating a solution of 
indium in sulphuric acid as a gelatinous mass, with three molecules of 
water. It forms an alum with ammonium sulphate. 
Indium Sulphide—In2S3—is precipitated by hydrogen sulphide as a 
yellow-colored compound from indium solutions. 
3. THALLIUM. 
T1 = 203 6. 
Thallium is rather widely distributed in nature, but in very 
small quantity. The very rare mineral crookesite contains 17 
per cent, of the metal, together with copper, selenium and 
silver. It is often found with potassium in sylvite and carnal- 
lite, in mineral springs, and in some varieties of pyrite and 
zincblendes. When These pyrites are roasted for the produc- 
tion of sulphuric acid, according to the chamber process, the 
thallium deposits as soot in the chimney and in the chamber 
sludge, and was discovered in the latter, in 1863, almost si- 
multaneously, by Crookes and Lamy, by means of the spec- 
troscope. 
To get the thallium, the chimney-dust is boiled with water 
or sulphuric acid, and thallous chloride precipitated from the 
solution by hydrochloric acid. The chloride is then converted 
into sulphate, and the metal separated from the latter by means 
of zinc or the electric current. Thallium is a white metal, as 
soft as sodium, and has the specific gravity 11.8. It fuses at 
290°, and distils at a white heat. It oxidizes very rapidly in 
moist air. It does not decompose water at ordinary tempera- 
tures. It is, therefore, best preserved under water in a closed 
vessel. By air access it gradually dissolves in the water, form- 
ing thallium hydroxide 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 
green. Thallium dissolves readily in sulphuric and nitric acids, THALLIUM. 
361 
but is only slightly attacked by hydrochloric acid, owing to the 
insolubility of thallous chloride. 
Thallium forms two series of compounds: thallous— 
1 m 
T1X and thallic—T1X3. The first are very similar to the 
compounds of the alkalies (and also those of silver). The 
solubility of the hydroxide and carbonate in water shows this; 
their solutions have an alkaline reaction. Again, many thal- 
lous salts are isomorphous with those of potassium, and afford 
similar double salts. In the insolubility of its sulphur and 
halogen compounds, monovalent thallium approaches silver 
and lead. 
In its compounds of the form T1X3 thallium is trivalent, 
like aluminium, but otherwise shows scarcely any similarity 
to-the latter. 
The heat of formation of some of the thallous compounds is: 
(T1,0) = 42.2 (T1,C1) = 48.5 (Tl,Br) = 41.2 (T1,I) = 30 
' (T1,0,H) = 56-9 (.T12,S,04) = 210.9 (T1,N,03) = 58 1. 
The heat of solution of all these compounds is negative. From the 
numbers cited above we can understand the deportment of thallium toward 
water and the acids. 
The heat of formation of the ic compounds in aqueous solution equals 
(Tl,Cl3Aq.) = 89.0 (Tl.Br3.Aq.) = 56.1 (Tl.I3.Aq.) = 10.5 
(T12,03)3H20) = 86.9. 
THALLOUS COMPOUNDS. 
, Thallous Oxide—T120—is formed by the oxidation of 
thallium in the air, or by heating the hydroxide to xoo°. It 
is a black powder which dissolves in water with formation of 
the hydroxide. 
Thallous Hydroxide—Tl(OH)—may be prepared by 
decomposing 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, 
affording strong alkaline solutions. 
Thallous Chloride—T1C1—is thrown down from solu- 
tions of thallous salts by hydrochloric acid as a white, curdy 
precipitate, which is very difficultly soluble in water. It sep- 
arates in small crystals from the hot solution. It fuses at 4270, 
and boils about 715°. Like potassium chloride, it affords an 
insoluble salt with platinic chloride—PtCl4,2TlCl. Thallous 
bromide forms a white, and thallous iodide a yellow precipi- 
tate. 362 
INORGANIC CHEMISTRY. 
Thallous Sulphate—T12S04—crystallizes in rhombic 
prisms, isomorphous "with potassium sulphate. It dissolves in 
20 parts of water at ordinary temperatures. It affords double 
salts with the sulphates of the metals of the magnesium group, 
of ferrous oxide, of cupric oxide, etc. (p. 314), e.g., 
MgS04Tl2S04 + 6H20; these are perfectly similar and anal- 
ogous to the corresponding double salts of potassium and 
ammonium. It affords thallium alum with the sulphates of 
the sesquioxides of the iron group, e.g., A1T1(S04)2 + i2H20 ; 
these are similar to potassium alum—A1K(S04)2 -j- i2H20. 
Thallous Carbonate—T12C03—is obtained from the 
oxide by the absorption of C02; it crystallizes in needles, which 
dissolve at ordinary temperatures in 20 parts of water. The 
solution has an'alkaline reaction. 
Thallous Sulphide—T12S—is precipitated from thallous 
salts by hydrogen sulphide as a black compound, insoluble in 
water. 
THALLIC COMPOUNDS. 
Thallic Chloride—T1C13—is produced by the action of 
chlorine upon T1 or T1C1 in water, and is very soluble in water. 
It decomposes at ioo° into T1C1 and Cl2. The alkalies precipi- 
tate from its solutions thallic hydrate, TIO.OH, a brown 
powder, which, at ioo°, passes into thallic oxide, ti2o3. 
Further heating decomposes the latter into thallous oxide and 
oxygen. 
The oxide and hydroxide are soluble in hydrochloric, nitric, 
and sulphuric acids, forming ThNOj^.Tl^SOJj/TlClg. 
On conducting chlorine through a solution of thallic hy- 
droxide in potassium hydrate, it 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. 
Tin and lead, with silicon and carbon, constitute a group, 
in which the transition from metalloidal to metallic character 
finds full expression ; it accords with the increasing atomic 
weights: C = 11.9, Si= 28, — = 73, Sn = 117.5, Pb= 206.4. 
The differences between these elements are quite gradual; 
TIN AND LEAD. 363 
TIN AND LEAD. 
the noticeably greater difference between metalloidal silicon 
and metallic tin is explained by the fact that the member cor- 
responding to arsenic of the nitrogen group, is not known (p. 
162). Its atomic weight would be near 73. 
Tin, like carbon and silicon, is tetravalent, and yields per- 
fectly analogous compounds, e.g , SnCl4,Sn02. The hydrates 
Sn(OH)4and SnO(OH)2, correspond to tin dioxide; only the 
second forms salts (p. 225). Stannic acid is, however, a weak 
acid ; its alkali salts react alkaline, and are not very stable. 
It also has a basic character, hence affords salts with acids. 
Therefore, stannic acid anhydride is called stannic oxide, and 
stannic acid, stannic hydroxide. The metallic salts of stannic 
acid (like Sn03Na2) are called stannates; those with acids 
are stannic salts. 
Tin also affords compounds of the form SnX2 (correspond- 
ing to CO), in which it figures as a dyad. The monoxide, 
called stannous oxide, possesses a decided basic character, and 
only yields salts with acids—stannous or stanno-salts. 
A further advance in metallo-.basic character is exhibited 
by lead, which, like tin, forms two series of compounds of the 
forms PbX4 and PbX2. While with tin the compounds SnX4 
are more stable than those of the form SnX2, the derivatives 
of lead are almost t xclusively of the form PbX2; in these the 
metal acts as a dyad. 
The tetravalence of lead is limited almost entirely to its 
metallo-organic compounds (as Pb(CH3)4,Pb(C2H5)4—p. 246), 
and lead dioxide Pb02. The latter does not form correspond- 
ing salts with the acids, but parts with one atom of oxygen with 
formation of salts of the monoxide, PbO. When heated with 
hydrochloric acid, it liberates chlorine; hence behaves like 
the peroxides, and is commonly called lead peroxide (p. 263). 
A distinct graduation shows itself in the series, Si02,Sn02, 
Pb02. The stability and acidity grow successively less, 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 anhydride of a plumbic acid, H2Pb03. 
That tin and lead belong to one group follows also from the heat of 
formation of their compounds: 
(Sn,Cl2) = 80.8 (Sn,0) = 68.0 (Sn,Cl4) = 127.0 (Sn,02) = 136.0 
(Pb,Cl2) = 82.7 (Pb,Oj = 50.3 (Pb,Cl4) = (Pb,02) = 74.0 
These numbers explain why, for example, tin is thrown out of solutions 
of stannic chloride (as well as from other salts) by lead, and conversely 
why tin precipitates lead, from alkaline solutions of lead oxide (p. 262). 364 
INORGANIC CHEMISTRY. 
Sn =: 117-5- 
I. TIN. 
Tin occurs in nature principally as dioxide (Cassiterite— 
tin stone) in England (Cornwall), Saxony, and India. To 
prepare the metal the oxide is roasted, lixiviated, and heated 
in a furnace with charcoal: 
Sn02 + 2C = Sn -|- 2CO. 
Thus obtained, it usually contains iron, arsenic, and other 
metals; to purify it the metal is fused at a low temperature, 
when the pure tin flows away, leaving the other metals. The 
tin obtained in the Indian isles (Malacca) is almost chemically 
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 ; 
and when a rod of it is bent it emits a peculiar sound (tin cry), 
due to the friction of the crystals. Upon etching a smooth 
surface of tin with hydrochloric acid, its crystalline structure 
is recognized by the appearance of remarkable striations. At low 
temperatures perfectly pure compact tin passes gradually into 
an aggregate of small quadratic crystals. The metal is tolerably 
soft, and very ductile, and may be rolled out into thin leaves. 
It becomes brittle at 200°, and may then be powdered. It fuses 
at 228°, and distils at a white heat (about 1700°); it burns with 
an intense white light when heated in the air, and forms tin 
dioxide. It does not oxidize in the air at ordinary temperatures, 
and withstands the action of many bodies, hence is employed 
in tinning copper and iron vessels for household use. 
The most interesting of the tin alloys, besides bronze and 
soft solder, is britannia metal. It contains 9 parts tin and 1 
part antimony, and frequently, also 2-3 per cent, zinc and 1 
per cent, copper. 
Tin dissolves in hot hydrochloric acid, to stannous chloride, 
with evolution of hydrogen gas :' 
Concentrated sulphuric acid, when heated, dissolves tin, 
with formation'of stannous sulphate. 
Somewhat dilute nitric acid oxidizes it to metastannic acid ; 
while anhydrous nitric acid, HN03, does not change it. It dis- 
solves when boiled with potassium or sodium hydrates, forming 
stannates: 
Sn -f 2HC1 = SnCl2 + 2H. 
Sn -f 2KOH +. H20 Sn03K2 + 2H2. TIN. 
365 
There are two series of tin compounds : the stannous, and 
stannic or compounds of stannic acid (p. 363). 
STANNOUS COMPOUNDS. 
Tin Dichloride—Stannous chloride, SnCl2—results when 
tin dissolves in concentrated hydrochloric acid. When its 
solution is evaporated it crystallizes with two molecules of 
water (SnCl, + 2H20) which it loses at ioo° C. It is used in 
dyeing, as a mordant, under the name of Tin Salt. The an- 
hydrous chloride, obtained by heating the metal in dry hydro- 
chloric acid gas, fuses at 250° and distils without decomposition 
about 620°. Its vapor density at 600-700° agrees with the 
formula Sn2Cl4; at 900° with SnCl2. 
Stannous chloride dissolves readily in water. Its'solution is 
strongly reducing, and absorbs oxygen from air with the sepa- 
ration of basic stannous chloride : 
3SnCl2 + O + Ha0 =± aSn/gj + SnCl4. 
In the presence of hydrochloric acid, only stannic chloride 
is produced. Stannous chloride precipitates mercurous chlo- 
ride and metallic mercury from solutions-of mercuric chloride 
(p. 326). It unites with chlorine to form stannic chloride, 
and with many chlorides to yield double salts, e.g. : 
Tin Monoxide—SnO, or Stannous oxide—is obtained by 
heatihg its hydrate, Sn02H2, in an atmosphere of carbon di- 
oxide ; it is a blackish-brown powder, which burns when 
heated in the air, and becomes stannic oxide. Sodium car- 
bonate added to a solution of stannolis chloride precipitates 
white 
Stannous Hydroxide—stanno-hydrate—Sn(OH)2: 
SnCl2.2KCI and SnCl2.2NH4Cl. 
SnCl2 + C03Na2 + H20 = Sn(OH)2 +' 2NaCl -f CC2. 
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; but, on boiling the 
solution, the hydrate decomposes into potassium stannate, 
K2SnOs, which remains dissolved, and metallic tin. 
The hydroxide affords salts by its solution in acids. Stannous 
chloride—SnCl2—or stanno-sulphate—SnS04—is formed when 
tin is warmed with concentrated hydrochloric or sulphuric 366 
INORGANIC CHEMISTRY. 
acid. The sulphate separates in small, granular crystals, when 
its solution is evaporated. 
Tin Monosulphide—Stannous sulphide—SnS—is pre- 
cipitated from stannous solutions by hvdrogen sulphide', as a 
dark-brown amorphous precipitate. Obtained by fusing tin 
and sulphur together, it is a lead-gray crystalline mass. It 
dissolves in concentrated hydrochloric acid, with liberation of 
H2S, and forms stannous chloride. It is insoluble in alkaline 
sulphides, but, if sulphur be and the solution bailed, it 
will dissolve as a sulpho-stannate (p. 367). 
SnS + S -f K2S = KjSnSs. 
STANNIC COMPOUNDS. 
Tin Tetrachloride—Stannic chloride — SnCl4—is pro- 
duced by the action of chlorine upon heated tin or stannous 
chloride—SnCl2. It is a colorless liquid (Spiritus fumans 
Libavii), fuming strongly in nooist air, of specific gravity 2.27, 
and boiling at 1140 ; its vapor density equals 129 (H = 1), 
corresponding to the molecular formula, SnCl4 = 258.9. It 
attracts moisture from the air and is converted into a crystal- 
line mass (Butter of Tin), SnCl4 -+- 3H2Q, readily soluble in 
water. Boiling decomposes the solution inJ:o metastannic acid 
(H2Sn03) and hydrochloric acid : 
SnCl* + 3H20 = H2SnOs + 4HCI. 
Stannic chloride possesses a salt-like nature, and.combines with 
metallic chlorides to the so-called double salts, e.g., SnCl4. 
2KCI and SnCl4.2NH4Cl; the latter compound is known as 
pink salt in calico printing. It also yields crystalline double 
salts with chlorides of the metalloids, e.g.,. SnCl4. PC15 and 
SnCl4.2SC14. 
Tin Bromide—SnBr4—forms a white, crystalline mass that melts at 
30° and boils at 200°. 
Tin Iodide—Snl4—consists of orange-red octahedra, fusing at 146° 
and boiling at 2950. It may be obtained by heating tin with iodine to 50°. 
Tin Fluoride—SnFl4—is only known in combination with metallic 
fluorides (e.g., K2SnFl6), which are very similar to and generally isomor- 
phous 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 it then forms a white, amorphous powder. 
It may be obtained' crystallized, by conducting vapors of the LEAD. 
367 
tetrachloride and water through a tube heated to redness. The 
dioxide is infusible, and not soluble in acids or alkalies. When 
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. If 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 ex- 
plained. The insoluble modification is commonly called 
metastannic acid. It is also obtained as a white powder by 
digesting .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 water decomposes them. The metal salts of stannic 
acid are more stable. The most important of these is sodium 
stannate—Na2Sn03 + 3H20—which is employed in calico 
printing, under the name of preparing salts. It is produced 
upon a large scale by fusing tin stone with sodium hydrate. 
On evaporating the solution, it crystallizes in large, transpar- 
ent, hexagonal crystals, containing three molecules of water. 
Tin Disulphide—Stannic sulphide—SnS2—is precipitated 
as an amorphous, yellow powder by H2S from stannic solu- 
tions. If a mixture of tin-filings, sulphur, and ammonium 
chloride be heated it is obtained in form of a brilliant crystal- 
line mass, consisting of gold-yellow scales. It is then called 
mosaic gold, and is applied in bronzing. Concentrated hydro- 
chloric acid dissolves the precipitated disulphide, forming 
stannic chloride ; nitric acid converts it into metastannic acid. 
The sulphides and hydrosulphides of the alkalies dissolve in 
tin disulphide forming sulphostannates (see p. 221). Sodium 
sulphostannate, SnS3Na2 + 2H20, crystallizes in colorless oc- 
tahedra. Acids decompose the sulphostannates with the sepa- 
ration of tin disulphide. 
Pb = 206.4. 
2. LEAD. 
Lead (.Plumbum) is found in nature principally as Galenite 
—PbS. The other more widely distributed lead ores are 368 
INORGANIC CHEMISTRY. 
Cerussite—PbCOs—Crocoisite (PbCr04) and Wulfenite (Mo04 
Pb). Galenite is the chief source of lead ; 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 pxide and sul- 
phate : 
and 
PbS + 30 = PbO + S02 
PbS + 04 = PbS04. 
Upon ignition, these two substances react with the lead sul- 
phide according to the follQwing equations: 
and 
2PbO + -PbS = 3Pb + S02 
S04Pb -f PbS =5= 2Pb + 2S02. 
Metallic lead has a bluish-white color, is very soft, and tol- 
erably ductile. A freshly cut surface has a bright lustre, but 
on exposure to air becomes dull- by oxidation. Its specific 
gravity is 11.37. It fuses at 3250, and distils at a white heat 
(about 1700°). It burns to lead oxide when heated in the 
air. 
In contact with air and water lead oxidizes to lead hydrox- 
idej Pb(OH)2, which is somewhat soluble in water. If, how- 
ever, the water contain carbonic acid and mineral salts—even 
in slight quantity, as in natural waters—no lead goes into so- 
lution, but it is covered with an insoluble layer of lead carbonate 
and sulphate. When much carbon dioxide is present the car- 
bonate is somewhat soluble in the water. This behavior is very 
important for practical purposes,, as lead pipes are frequently 
employed in conducting water. 
Sulphuric .and hydrochloric acid have little effect on the 
metal, owing to the insolubility of its sulphate and chloride; 
yet, if the lead be in the form of powder both acids will dis- 
solve it. It forms lead nitrate with dilute nitric acid. Zinc, tin, 
and iron precipitate it, as metal, from its solution; a strip of zinc 
immersed in a. dilute solution of lead acetate is covered with 
an arborescent mass, consisting of shining crystalline leaflets 
(lead tree). 
Alloys.—An alloy of equal parts lead and tin fuses at 1860, 
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 usual lead compounds are constituted according to the 
type PbX2, and are called plumbic (p.363). They show a LEAD. 
369 
slight similarity to the stannous derivatives. Many of the 
lead salts are isomorphous with those of barium ; the sulphates 
of both metals are insoluble. 
The heat of formation of some of the lead compounds equals: 
(Pb,Cl2) = 82.7 (Pb,Br2) = 64.4 (Pb,I2) — 39-6 (Pb,0) = 50.3. 
(Pb,S) = 20.4 (Pb,S,04) = 216.2 (Pb,N2,Og\ = 105.5 
If we include the heat of solution with the above numbers, they will af- 
ford us an explanation for the deportment of lead toward acids, as well 
as for the various transpositions of its compounds. 
Lead Oxide—PbO—is produced when lead is heated in 
air. After fusion it solidifies to a reddish-yellow mass of 
rhombic scales (litharge). When lead is carefully roasted, or 
the hydrate or nitrate ignited, we obtain a yellow amorphous 
powder called massicot. Lead oxide has strong basic prop- 
erties; it absorbs carbon dioxide from the air, and imparts an 
alkaline reaction to water as it dissolves as hydrate. Like 
other strong bases it saponifies fats. It dissolves in hot po- 
tassium hydrate, and on cooling crystallizes from solution in 
rhombic prisms. 
Lead Hydroxide—Pb(OH)2.—Alkalies throw it out of 
lead solutions as a white, voluminous precipitate. 
It imparts an alkaline reaction to water, as it is somewhat 
soluble, and absorbs carbon dioxide with formation of lead 
carbonate. When heated to 130° it decomposes into lead 
oxide and water. 
If lead or the amorphous oxide be heated to 300-400°, 
for some time, in the air, it will absorb oxygen and be con- 
verted into a bright red powder, called red lead, or minium. 
Its composition corresponds to the formula, Pb304; it is con- 
sidered a compound of PbO with lead peroxide (Pb304 = 
2PbO + Pb02). When minium is treated with somewhat 
dilute nitric acid, lead nitrate passes into solution, while a dark- 
brown amorphous powder—lead peroxide, Pb02—remains. 
This oxide is more conveniently obtained by adding bleach- 
ing lime (and milk of lime) to a concentrated lead chloride 
solution (in calcium chloride): 
2PbCl2 + Ca (0C1)2 + 2H20 = 2Pb02 + CaCl2 + 4HCI. 
Or chlorine is conducted into a mixture of lead chloride (2 
mols.) and lime hydrate (3 mols.) with water. 370 
INORGANIC CHEMISTRY. 
Lead peroxide dissolves in cold hydrochloric acid to a 
reddish-yellow liquid containing lead tetrachloride, PbCl4. 
The latter also results from the action of chlorine gas on a 
mixture of lead chloride with hydrochloric acid. The tetra- 
chloride cannot be obtained free, because its solution readily 
decomposes into PbCl2 and chlorine. 
Oxygen is disengaged when sulphuric acid acts upon it, and 
lead sulphate (PbS04) formed. When dry sulphur dioxide is 
conducted over it, glowing sets in and lead sulphate results: 
Pb02 + SO, = PbS04. 
When ignited Pb02 breaks up into PbO and oxygen. 
As previously mentioned fp. 363), 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—separate 
out; these are perfectly analogous to potassium stannate—K2SnOs-(- 
3H20. An alkaline lead oxide solution added to a solution of potassium 
plumbate produces a yellow precipitate (Pb304-|-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(OH)4, and silicic, SqOH )4, acids : 
Another oxide—Pb203—which is precipitated as a reddish-yellow 
powder on the addition of sodium hypochlorite to an alkaline lead solu- 
tion, is very probably the lead salt of metaplumbic acid : Pb203 = PbPb03. 
Nitric acid decomposes it into lead nitrate and peroxide. It dissolves in 
hydrochloric acid without liberation of chlorine ; this gas escapes, however 
when the solution is heated. 
Pb304 = Pb2PbO+. 
Lead Chloride—PbCl2—separates as a white precipitate, 
when hydrochloric acid is added to the solution of a lead salt. 
It is almost insoluble in cold water; from hot water, of which 
it requires 30 parts for solution, it crystallizes in white, shining 
needles. It melts about 500° and solidifies to a horn-like mass. 
It is volatile at a white heat; its vapor density corresponds to 
the formula PbCl2. 
Lead Iodide—Pbl2—is thrown down as a yellow precipi- 
tate from lead solutions by potassium iodide; it crystallizes 
from a hot solution in shining, yellow leaflets, melting at 383°. 
Lead Nitrate—Pb(N03)2, obtained by the 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. 
When boiled with lead oxide and water, it is converted into 
the basic nitrate, which separates in white needles. LEAD. 
371 
Lead Sulphate—PbS04, occurs in nature as Anglesite, in 
rhombic crystals, isomorphous with barium sulphate. It is 
precipitated from lead solutions as a white crystalline mass by 
sulphuric acid. It dissolves with difficulty in water, more 
readily in concentrated sulphuric acid. When ignited with 
carbon, it is decomposed according to the following equa- 
tion : 
Lead Carbonate—PbC03, occurs as Cerussite in nature. It 
is precipitated by ammonium caibonate from lead nitrate so- 
lutions. Potassium and sodium carbonates precipitate basic 
carbonates, the composition of which varies with the temper- 
ature and concentration of the solution. A similar basic salt, 
whose composition agrees best with the formula : 
PbS04 + 2C = PbS + 2C02. 
co/?bOH 
2PbC03,Pb(0H)2 = 3>Pb 
CO / 
UUs\PbOH, 
is prepared on a large scale by the action of carbon dioxide 
upon lead acetate. It bears the name of white lead. 
White lead was formerly manufactured by what is known as 
the Dutch process. Rolled lead sheets were moistened in 
earthenware pots, with acetic acid, and then covered with 
manure and permitted to stand undisturbed for some time. 
In this way, the action of the acetic acid and air upon the 
lead produced a basic acetate, which the C02, evolved from 
the decaying manure, converted into basic lead carbonate. 
At present it is prepared by dissolving litharge in acetic acid, 
and transferring the resulting basic acetate into a carbonate 
by conducting carbon dioxide into it. 
White lead is employed for the manufacture of white oil 
colors. As it is poisonous, and blackened by the hydrogen 
sulphide of the air (formation of lead sulphide), it is being 
replaced more and more by zinc white and permanent white 
(BaSOj. 
Lead Sulphide—PbS—occurs crystallized in metallic, 
shining cubes and octahedra. Hydrogen sulphide precipi- 
tates it as an amorphous black powder. It is insoluble in di- 
lute acids. 
The lead compounds are very poisonous. The soluble salts 
have a sweetish, astringent taste. They are readily recognized 
by the following reactions : sulphuric acid precipitates white 
lead sulphate, which is blackened by hydrogen sulphide ; po- 
tassium iodide precipitates yellow lead iodide. 3 72 
INORGANIC CHEMISTRY. 
BISMUTH. 
Bi = 207. 
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. 145) ; the acid nature of 
the oxides passes into a basic. Antimony oxide (Sb2Os) is a 
base, while the higher oxide, Sb2Os, represents an acid anhy- 
dride. In bismuth, the metallic nature attains its full value. 
This is manifest in its inability to unite with hydrogen. Bis- 
muth trioxide is a base, and the pentoxide possesses a very 
feeble acid character, yielding indefinite compounds with the 
alkalies ; it behaves more like a metallic peroxide, and in its 
properties exhibits great similarity to lead peroxide. 
Bismuth usually occurs native, and in combination with sul- 
phur, as bismuthinite. To obtain the metal, the sulphide is 
roasted in the air, and the resulting oxide reduced with char- 
coal. 
Bismuth is a reddish-white metal, of specific gravity 9.9. 
It is brittle and may be easily pulverized. It crystallizes in 
rhombohedra. It fuses at 267° and distils at a white heat 
(about 1300°). It does not change in the air at ordinary tem- 
peratures. When heated it burns to bismuth oxide—Bi203. It 
is insoluble in hydrochloric acid, but dissolves in boiling sul- 
phuric acid with formation of sulphate of bismuth, and the 
evolution of sulphur dioxide. Nitric acid dissolves it readily 
in the cold. 
Water decomposes bismuth solutions in the same manner as 
those of antimony; insoluble basic salts are precipitated but 
these are not dissolved by tartaric acid. 
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 fuses at 230° and 
boils about 4350. It deliquesces in the air. Water renders 
its solution turbid, a white, crystalline precipitate of Bismuth 
Oxychloride—BiOCl—separating at the same time: 
Bids + H20 = BiOCl + 2HC1. 
The metalloidal character of bismuth is indicated by this 
reaction. . CHROMIUM GROUP. 
373 
The compounds BiBr3 and Bil3 are very similar to bismuth 
chloride. All three combine with many metallic haloid salts 
to form double halogen derivatives. 
Halogen derivatives of pentavalent bismuth are unknown. 
Bismuth Oxide—Bi2Os—prepared by burning bismuth or 
heating the nitrate, is a yellow powder, insoluble in water and 
the alkalies. 
Normal bismuth hydroxide—Bi(OH)3—is not known in 
a free state. Potassium hydrate added to a bismuth solution 
precipitates a white amorphous metahydrate—BiO.OH. 
Chlorine conducted through a concentrated potassium 
hydrate solution in which bismuth oxide is suspended pre- 
cipitates red bismuthic acid (Bi03H or Bi2H407), which 
when gently heated becomes Bi2Os, bismuthic oxide. Strong 
ignition converts the latter into Bi203 and 02: hydrochloric 
acid dissolves it to bismuth chloride, with liberation of 
chlorine. 
Bismuth Nitrate—Bi(N08)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 
fN°3 (NO, 
precipitation of white, curdy basic salts: Bi < N03and Bi -< OH. 
(OH (OH 
The precipitate is employed in medicine under the name of 
Magisterium bismuthi {subnitrate). 
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 
hydrogen sulphide. Unlike antimony and arsenic sulphides, 
it does not form sulpho-salts. 
The alloys of bismuth are nearly all readily fusible. An 
alloy of 4 parts Bi, 1 part Cd, 1 part Sn and 2 parts Pb, fuses 
at 65° (Wood’s metal). The alloy of 2 parts bismuth, 1 part 
lead and 1 part tin (Rose’s metal) fuses at 940. 
CHROMIUM GROUP. 
We observed that a group of more metallic analogous ele- 
ments attached itself to the metalloidal elements, carbon, 
silicon and tin (p. 236); and further that there was an analo- 
gous group of metallic elements corresponding to the metal- 374 
INORGANIC CHEMISTRY. 
loidal group of phosphorus (p. 224). We now meet a group 
of metals, consisting of chromium, molybdenum, tungsten, 
and probably uranium, that bears a like relation to the ele- 
ments of the sulphur group (see Periodic System of the Ele- 
ments). The resemblance of these elements to sulphur and 
its analogues is plainly manifest in their highest oxygen deriv- 
atives (see also manganese). As the elements of the sulphur 
group in their highest oxygen compounds-are hexavalent, so 
chromium, molybdenum, tungsten and uranium form acid 
oxides—Cr03, Mo03, W03, Ur03. Many of the salts corre- 
sponding to these are very similar to and isomorphous with 
the salts of sulphuric acid. Sodium chromate, like sodium 
sulphate, crystallizes with 10 molecules of water; the potas- 
sium salts of both groups form isomorphous mixtures; their 
magnesium salts, as well as that of tungstic acid, have the same 
constitution : 
MgS04 7H2Q and MgCr04 -f* 7H20. 
Corresponding to the acid oxides are the chlorine derivatives: 
S02C12, Cr02Cl2, Mo02C12, MoOC14, WOCl4 and WCI6, 
which are perfectly analogous, so far as chemical deportment 
is concerned. Besides these highest oxides the elements of 
the sulphur group form the less 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 tetravalent, possess an undetermined, neither acid nor 
basic, character. 
The most important basic oxide of chromium is its sesqui- 
oxide. This affords salts with the acids, and they are per- 
fectly similar to those of the sesquioxides of iron (Fe203), 
manganese (Mn203), and aluminium (A1203) (p. 331). 
Since the vapor density of ferric chloride declares its formula to be 
Fe2Cl6, we assume that in their sesquioxide compounds Cr, Mn, and Fe 
are tetravalent, and that these contain a hexavalent group consisting of two 
atoms of metal: 
IV IV 
=Cr — Cr= 
IV IV 
CI3Cr — CrCl3 etc. 
We can, however, conceive these derivatives, as well as those of alumi- 
nium, as derived from trivalent metallic atoms, and then make use of the 
simple, instead of the double formulas (CrX3 and FeX3). (Compare pp. 
340 and 348.) CHROMIUM. 
375 
Finally, compounds of chromium, CrX2, are known in 
which the metal figures as a dyad. These so-called chro- 
mous compounds are very much like (p. 330) the derivatives 
of the metals of the magnesium group, especially the ferrous 
salts (FeX2). They are very unstable, and are oxidized by 
the air into chromic compounds. 
Salts of molybdenum and tungsten, corresponding to the 
states of lowest oxidation, are not known, because these metals 
occur as hexads in most of their derivatives. Uranium, which 
has the highest atomic weight of the group, shows some varia- 
tions from its analogues; these are explained, as in similar 
cases by its high atomic weight. 
i. CHROMIUM. 
Chromium is found principally as chromite in nature. This 
is a combination of chromic oxide with ferrous oxide— 
Cr203FeO—and occurs in North America, Sweden, Hungary, 
and in large quantities in the Urals. Crocoisite, or lead 
chromate (PbCr04), is not met with so frequently. Chromic 
iron is used almost exclusively for the preparation of all other 
chromium derivatives, as it is first converted into potassium 
chromate (see this) by fusion with potassium carbonate and 
nitrate. 
Metallic chromium may be isolated by the very strong igni- 
tion of the oxide with charcoal. It is more conveniently 
obtained by igniting a mixture of chromium chloride, potas- 
sium chloride or sodium chloride with zinc, in a closed cru- 
cible. The separated chromium dissolves in the molten zinc, 
and when the latter is dissolved in nitric acid the chromium 
remains behind as a gray, metallic, crystalline powder, of spe- 
cific gravity 6.8. It is very hard (cuts glass), and fuses with 
difficulty. When heated in the air it slowly oxidizes to chro- 
mic oxide; ignited in oxygen it burns with a bright light. 
It dissolves readily in hydrochloric and warm dilute sulphuric 
acid, with elimination of hydrogen; it is not altered by nitric 
acid. 
Three series of chromium compounds are known ; chromous 
—CrX2, chromic—Cr2X6, and the derivatives of chromic acid, 
called chromates. All chromium compounds are brightly col- 
ored, hence, the name chromium (from xpibtia, color). 
Cr = 52.4. INORGANIC CHEMISTRY. 
CHROMOUS COMPOUNDS. 
These are very unstable, and by oxidation pass readily 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 col- 
ored. The alkalies precipitate yellow chromous hydrate, Cr(OH)2, from 
it. This is readily oxidized. When heated it parts with hydrogen and 
water and becomes chromic oxide : 2Cr(OH)2 = Cr203 + H2 + H20. 
Chromic Chloride—Cr2Cl6, like A12C16, isobtained by igni- 
tion of the oxide and charcoal in a current of chlorine. When 
raised to a redheat in this condition it sublimes in shiningviolet 
leaflets, which are transformed into chromic oxide by ignition 
in the air. Pure chromic chloride only dissolves in water after 
long-continued boiling; if, however, it contains traces of CrCl2, 
it dissolves readily at ordinary temperatures. Green crystals 
of Cr2Cl6 -f i2H20 separate from the green solution on evap- 
oration ; these deliquesce in the air. The same crystals may- 
be obtained from solutions of chromic hydrate, Cr2(OH)6, in 
hydrochloric acid. When they are dried intermediate oxy- 
chlorides, Cr2Cl4(OH)2 and Cr2Cl2(OH)4, and at lastCr2(OH)6 
result. 
Chromic Hydrate—Cr2(OH)6 or (CrOH)3. It is precipi- 
tated by ammonium hydrate from chromic solutions as a volumi- 
nous, bright brown, hydrous mass. The green precipitates pro- 
duced by sodium and potassium hydroxides, contain alkali that 
can not be removed even by boiling water. They dissolve readily 
with an emerald green color, in an excess of KOH or NaOH 
(slightly in ammonia), but are reprecipitated upon boiling 
their solutions. When it is heated to 200° in a current of hy- 
drogen, the product is the hydroxide,CrO.OH, which is a 
grayish-blue powder, insoluble in dilute hydrochloric acid. 
When chromium hydroxide is ignited, it becomes chromic 
oxide. 
Chromic Oxide—Cr203—is a green, amorphous powder. 
It is also formed by the ignition of chromium trioxide: 
CHROMIC COMPOUNDS. 
2Cr03 = Cr203 -f 3O. 
or of ammonium bichromate :— 
(NH4)2Cr20T = Cr203 + 4H20 + N2. CHROMIUM. 
3 77 
It may be obtained in black, hexagonal crystals, by conduct- 
ing the vapors of the oxychloride through a tube heated to 
redness: 
2Cr02Cl2 — Cr203._-f- 2C12 -(- O. 
Ignited chromic oxide is insoluble in acids. When fused with 
silicates, it colors them emerald green, and serves, therefore, 
to color glass and porcelain. 
Guignef s green is a beautifully green-colored chromium 
hydroxide, which is applied as a paint. It is obtained by ig- 
niting a mixture of one part potassium bichromate with three 
parts boric acid ; after treating the mass with water, which 
dissolves potassium borate, there remains a green powder, the 
composition of which corresponds to the formula: 
Cr20(0H), =• Cr203.2H20. 
The predominating properties of chromic oxide are basic, 
as it readily affords salts with acids; yet its basic nature, 
like that of all sesquioxides, is but slight, so that it does not 
afford salts with weak acids (p. 378). In addition to all this 
it possesses a slightly acidic character, and metallic salts 
are derived from it, generally from the hydrate, CrO.OH, 
which are analogous to the aluminates (p. 351). Salts, like 
(CrO.O)2Mg and (CrO.O)2Zn, can be obtained crystallized in 
regular octahedra by fusing chromic oxide with metallic ox- 
ides and boron .trioxide (as flux). Chromic iron is such a 
salt: 
Cr203,Fe0 = (CrO.O)2,Fe. 
Chromium Sulphate—Cr2(S04)3—is obtained by dis- 
solving the hydroxide in concentrated sulphuric acid. The so- 
lution, green at first, becomes violet on standing, and deposits 
a violet-colored crystalline mass. This may be purified by 
solution in water and precipitation by alcohol: This salt 
crystallizes from very dilute alcohol in bluish-violet octahedfa 
containing 15 molecules of water. If the aqueous solution of 
the violet salt be heated, it assumes a green color, because the 
salt breaks up into free acid and a basic salt which, upon evapo- 
ration, separates as a green amorphous mass, soluble in alco- 
hol. When the green solution stands, it reverts to the viplet 
of the neutral salt. The other chromic salts, the nitrate and 
the alum, behave in a similar manner. 378 
INORGANIC CHEMISTRY. 
Chromium sulphate forms double salts with the alkaline sul- 
phates—the chromium alums (p. 354). 
Potassium Chromium Alum—Cr2(S04)3K2S0424H20 
—crystallizes in large, dark violet octahedra. It may be most 
conveniently prepared by acting upon a solution of potassium 
bichromate mixed with sulphuric acid, with sulphur dioxide: 
K2CrA + H2S04 + 3S02 = Cr2(S04)3‘.K2S04 + H20. 
At 8o° the .violet solution of the salt becomes green, and on 
evaporation affords an amorphous green mass. 
As chromium hydroxide possesses only a slightly basic nature, 
salts with weak acids, like C02, S02,H2S (see Aluminium, p. 
353) do not exist. Therefore, the alkaline carbonates and 
sulphides precipitate chromium hydroxide 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. 
DERIVATIVES OF CHROMIC ACID. 
In its highest oxygen derivative, Cr03, chromium possesses 
a complete metalloidal, acid-forming character. Chromic acid, 
H2Cr04, is perfectly analogous to sulphuric acid, H2S04, but 
has not been obtained free, since when liberated from its salts 
it at once breaks up into the oxide-and water: 
H2Cr04 = CrOs + H20. 
The chromates are often isomorphous with the sulphates 
(p. 374). Polychromates also exist, and are derived from 
polychromic acids produced by the condensation of several 
molecules of the normal acid (see Disulphuric acid, p. 191): 
Potassium chromate. 
K2Cr04, 
Potassium bichromate. 
K2Cr207, 
Potassium trichromate. 
K2Cr3O10, etc. 
The constitution of these salts is expressed by the following 
formulas: CHROMIUM. 
379 
/OK 
Cr02( 
>° 
0rO2< 
X)K 
/OK 
CrO/ 
>° 
Cr02<( 
>o 
CrO,< 
xOK 
CrO 
LrU!\OK 
The free polychromic acids are not known, because as soon 
as they are separated from their salts, they immediately break 
up into the acid oxide and water: 
H2Cr3O10 = 3Cr03 + Ha0. 
The polychromates are frequently, but incorrectly, called 
acid salts; true acid or primary salts, in which only one H 
atom is replaced by metal (Cr04KH), are unknown for chromic 
acid. 
The salts of normal chromic acid are mostly yellow-colored, 
while the polychromates are red. The latter are produced 
from the former by the action of acids: 
Conversely, by the action of the alkalies, the polychromates 
pass into the normal salts: 
2K2Cr04 + 2HNO3 = K2Cr207 + 2KNOs + H20. 
K,Cr207 + 2K0H = 2K2Cr04 + 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: 
Cr04K2 CrOs = K2Cr207. 
When there is an excess of acid the anhydride (Cr03) is set 
free. 
Chromium Trioxide-j—acid Anhydride—Cr03. 
It consists of long, red, rhombic needles or prisms, obtained by 
adding.sulphuric acid to a concentrated potassium bichromate 
solution. The crystals deliquesce in the air and are readily 
soluble in water. When .heated, they blacken, melt, and at 
about 250° decompose into chromic oxide and oxygen : 
2Cr03 = Cr2Os + 3O. 
Chromium trioxide is a powerful oxidizing agent, and de- 
stroys organic matter; hence its solution cannot be filtered 
through paper. When alcohol is poured on the crystals, de- 
tonation takes place, the alcohol burns, and green chjomic 
oxide remains. By the action of acids, e.g., sulphuric, the 
trioxide deports itself like a peroxide; oxygen escapes and a INORGANIC CHEMISTRY. 
chromic salt results. When heated with concentrated hydro- 
chloric acid chlorine is evolved : 
2Cr03 + 12HCI = Cr2Cl6 + 6H,0. + 3Clr 
Reducing substances, like sulphurous acid and hydrogen sul- 
phide, convert chromic acid into oxide :— 
2Cr03 + 3H2S = Cr203 + 3H20 -f 3S. 
Chromate of Potassium—K2Cr04—is obtained by add- 
ing potassium hydrate to potassium bichromate. It forms 
yellow rhombic crystals, isomorphous with potassium sulphate 
(K2S04); isomorphous mixtures crystallize out from the solu- 
tion of the two salts. 
Bichromate of Potassium—K2Cr207—called acid potas- 
sium chromate, is manufactured on a large scale, and bears 
the name red chromate of potash in commerce. It is obtained 
by igniting pulverized chromite, Cr2OsFeO, 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 acetic or nitric acid 
(see p. 379), when potassium bichromate crystallizes from the 
concentrated solution. 
2Cr03 -f- 3S02 -|- 311,0 — Cr203 -f" 3S04H2. 
In practice the above method is advantageously replaced by the following: 
The pulverized chromite is ignited, together with lime, in furnaces allowing 
air access. Caloium chromate (CaCr04) (together with ferric oxide) is 
produced. This dissolves in dilute sulphuric acid to bichromate, CaCr207; 
the latter is converted by potassium carbonate into potassium 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 a very high 
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 -f Cr203 -f 03. 
K2Cr207 + 4H2S04 = Cr2(S04)3.K2S04 + 4H20 + 3O. 
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 + ioH20—forms deli- 
quescent crystals, and is analogous to Glauber’s salt (Na2S04 
+ ioH20). Barium and Strontium Chromates—BaCrG4 CHROMIUM. 
381 
and SrCr04—are almost insoluble in water. Calcium Chro- 
mate—CaCr04—is difficultly soluble m water, and crystal- 
lizes like gypsum with two molecules of water. The magnesium 
salt, MgCr04 7H20, dissolves readily and corresponds to 
Epsom Salt. The chromates of the heavy metals are insoluble 
in water, and are obtained by transposition. 
Chromate of Lead—PbCr04— is obtained by the pre- 
cipitation of soluble lead salts with potassium chromate. It is 
a yellow amorphous powder which serves as a yellow paint. 
When heated it melts undecomposed, and solidifies to a brown, 
radiating, crystalline mass. It oxidizes and easily decomposes 
all the carbon compounds, and is therefore used in their 
analysis. In nature lead chromate exists as crocoisite. 
The oxide, C1O2, called peroxide, may be obtained by the careful igni- 
tion of chromium trioxide, and is, most likely, a salt-like compound : 
Cr203.Cr03 or CrO.CrO,. 'Its hydrate is precipitated from chromium solu- 
tions upon the addition of potassium chromate. On warming the peroxide 
with hydrochloric acid, chlorine is evolved. 
Chromic Acid Chloranhydrides.—Chromic acid forms 
chloranhydrides similar to those of sulphuric acid (p. 192). 
Corresponding to S02C12, we have chromyl chloride, Cr02Cl2; 
and for the first sulphuric acid chloranhydride, S02 j is 
the salt, Cr02 j : 
vi /Cl 
Cr°2\Cl 
vi /Cl 
Cr°2\OK 
VI /OK 
Cr02\OK, 
Chromyl Chloride — Cr02Cl2 — Chromium oxychloride, 
is produced by heating a mixture of potassium bichromate and 
sodium chloride with sulphuric acid; it distils over as a dark 
red liquid, of specific gravity 1.92, and fumes strongly in the 
air. It boils at 1160; its vapor density equals 77 (H = 1), 
corresponding-to the molecular formula Cr02Cl2= 154. Chro- 
myl chloride has a strong oxidizing action. With water it is 
decomposed according to the following equation : 
Cr02Cl2 + H20 = Cr03 + 2HCI. 
If potassium bichromate be heated with concentrated hy- 
drochloric acid, the salt, Cr02 j , called potassium chloro- 
chromate crystallizes from the solution oa cooling. Heated 382 
INORGANIC CHEMISTRY. 
to ioo° it gives up chlorine. It is decomposed by water into 
potassium chloride and chromium trioxide. 
The chromium compounds can be readily recognized by 
their color. The following reaction is very characteristic for 
chromic acid: On adding hydrogen peroxide to a solution of 
chromium trioxide, or the acidified solution of a chromate, 
the red liquid is colored blue. The nature of the compound 
causing this coloration is not known; it is usually assumed to 
be a higher oxide of chromium. On shaking the blue solution 
with ether, the latter withdraws the blue compound and is 
beautifully colored in consequence. The ethereal solution is 
somewhat more stable than the aqueous. Both are gradually 
decolorized, with liberation of oxygen. 
2. 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. When heated in the air it oxidizes to molyb- 
denum trioxide. It is soluble in concentrated sulphuric and nitric acids. 
It is also converted by the latter into insoluble MoOa. 
Like chromium, molybdenum affords compounds of the forms MoX2, 
MoX4 and MoX8; besides which derivatives are known in which it 
appears to act as a pentad and also as a triad. 
Molybdenum Dichloride—MoCI2—resulting from the trichloride, 
MoC)3,when heated in a stream of carbon dioxide (together with MoC14 ', 
is a bright yellow, non-volatile powder. It is converted by potassium 
hydrate into the hydrate, Mo(OH)2, a black powder. 
.Molybdenum Trichloride— MoC13 or Mo2CI8—produced by gentle 
heating (at 250°) of MoC15 in a current of H or C02, is a reddish-brown 
powder, which, when strongly ignited, yields a dark-blue vapor. It dis- 
solves with a beautiful blue color in concentrated sulphuric acid, upon heat- 
ing, with an emerald green color. Potassium hydrate converts it into the 
hydrate, Mo^OH)3 or Mo2iOH)6, which forms salts with acids. The 
ignition of the hydrate affords the black oxide, Mo203. Strong heating 
of the trichloride in a current of C02 leaves MOCl2 and it sublimes. 
Molybdenum Tetrachloride—MoC14—is a brown, crystalline pow- 
der, which appears to break up by evaporation into MoCl5 and MoC13. It 
yields a hydrate with ammonium hydrate, forming salts with acids.. The 
brown solution of the salts readily assumes a blue color by oxidation in 
the air. The igni.ion of the hydrate leaves-the dioxide, M0O2, which is 
converted by nitric acid into the trioxide, MoOs. Molybdenum disulphide, 
M0S2, is produced by the ignition of the trisulphide, MoS3, away from air. 
Mo = 95.8. MOLYBDENUM. 
383 
It is a shining black powder, which occurs native as molybdenite, in 
hexagonal, graphite-like crystals, with a specific gravity of 4.5. 
Molybdenum Pentachlcride—MoC15—is prepared by heating MoS2 
or molybdenum in dry chlorine gas. It is a metallic, shining, black, crystal- 
line mass, fusing at 1940 and distilling at 268° ; its vapor density equals 
136, corresponding to the molecular formula ’MoC15 = 272.6. It fumes 
and deliquesces in the air, and dissolves in water with hissing. Its aqueous 
solution has a brown color. It dissolves in absolute alcohol and ether with 
a dark-gneen color. 
The kexachloride, 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 
ioo° and yielding a dark-red vapor. 
Bromine forms perfectly analagous compounds with molybdenum. 
Molybdenum Trioxide— MoOa results on roasting metallic molyb- 
denum 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 carbon- 
ates, salts are produced, partly derived from the normal acid, H2Mo04, and 
partly from the polyacids, and correspond to the polychromates: 
The ammonium salt—(NH^)2Mo04—is obtained by dissolving the 
trioxide in concentrated ammonfum 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 (NH416Mo}024-(- 
4H2O crystallizes out. Both salts are decomposed by heat, leaving 
molybdenum trioxide. 
Hydrochloric acid added to a concentrated solution of a molybdate 
precipitates 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 (formation of sesquioxide), in conse- 
quence of the formation of lower oxides (like Mo308 =-2Mo03, M0O2), 
and finally brownish-red and yellow, when a suboxide, Mo5Ot = 2Mo203, 
MoO, is produced. Potassium permanganate converts all these lower 
oxides into molybdic acid. 
Molybdic acid can also form polyacids with phosphoric and arsenic 
acids, e.g., H3P04.iiMoQ3. These complex phosphomolybdic acids are 
distinguished by the fact that they form salts insoluble in dilute acids 
with the metals of the-potassium group, with ammonia and with organic 
bases. On adding a solution containing phosphoric (or arsenic) acid to 
the nitric acid solution of ammonium molybdate, there is produced a 
yellow crystalline precipitate.of arrtmonium phospho-molybdate—(NH4)3 
P04. iiMo03 -)- 6H20. This reaction serves for the detection and sepa- 
ration of phosphoric aoid. 
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. 
K2Mo04, K2Mo207, K2Mo3O10, Na2Mo4013, K6Mo7024, etc. 384 
INORGANIC CHEMISTRY. 
3. TUNGSTEN. 
Tungsten is found in nature in the tungstates: as wolframite, FeW04, 
as scheelite, CaW04, and as stolzite, PbW04. 
The metal is obtained, like molybdenum, by the ignition of the oxides 
or chlorides in a stream of hydrogen, in the form of a black powder, or in 
steel-gray crystalline leaflets, having a sp. gr. 19.1. It is very hard and 
difficultly fusible. It becomes trioxide when ignited in the air. 
Tungsten forms the following chlorides; WC12, WC14, WC15 and WC16. 
The Dichloride—WC12—arises by strong ignition of WC16 and WC14 
in a current of carbon dioxide, and is a bright gray, non-volatile mass. 
The Tetrachloride—WC14—obtained by gentle ignition of WC16 and 
\VC15, in a current of hydrogen or carbon dioxide, is grayish-brown and 
upon sublimation decomposes into WC12 and volatile WC15. It forms a 
brown oxide (W02) with w'ater. 
The Pentachloride—WC15—is obtained by the distillation' of WC16 
in a current of hydrogen or carbon dioxide, and consists of shining, 
black, needle-Like crystals It fuses at 248° and boils at 2750, forming a 
dark brown vapor, with the density 180 (WC15 = 360.4). It affords an 
olive-green solution and a blue oxide, W205, with water. It dissolves 
with a deep blue color in carbon disulphide. ’ 
Tungsten Hexachloride—WC16—is produced when the metal or a 
mixture of wollramite with carbon is heated in a current of chlorine. It 
forms a dark violet, crystalline mass, fusing at 2750 and boiling at 346°. 
The vapor density equals 198 (WC16 = 395.8). It dissolves in carbon 
disulphide with a reddish-brown color; it forms W03 with water. 
The Oxychloride—WC140— consists of red crystals, fusing at 210° 
and boiling at 2270; its vapor density equals 170 (WC140 = 341). The 
Dioxychloride—WC1202—sublimes in bright yellow, shining leaflets. 
Tungsten Trioxide—WOs—is thrown out of the hot solution of tung- 
states by nitric acid, as a yellow precipitate insoluble in acids, but dissolv- 
ing readily in potassium and sodium hydrates. Tungstic acid, WO(OH)4, 
is, however, precipitated from the cold solution, but on standing over sul- 
phuric acid it becomes W02(OH)2 and at ioo° passes into ditungstic 
acid, W207H2 = W2C>5(OH)2. 
When tungstic acid is reduced in hydrochloric acid solution by zinc it 
first becomes blue (formation of W2051 and then brown, when the salt of 
the dioxide, is- formed. Potassium permanganate oxidizes this to 
tungstic acid. 
The salts of tungstic acid are perfectly analpgous 30 the molybdates and 
are derived from the normal acid or the polyacids. The normal sodium 
salt, Na2W04 -)- 2H20, and the so-called meta-tungstate of sodium, 
Na2W4013 ioII20, are applied practically. Materials saturated with 
their solutions do not burst into a flame, but smoulder away slowly. 
The reduction of the tungstates (by fusion tin, etc.) affords pecu- 
liar compounds, e.g., K2W309 or K2W4012; these have various colors, 
possess metallic lustre, and are applied as tungsten bronzes. 
Tungstic acid also combines with phosphoric and arsenic acids, forming 
derivatives analogous to those of molybdic acid with the same acids. 
The metal is used in the manufacture of tungsten steel: a slight quan- 
tity of it increases the hardness of the latter very considerably. 
W =.183.6. URANIUM. 
385 
4. URANIUM. 
In nature it occurs chiefly as uraninite, a compound of uranic and 
qranous oxides—U02, 2U03 = U308. 
The metal is obtained by heating uranous chloride with sodium. It has 
a steel-gray color and a specific gravity of 18.7. When heated in the air 
it burns to uranous-uranic oxide. Its specific heat equals 0.0267, and its 
atomic volume is therefore 6.6. There are two serie? of uranium com- 
pounds. In the one, the metal is a tetrad UX4; these uranous or urano- 
compounds are very unstable, and pass readily into the uranic or deriva- 
tives of hexavalent uranium. Uranous oxide is of a basic nature, and only 
forms salts with acids. 
The compounds of hexavalent uranium are called the uranic com- 
pounds. U03 and U02<0H)2 have a predominant basic character, but 
are also capable of forming salts with bases which are called uranates. In 
the salts derived from acids, e.g., U02' N03)2 and U02S04, the group UO,2 
plays the role of a metal; it is called tiranyl, and its salts are termed 
uranyl salts. They may also be regarded as basic salts. 
Ur or U = 239. 
URANOUS COMPOUNDS. 
Uranous Chloride—UC14—is obtained by heating metallic uranium 
in a stream of chlorine, or uranous oxide in hydrochloric acid. It con- 
sists of dark green octahedra with metallic lustre. It volatilizes at a red 
heat, forming a red vapor, whose density agrees with the formula UC14. 
It deliquesces in the air, and dissolves with hissing in water. Uranous 
hydroxide remains when the solution is evaporated. 
Uranous Oxide—U02—is formed when the other oxides are heated 
in a current of hydrogen. It Is a'black .powder, which becomes uranous- 
uranic oxide, U02.2U03, when heated in the air. 
Uranous oxide dissolves with a green color in hydrochloric and con- 
centrated sulphuric acids. Uranous sulphate, U(S0412 + 8H20, consists 
of green crystals. From the salts the alkalies precipitate the voluminous, 
bright green uranous hydroxide, U(OH)4, which becomes brown on ex- 
posure. 
Uranium Hexachloride—UC16—has not been obtained, but the oxy- 
chloride, U02C12 (Uranyl chloride), exists; it is obtained by heating U02 
in dry chlorine gas, or by the evaporation of uranyl nitrate with hydro- 
chloric acid. It is a yellow crystalline mass, deliquescing in the air. 
Uranic Oxide, U03, or Uranyl oxide, U02 O—is a yellow powder, 
and is obtained by heating uranyl nitrate to 250°. When warmed with 
nitric acid it becomes uranyl hydrate or tiranic acid, UO^(OiI)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 
HEXAVALENT URANIUM COMPOUNDS 386 
INORGANIC CHEMISTRY. 
readily soluble in water and alcohol. On adding sulphuric acid to the 
solution, Uranyl sulphate—UCUjSO* 6H20—crystallizes out, on evap- 
oration, in lemon yellow needles. 
If sodium or potassium hydrate be added to the solutions of uranyl salts, 
yellow precipitates of the uranates—U2O.K2 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 (ura- 
nium glass) and porcelain. The uranates can be obtained in crystalline 
form, by igniting uranyl chloride with alkali chlorides in the presence of 
ammonium chloride. Zinc and sulphuric acid reduce uranic to uranous 
compounds. 
The so-called uranic-uranous oxide, which constitutes uraninite, and is 
formed by the ignition of the other oxides in the air, must be viewed as 
• VI IV 
uranous uranate—2U0S.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 blp.ck color to glass and porcelain. 
Besides these compounds, in which uranium appears to be tetravaJent 
and hexavalent, it also affords a pentachloride, UC15, like molybdenum and 
tungsten. The same results on conducting chlorine gas over a moderately 
heated mixture of carbon with one of the uranium oxides. It consists of 
dark needles, which, in direct light, are metallic green, but in transmitted, 
ruby red. It deliquesces in the air to a yellowish-green liquid ; upon heat- 
ing it is dissociated into UC14 and Cl (at i2O0-235°). 
There is also a tetroxide, U04, which, like the trioxide, U03, yields 
salts with the bases. 
MANGANESE. 
According to its atomic quantity, manganese bears the same 
relation to the elements of the chlorine group as chromium 
to the elements of the sulphur group. The relationship mani- 
fests itself distinctly in the higher states of oxidation. Per- 
manganic oxide, Mn207, and acid, HMn04, are perfectly anal- 
ogous to C1207 (or I207) and HC104. The permanganates 
and the perchlorates are very similar, and for the most part 
are isomorpliQus. The manganese in them appears to be hep- 
tavalent, like the halogens in their highest state of oxidation. 
The similarity of manganese to the halogens is restricted to 
this one point of resemblance. In the rest of its derivatives, 
manganese shows great resemblance to the elements standing 
in the same horizontal series of the periodic system, viz., with 
iron and chromium (p. 347). . Like these two elements, it 
forms three series of compounds. 
Mn = 54.8 (55.0). MANGANESE. 
1. In the manganous derivatives—MnX2—the metal is di- 
valent. These salts are the more stable, and comprise the 
most common manganese compounds. They resemble and 
are usually isomorphous with the ous salts of iron and chro- 
mium, 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; they are, however, less stable, and easily reduced to the 
manganous state. Their composition is due to the tetravalent 
nature of manganese (p. 374). 
3. The derivatives of manganic acid—H2Mn04 = Mn02 
(OH)2, in which manganese is hexavalent—are analogous to 
those of ferric (H2Fe04) and chromic (H2Cr04), and, 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. 330). In its ous condition, manganese, 
like the elements of the magnesium group, has a rather strong 
basic character, which diminishes considerably in the ic state. 
Hexavalent manganese has a metalloidal acidic character,.and, 
in manganic acid, approaches sulphur. By the further addition 
of oxygen, manganese finally (in permanganic acid) acquires 
the metalloidal character of the halogens. We have already 
noticed that many other metals, especially chromium and 
iron, exhibit a similar behavior. Osmium tetroxide, 0s04, 
wholly resembles the halogens. 
On the other hand, the metalloidal and the weak basic metals acquire a 
strong basic, alkaline character, by the addition of hydrogen, or hydro- 
carbon groups (CH3,C2H5b The groups, NH4 (ammonium), Pi CH3)4 
(tetramethylphosphonium), S(C2H5)3 (triethylsulphine), Sn(C2H-)3 (tin tri- 
ethyl), etc., are of metallic nature, because their hydroxides, P(CH3'4.OH, 
S(C2H5)3.OH, Sn(C2H5)3.OH, are perfectly similar to the hydroxides 
(KOH,NaOH) of the alkali metals. 
Manganese is widely distributed in nature. It is found na- 
tive in meteorites. Its most important ores are pyrolusite, 
Mn02, hausmannite, Mn304, braunite,Mn203, manganite, 
Mn203.H20, and rhodochroisite, MnCOa. 
Metallic manganese is obtained by igniting tlie oxides with 
charcoal. It has a grayish-white color, is very hard, and dif- 
ficultly fusible: specific gravity 7.2. It oxidizes readily in 
moist air. It decomposes water on boiling, and, when dis- 
solved in acids, forms manganous salts. 388 
INORGANIC CHEMISTRY. 
The heat of formation of the most important manganese compounds 
corresponds to the symbols : 
(Mn,0,H20) = 94.7 (Mn,Cl2) = 111.9 (Mn,Cl2,Aq.) — 128.0 
(Mn,02,H20) = 116.2 (Mn,S,04) = 249.8 (Mn,04,K) = 194.8 
MANGANOUS COMPOUNDS. 
Manganous Oxide—MnO—results from ignition of the 
carbonate, with exclusion of air, and by heating all manganese 
oxides in hydrogen. It is a greenish, amorphous powder, 
which, in the air, readily oxidizes to Mns04. 
Manganous Hydroxide—Mn(OH)2—is a voluminous,’ 
reddish-white precipitate, formed by the alkalies in manganous 
solutions. When exposed to the air, it oxidizes quickly to 
manganic hydrate, Mn2(OH)6. 
Manganous salts usually have a pale, reddish color, and are 
formed by the solution of manganese or manganic oxides in 
acids. 
Manganous Chloride—MnCl2—crystallizes with four 
molecules of water in reddish tables. On drying, it is decom- 
posed with separation of hydrochloric acid. Anhydrous 
manganous chloride is prepared by igniting the double salt 
MnCl22NH4Cl + H20 (see Magnesium Chloride) or by heat- 
ing manganese oxides in hydrochloric acid gas; it is a crys- 
talline, reddish mass, which deliquesces in the air. 
Manganous Sulphate—MnS04—crystallizes below + 6° 
with 7 molecules of H20 (like magnesium and ferrous sul- 
phates), and at ordinary temperatures with 5H20 (like copper 
sulphate) ; the last molecule of water does not escape until 
2oo°. It forms double salts with the alkaline sulphates, e.g., 
MnS04. K2S04 + 6H20. 
Manganous Carbonate—MnC03—exists in nature as 
rhodochroisite, and is precipitated by alkaline carbonates from 
manganous solutions, as a white powder, which. turns brown 
on exposure. 
Manganous Sulphide—MnS—is found in nature as ala- 
bandite or manganese blende. Alkaline sulphides precipitate 
a flesh-colored sulphide from manganese solutions. It be- 
comes brown in the air. MANGANESE. 
389 
Manganic Oxide—Mn203, manganese sesquioxide—is a 
black powder produced by the ignition of the ‘manganese 
oxides in a current of oxygen gas. It occurs as Braunite in 
dark-brown quadratic crystals. 
Manganic Hydroxide—Mn2(OH6) or Mn(OJI)3, man- 
ganic hydrate—is precipitated by ammonium hydrate from 
manganous solutions containing ammonium chloride as a dark- 
brown mass. It dissolves in cold hydrochloric acid to a dark- 
brown liquid, containing, in all probability, manganic chlo- 
ride, MnCl, or Mn2Cl6. When this is heated it decomposes 
into MnCl2 and chlorine. 
Manganite, occurring in iron-black crystals, is the hydrox- 
ide. Mn202(0H)2 or MnO.OH. 
Manganous-manganic Oxide—Mn304 = Mn0,Mn203. 
It constitutes the mineral hausmannite, crystallized in dark- 
gray quadratic octahedra, and is obtained as a reddish-brown 
powder by the ignition of all other manganese oxides in the 
air. It reacts with hydrochloric acid, according to the equa- 
tion : 
MANGANIC COMPOUNDS. 
Mn304 + 8HC1 = 3MnCl2 + 4H20 + Cl2. 
Since manganic oxide is quadratic jn its crystallization, while all other 
sesquioxides (like corundum and hematite) are rhombohedral, and since 
the first is decomposed by dilute nitric and sulphuric acids into M11O2 and 
a-manganous salt, it has been generally supposed that manganic oxide is 
not a sesquioxide, but rather a’compound of the dioxide with manganous 
oxide: 
IV /°\ 
Mn04. MnO = MnO< >Mn. 
xv 
Hausmannite is quadratic, while other metallic oxides (the spinels, p. 
352 and p. 377, and magnetite,.Fe304) are isometric; therefore the former 
is not considered a compound of manganese sesquioxide and protoxide: 
MnO.CX 
Mn,Oo.MnO = . )Mn, 
MnO.O/ 
but as manganous oxide and the dioxide : 
IV o Mru ' 
MnO,.2MnO = MnO<; \0. 
\0.Mn/ 
This is shown by its behavior toward dilute nitric and sulphuric acids, 
which decompose it into manganese dioxide, and two molecules of man- 
ganous oxide. Chrysoberyl, unlike other spinels, is trimetric, and other re- 
actions clearly prove (chiefly their deportment with concentrated sulphuric 390 
INORGANIC CHEMISTRY 
acid) that manganic and mangano-manganic oxides are to be regarded as 
sesquioxide derivatives. 
Manganic oxide, like the other sesquioxides, is a very feeble 
base, which does not afford salts with dilute or weak acids, and 
by separation of oxygen reverts to the manganous condition. 
Its salts are very unstable. 
Manganic Sulphate—Mn2(S04)3—is obtained by the so- 
lution of manganic oxide, hydroxide, or better, manganous- 
manganic oxide in concentrated sulphuric acid. When the 
last oxide is employed manganous sulphate also results. The. 
best procedure is to heat the hydrate of manganese dioxide (see 
below) with concentrated sulphuric acid to 1680, when the 
sulphate will separate as an amorphous, dark-green powder. 
It dissolves with a dark-red color in a little water. It 
forms alums, with potassium and ammonium sulphates—e.g., 
Mn2(S04)3,K2S04 + 24H20. Much water will decompose 
these with exhaustion of manganic hydroxide. 
Manganese Dioxide—Mn02—peroxide. This is the min- 
eral 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 igni- 
tion into manganous-manganic oxide : 
3Mn0.2 ;= Mn304 -f- 20. 
It is used for making oxygen. Manganous oxide results at a 
white heat. Chlorine escapes when it is warmed with hydro- 
chloric acid: 
Mn02 ■+ 4HCI = MnCl2 -f 2H20 + Cl2. 
The dioxide may be obtained artificially by heating mangan- 
ous nitrate to 150-160°. Its hydrates—Mn02,H20 and Mn02, 
2.H20—are produced on adding a hypochlorite to the solution 
of a manganous salt, or if chlorine be conducted through a solu- 
tion of manganese containing sodium carbonate, or by adding 
KMn04 to a boiling solution of a manganous salt. The pre- 
cipitated dioxide dissolves in cold hydrochloric acid, without 
liberating chlorine, as MnCl4 is probably formed; when heat 
is applied it breaks up into MnCl2atid Cl2. This deportment 
would indicate that manganese is. tetravalent in the dioxide. 
Manganese dioxide also unites with bases, affording the so- 
called manganites, e.g., Mn205Ba and Mn206Kg. MANGANESE. 
391 
Manganese peroxide (also Mn203 and Mn304), serves chiefly for the 
manufacture of chlorine gas, and it is, therefore, important from a technical 
point to estimate the quantity of chlorine which a given dioxide of manga- 
nese is able to set free. This is done by boiling the oxide with hydro- 
chloric acid, conducting the liberated chlorine into a potassium iodide 
solution, and determining the separated equivalent amount of iodine by 
means of sodium hyposulphite. Or the oxide is heated in a flask with 
oxalic and sulphuric acids, when the oxalic acid is oxidized to carbon 
dioxide, and from the quantity of this set free we can calculate the quan- 
tity of active or available oxygen in the manganese oxide". 
In the preparation of chlorine the manganese is found in the residue as 
manganous chloride. With the relatively high value of pyrolusite, it is 
important for trade that the peroxide be recovered from the residue. This 
regeneration is at present largely executed by the method proposed by 
Weldon, according to which the manganous chloride, containing excess of 
hydrochloric acid, is neutralized 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, Mn03Ca = 
MnC>2 CaO, is precipitated as a black mud : 
MnCl2 -j- 2CaO -f- O = Mn03Ca -)- CaCl2. 
The calcium chloride solution is run off, and the residual calcium man- 
ganite employed for the preparation of when it conducts itself as 
a mixture of MnC>2 T CaO. 
When oxygen compounds of manganese are heated in the 
air in contact with potassium hydrate, or, better, with oxi- 
dizing substances, like nitre or potassium chlorate, a dark 
green amorphous mass is produced, which dissolves in cold 
water, with a dark green color. When this solution is evapo- 
rated under the air-pump, dark green metallic rhombic prisms 
of potassium manganate—K2Mn04—crystallize out. This 
salt is isomorphous with potassium sulphate and chromate. It 
suffers no change by solution in potassium or sodium hydrate, 
but is decomposed by water, brown hydrated manganese dioxide 
separating, and the green solution of the manganate changing 
into a dark red solution of the permanganate, KMn04: 
COMPOUNDS OF MANGANIC AND PERMANGANIC ACID. 
3K2Mn04 + 3H20 = 2KMn04 +• Mn02.H20 + 4KOH. 
A similar conversion of the green manganate into red per- 
manganate occurs more rapidly under the influence of acids: 
3lC2Mn04 -f 4HNO3 = 2KMn04 -f Mn02 + 4KNO3 + 2H20. 
Owing to this ready alteration in color the solution of the 
manganate is called chameleon mineral. 392 
INORGANIC CHEMISTRY. 
Potassium Permanganate—KMnO-4—is best prepared 
by conducting C02 into the manganate solution until the 
green color has passed into a red. When the solution is con- 
centrated the salt crystallizes in dark-red. rhombic prisms 
isomorphous with potassium perchlorate,KC104. It is soluble 
in twelve parts of water at ordinary temperatures. 
The permanganate solution is a strong oxidizing agent, con- 
verting lower, oxygen compounds into higher, and in doing 
this it is reduced to a colorless manganous salt. When a 
permanganate solution is added to an acidulated ferrous solu- 
tion, the former is decolorized, and there results a faintly 
yellow-colored solution of ferrit and manganous salts: 
2KMn04 -f- ioFeS04 -f- 8H2S04 = 2MnS04 -(- 5Fe2(S04)3 -j- 
8H20 + K2S04. 
Hence the solution of this salt serves for the volumetric 
estimation of ferrous salts. 
In the same manner, the permanganate oxidizes and de- 
stroys many organic substances, therefore its solution cannot 
be filtered through paper'; it serves as a disinfectant. 
The permanganate is also reduced by hydrogen peroxide 
(p. ioi) ; the reaction proceeds according to the following 
equation: 
Mn207K20 -j- 5B2O2 — 2MnO -j- K20 -)- 5H20 5O2» 
the formation of oxides requires the presence of acids (sul- 
phuric-acid) for the completion of the reaction. 
The remaining permanganates are similar to and isomorph- 
ous with the perchlorates. The sodium salt is very soluble in 
water, and does not crystallize well. 
Very cold sulphuric acid added to dry permanganate 
causes the separation of Manganese Heptoxide—Mn207— 
an oily, dark-colored liquid. By careful warming it is con- 
verted into dark violet vapors, which explode when heated 
rapidly. Manganese heptoxide has a violent oxidizing action ; 
paper, alcohol and other organic matter are inflamed by mere 
contact with it. 
METALS OF GROUP VIII. 
Of the known elements, those standing in the eighth column 
of the periodic system remain for consideration (p. 244). METALS OF GROUP VIII. 
393 
Fe = 55.9 Co — 58.6 Ni = . 58.6 
Ru i= 103.0 Rh •= 104.0 Pd == 106.2 
Os = 195.0* Ir == 192.5 Pt = 194.3 
These elements are the middle members of the three great 
periods, and they have no analogues in the two short periods 
(pages 241, 243). 
As regards both atomic weights and physical and chemical 
deportment, these elements 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, 
p. 252). 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 gravi- 
ties. They are, therefore, usually arranged in groups, and 
distinguished as, (1) the iron group (Fe, Co, Ni), with the 
specific gravity 7.8-8.6; (2) the group of the light platinum 
metals (Ru, Rh, Pd), with the specific gravity 11.8-12.1, and 
(3) the group of the heavy platinum metals (Os, ir, Rt), with 
the specific gravity 21.1-22.4. 
On the other hand, the homologous elements (Fe,* Ru, Os; 
Co, Rh, Ir; and Ni, Pd, Pt ) show a like similarity in their chem- 
ical properties, as do the other homologous groups, and there- 
fore 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. 330). We know 
that the metals of group VI.-(chromium, molybdenum, tung- 
sten) and of group VII. (manganese) form the highest oxides 
(Me03 and Me207) having an acidic nature. In the adjacent 
elements of group VIII. (iron, ruthenium, and osmium) we find 
salts: 
Fe04K2, Ru04K2, Os04K2, 
derived from the unstable trioxides FeOs, RuOs and OsOs. 
This acid-forming function disappears in the following mem- 
bers, Co, Rh, Ir, and Ni, Pd, Pt; their chemical valence dimin- 
ishes rapidly and they attach themselves to Cu, Ag and Au. 
Consequently the whole physical and chemical deportment 
* As already mentioned (p. 245), osmium, from its position in the 
periodic system, must have a lower atomic weight than that determined 
experimentally (heretofore 198: recently 195'. But as it shows the same 
relation in its whole deportment to Ir and Pt as Ru to Rh and Pd and Fe to 
Co and Ni, it is extremely probable that its atomic weight is not correctly 
determined, and that- it will show itself somewhat less than that of iridium 
(192.5). 394 
INORGANIC CHEMISTRY. 
of the 9 elements about to be considered is governed by their 
position in the periodic system. 
As mentioned on pp. 242, 247 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, and may be observed in the following tabula- 
tion of both classes of oxides of the middle members of the gn at periods: 
V 
V2°5 
VI 
Cr03 
VII 
Mn207 
VI 
Fe03 
IV 
Co203 
ii 
II 
CuO 
ii 
III 
IV 
III 
ii 
ii 
ii 
ii 
NiO 
i 
ZnO 
Ga203 
II 
v2o3 
CrO 
MnO 
FeO 
CoO 
Cu20 
. — 
V 
VI 
• 
VI 
IV 
IV 
Nb,(X 
Mo03 
RuOg 
Rh203 
ii 
I 
ii 
m 
Sn02 
l 0 
III 
II 
II 
IT 
PdO 
Ag20 
CdO 
ln203 
II 
(Nb203) 
MoO 
— . 
RuO 
RhO 
SnO 
V 
VI 
VI 
IV 
IV 
III 
ii 
hi 
IV 
Ta205 
wo3 
— 
0s03 
Ir02 
Pt02 
Au203 
HgO 
ti2o3 
Pb02 
hi 
II 
II 
ii 
II 
i 
I 
I 
ii 
(Ta20)3 
(WO) 
— 
OsO 
IrO 
PtO 
Au20 
Hg20 
T120 
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 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 acidie character, and the derivatives of ferric acid 
(H2Fe04) are perfectly similar to those of chromic and man- 
ganic acids (p. 387); they are, however, less stable than the 
latter. Their analogues with cobalt and nickel are not known. 
The ferric compounds—Fe2X6—containing the hexavalent 
VI 
group Fe2 (p. 374) are much like the alumihium, chromic and 
manganic derivatives. They are generally isomorphous with 
them. They are characterized among iron salts by their 
relative stability. The highest oxides of cobalt are far less 
stable, and only a few double salts of this form are known, 
while the higher salts with nickel are unknown. 
Again, iron, cobalt, and nickel afford ous compounds, 
(FeX2,CoX2,NiX2) in which they appear to be dyads. They 
resemble the compounds of chromium, manganese, and copper 
of the same form, and those of the magnesium metals. The 
ferrous salts are not as stable as the ferric; they are readily 
oxidized to the latter. 
The cobaltous and nickelous compounds are quite stable, IRON. 
395 
and in this respect these metals ally themselves with copper 
and zinc. 
1. IRON. 
This metal, of such great practical importance, is very 
widely distributed in nature. It is found native on the earth’s 
surface almost exclusively in meteorites; it is, however, present 
in great masses in other worlds which (like the sun) are sur- 
rounded by an atmosphere of hydrogen. 
The most important iron ores are: magnetite (Fe304), 
hematite (Fe203), brown iron ore and limonite (hydrates of 
the oxide) and siderite (FeC03). These ores constitute 
almost the sole material for the manufacture of iron; the 
sulphur ores, like pyrite, are less adapted to this end. 
In commerce there are three varieties of iron: cast-iron, 
steel, and wrought-iron. Their chief chemical difference is 
in the variable quantity of carbon contained in them. 
Cast iron contains 3-6 per cent, carbon, in part chemically 
combined, and in part mechanically mixed in the form of gra- 
phite. When molten cast-iron is cooled rapidly it yields the 
so-called white iron, in which the greater portion of the carbon 
is chemically combined with the iron. R.has a whitish color, 
exhibits a granular crystalline structure upon fracture, and is 
very hard and brittle. Its specific gravity is 7.1. It fuses to 
a pasty mass about 1200° and is on this account not suited 
for castings. The chemically combined carbon in it can 
easily be removed by oxidation, and, therefore it is adapted 
for the manufacture of steel or wrought iron. 
When molten cast-iron is allowed to cool slowly, the greater 
part of the carbon in it separates in the form of small leaflets 
of graphite. The gray cast-iron produced in this way has a 
darker gray color, is not so hard and brittle, fuses more readily 
(about .1150°) than white cast-iron, and serves for the manu- 
facture of castings. Neither variety can be forged or welded, 
on account of its brittleness. 
Steel contains 0.8-1.8 per cent, of carbon, all of which is 
chemically 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°) then cast-iron, but 
easier than wrought-iron. When molten steel is rapidly cooled, 
it becomes very hard and brittle. In this process more car- 
bon is chemically combined. If cooled slowly, it is soft and 
Fe = 55-9- 396 
INORGANIC CHEMISTRY. 
malleable, and may be forged and welded. Welding becomes 
more and more difficult with the addition of carbon. 
Wrought-iron contains the least amount of carbon, 0.2- 
0.6 per cent. 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 fibrous texture, while the forged is fine-grained ; 
the former is more compact and tenacious. Wrought-iron 
fuses at a bright white heat (1500°). 
Metallurgy of Iron.—-The extraction of iron from its oxygen ores is 
based upon the reduction of the same by carbon at a red heat. In the 
oldest method, the ores were heated wdth carbon in wind furnaces; in this 
way the excess of air consumed the greater portion of the carbon, and the 
product was an iron poor in carbon, wrought-iron, a spongy mass, wdiich 
was then forged under the hammer. The present methods were adopted 
since the beginning of the previous century. According to these cast-iron 
is first prepared from the ores, and this afterwards converted into steel or 
wrought-irPn. The smelting of the ores is executed in large, walled blast 
furnaces, that permit the process to proceed without interruption. The 
furnaces are filled from openings above, wnth alternating layers of coal, 
broken ore and-fluxes containing silica and lime ; the latter facilitate the 
melting together of the reduced iron. The air necessary for the process 
is blown into the contracted portion of the furnace by means of a blast 
engine. The combustion of the coal affords carbon monoxide, which 
reduces the iron oxides to metal: 
Fe2©3 +.3 CO = 2Fe -f 3C02. 
As the reduced iron sinks in the furnace it comes in contact w ith the 
coal, takes up carbon and forms cast-iron, wdiich fuses as it sinks lower 
and flows into the hearth of the furnace. Protracted and strong heating 
converts the chemically united carbon into the graphitic form, and thus 
accelerates the formation of the gray cast-iron. The earthy impurities of 
the ores combine with the fluxes to a readily fusible slag, w'hich envelops 
the fused iron and protects it from oxidation. 
To convert the cast iron thus produced into steel or wrought-iron, 
carbon must be withdrawn from it. In making the the cast- 
iron is fu*sed in open hearths (refining process), or in reverberatory 
furnaces wfith 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 otlier admixtures, like silicon, 
sulphur, and phosphorus, present in small quantities, are oxidized. The 
wfrought-iron is then w'orked up by rolling, or under the iron hammers 
(bar-iron). 
Steel was formerly manufactured from wrought-iron (not cast-iron), 
by cementation. The iron bars, mixed wdth fine charcoal, were exposed 
to a red heat, when the iron took up carbon from the surface. The 
bars were then reforged, again heated with fine charcoal, and the 
process repeated until the mass became as homogeneous as possible 
(cementation steel). A more uniform steel is obtained if it be fused in 
crucibles (cast-steel). IRON. 
397 
At present, steel is chiefly prepared directly from cast-iron, by the 
method invented by Bessemer, somewhere in 1850. Itconsists in blowing 
air, under high pressure, into the molten iron, until the necessary amount 
of carbon has been consumed (Bessemer steel). 
Puddle steel is obtained from cast-iron by the puddling process, and is 
not so fully decarbonized. Uchatius steel is prepared by fusing cast-iron 
together with some iron ore and pyrolusite. 
The various and more recent processes’ for manufacturing steel and 
wrought-iron, the knowledge that their difference is mainly in hardness, and 
that the so-called Bessemer steel is not tempered, have led to the introduc- 
tion of a new division and nomenclature for these substances (which are 
difficultly fusible and malleable compared with cast-iron). We now distin- 
guish : 
(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, form'erly Bessemer steel; (4) ingot steel, 
has been fused and tempered. 
Ordinary iron, even the purest wire, always contains foreign 
ingredients, principally carbon and manganese, and minute 
quantities of silicon, sulphur, phosphorus, nitrogen, nickel, 
cobalt, titanium and others. 'The quantity of manganese is 
purposely increased (to 30 per cent), as by this means the iron 
acquires valuable technical properties; it becomes more com- 
pact and solid. When iron, containing carbon, is dissolved 
in hydrochloric acid the chemically combined carbon unites 
with hydrogen,forming hydro-carbons, while the mechanically 
admixed graphite remains behind. The whole quantity of 
carbon is determined by the solution of the iron in bromine 
or cupric chloride, when all the carbon remains behind. 
To prepare chemically pure iron, heat the pure oxide or the 
oxalate in a current of hydrogen : 
the iron then remains as a fine black powder. If the reduc- 
tion 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 sul- 
phate 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 at i8oo°. Ordinary 
iron justs rapidly in moist air, as it covers itself with a thin 
layer of ferric hydrate. When ignited in the air it is coated 
with a layer of ferrous-ferric oxide (Fe304) which is readily 
detached. It burns with an intense light in oxygen. 
Fe203 + 3H2 = 2Fe + 3H20; 398 
INORGANIC CHEMISTRY. 
In contact with a magnet iron becomes magnetic; steel 
alone retains the magnetism, while cast-iron and wrought-iron 
soon lose the property after the removal of the magnet. 
Iron decomposes water at a red heat, with the formation of 
ferrous-ferric oxide, and the liberation of hydrogen : 
The metal dissolves without trouble in hydrochloric and 
sulphuric acids, with evolution of hydrogen ; the latter has a 
peculiar odor, due to hydrocarbons that are liberated at the 
same time. Iron dissolves in nitric acid with separation of 
nitric 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 probably due to the 
production of ferrous oxide upon its surface. 
3Fe + 4H20 = Fe364 + 4H2.. 
These are produced by the solution of iron in acids, and may 
also be obtained by the reduction of ferric salts : 
FERROUS COMPOUNDS. 
In the hydrous state they are usually of a green.color-; in 
the air they oxidize to ferric salts: 
Fe2Cl6 + Zn = 2FeCl2 +* ZnCl2. 
Ferrous Chloride — FeCl2— crystallizes from hydro- 
chloric acid solutions in green monoclinic prisms, with four 
molecules of water. These deliquesce in the air and oxi- 
dize. When dried they sustain a partial decomposition. 
The anhydrous salt is formed by conducting hydrogen chlo- 
ride over heated iron. It is a white mass, which fuses on 
application of heat and sublimes at a red heat in white, six- 
sided leaflets. 
It forms double salts with the alkaline chlorides, e.g.: 
2FeO -j- O = Fe203. 
FeCl2.2KCl + 2H20. 
Ferrous Iodide—Fel2—is obtained by warming iron 
with, iodine and water. It crystallizes with four molecules of 
water. 
Ferrous Oxide—FeO—is a black powder, resulting from 
the reduction of ferric oxide by carbon monoxide. When warmed 
it oxidizes readily. Ferrous Hydroxide—Fe(OH)2—is 
thrown out of ferrous solutions by the alkalies, as a greenish- 
white precipitate. Exposed to the air, it oxidizes, becoming IRON. 
399 
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 H20 in large, greenish, monoclinic prisms, and is generally 
called green vitriol. The crystals effloresce somewhat in dry air. 
They oxidize in moist air, and become coated with a brown 
layer of basic ferric sulphate. At ioo° they lose 6 molecules 
of H20, and change to a white powder. The last molecule of 
water escapes at 300°. Therefore, ferrous sulphate behaves just 
like the sulphates of the metals of the magnesium group. Like 
them, it unites with alkaline sulphates to double sulphates, which 
contain six molecules of water, e.g., S04Fe,S04K2 + 6H20. 
These are more stable than ferrous sulphate, and oxidize very 
slowly in the air. 
Ferrous sulphate is obtained by dissolving irqn in dilute 
sulphuric acid ; or from pyrites (FeS2). When the latter are 
roasted they lose one molecule. of sulphur, and are converted 
into ferrous sulphide (FeS), which, in the presence of water, 
absorbs oxygen from the air, and is converted into sulphate, 
which may then be extracted by water. 
Iron vitriol has an extended practical application ; among 
other uses, it is 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 sul- 
phuric acid (p. 192), and of colcothar. 
Ferrous Carbonate—FeC03—exists in nature as. siderite, 
crystallized in yellow-colored r-hombohedra, isomorphous with 
calcite and smithsonite. Sodium carbonate added to ferrous 
solutions precipitates a white voluminous carbonate, which 
rapidly oxidizes in the air to ferric hydrate. Ferrous car- 
bonate is somewhat soluble in water containing carbon diox- 
ide, hence present in many natural waters.' 
Ferrous Phosphate—Fe3(P04)2+ 8H20—occurs crystal- 
lized in bluish monoclinic prisms as Vivianite. Precipitated 
by sodium phosphate from ferrous solutions, it is a white 
amorphous powder, which oxidizes in the air. 
Ferrous Sulphide—FeS—is a dark-gray, metallic mass, 
obtained by fusing together iron and sulphur. It is made use 
of in laboratories for the preparation of hydrogen sulphide. 
If an intimate mixture of iron filings and sulphur be mois- 
tened with water, the union will occur even at ordinary tern- 400 
INORGANIC CHEMISTRY. 
peratures. Black ferrous sulphide is precipitated from ferrous 
solutions by alkaline sulphides. When the moist sulphide is 
exposed to the air it oxidizes to ferrous sulphate. The alka- 
line sulphides also precipitate ferrous sulphide from ferric 
salts, but the latter first suffer reduction: 
and 
Fe2Ci6 + (nh4)2S = 2FeCl2 + 2NH4C1 + S, 
FeCl2 -f (XH4)2S = FeS 4- 2NH4C1. 
FERRIC COMPOUNDS. 
Ferric Oxide—Sesquioxide of Iron—Fe203—exists in na- 
ture, in compact mass as hematite, and as iron mica, in dark 
gray metallic rhombic prisms. It may be prepared by heating 
the iron oxygen compounds in the air, and is obtained on a 
large scale by the ignition of green vitriol. It is then a dark- 
red powder {colcothar or caput mortuum) used as a paint and 
for polishing glass. 
Ferric Hydroxide—Fe2(OH)6—is precipitated by al- 
kalies from ferric solutions as a voluminous, reddish-brown 
mass. On boiling, it becomes more compact, gives up water 
and is converted into the hydrate, Fe20(OH)4. Many iron 
ores, like bog-iron ore, Fe20(0H)4, pyrosiderite, Fe202(OH)2 
(isomorphous with diaspore), and brown hematite Fe403(0H)6, 
are derived in a similar manner. 
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 hydroxide. All of the latter is 
precipitated as a jelly from such a solution upon the addition 
of a little alkali or acid. 
Ferrous-Ferric Oxide—Fe304 = FeO.Fe203—occurs in 
nature crystallized in black regular octahedra—magnetite. It 
is abundant in Sweden, Norway, and the Urals. It maybe ob- 
tained artificially by conducting steam over ignited iron (p. 
398). Magnetite constitutes the natural magnets. 
Ferric hydroxide, like other sesquioxides, is a feeble base, 
and does not yield salts with weak acids, like carbonic or 
sulphurous (p. 353). 
Ferric salts arise by the solution of ferric oxide in acids, 
or by the oxidation of ferrous salts in the. presence of free 
acids (best by chloric or nitric acids) : 
2FeS04 + H2S04 + O = Fe2(S04)3 + H20. IRON. 
401 
They generally have a yellow-brown color, and are converted 
by reduction into ferrous salts: 
Fe2Cl6 -f H2S = 2FeCl2 -f 2HCI -f S. 
Ferric Chloride—Fe2Cl6.—It is obtained in aqueous so 
lution by conducting chlorine into a solution of ferrous chlo 
ride : 
2FeCl2 + Cl, = Fe2Cl6. 
The hydrate—Fe2Cl6 + 6H20—remains upon evaporation. 
It is a yellow crystalline mass, readily soluble in water, alco- 
hol, and ether. It is partially decomposed when heated; hy- 
drogen chloride escapes, and a mixture of chloride and oxide 
remains. 
Anhydrous ferric chloride is produced by heating iron in a 
current of chlorine gas ; it sublimes in brownish-green, metal- 
lic, shining, six-sided prisms and scales, which deliquesce in 
the air. The specific gravity of their vapor is 161.5 (H = 1) 
corresponding to the molecular formula Fe2Cl6 = 323. 
Ferric Sulphate—Fe2(S04\—is obtained by dissolving 
the oxide in sulphuric acid. When its solution is evaporated, 
it remains as a white mass, which gradually dissolves in water, 
with a reddish brown color. It forms alums (p. 348) with 
alkaline sulphates, 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 in- 
soluble in water and acetic acid. 
Ferric Sulphide—FeS2—occurs in nature as pyrites, crys- 
tallized in yellow, metallic, shining, regular cubes or octahe- 
dra. It is employed in the manufacture of sulphuric acid and 
green vitriol. The artificial sulphide can be prepared in many 
ways. 
COMPOUNDS OF FERRIC ACID. 
On fusing iron filings with nitre, or by conducting chlorine 
into potassium hydrate, in which ferric hydroxide is suspen- 
ded, Potassium Ferrate, K2Fe04, is produced, and crystallizes 
from the alkaline solution in dark-red prisms. This salt is iso- 
morphous with potassium chromate and sulphate. It dissolves 
quite easily in water; but the dark-red liquid soon decom- 
poses with separation of ferric hydroxide and oxygen. The free 402 
INORGANIC CHEMISTRY. 
acid is not known, as it immediately breaks up when liberated 
from its salts. 
CYANOGEN DERIVATIVES OF IRON. 
Iron unites with the cyanogen group to form compounds 
which are very characteristic, and important in a commercial 
sense. When potassium cyanide is added to aqueous solutions 
of the ferrous or ferric salts, the cyan ides, Fe(CN)2and Fe2(CN)6, 
are thrown down as white precipitates, but decompose rapidly 
in the air. They dissolve in an excess of potassium cyanide to 
form thedoublecyanides,Fe(CN)2-4KCN and Fe2(CN)6.6KCN. 
When acids are added to these solutions the hydrogen com- 
pounds, H4FeCy6* (= FeCy2-4HCy) and Fe2Cy12H6(= Fe2Cy6. 
6HCy) separate. 
These are of acid nature, and form salts by exchanging 
their hydrogen for metals. The iron and the cyanogen group 
in these salts and the free acids cannot be detected by the 
usual reagents (e.g., iron is not precipitated by the alkalies). 
It is supposed that compound groups of peculiar structure are 
present in these double cyanides, and that they conduct them- 
selves like the halogens. The group, FeCv6, in the ous com- 
pounds is called ferrocyanogeti, that of Fe2Cy12 in the ic,ferri- 
cyanogcn. The ferro- behave toward the ferri-compounds the 
same as the ferrous toward the ferric salts; oxidizing agents 
convert the former into the latter, and reducing agents trans- 
form the latter into the former: 
2FeCy6K4 -f Cl2 = K6Fe2Cy12 -f 2KCI 
and 
Fe2Cyi2K« + 2KOH -j- H2 = 2K4FeCy6 + 2H20 
Cobalt, manganese, chromium and the platinum metals 
afford similar cyanides. 
Pjotassium Ferrocyanide—Yellow Prussiate of Potash— 
K4FeCy6, is produced by the action of potassium cyanide 
upon iron compounds, or upon free iron (in which case the 
oxygen of the air or of water takes part). It is prepared 
commercially by igniting carbonized nitrogenous animal mat- 
ter (blood, horns, hoofs, leather offal, etc.) with potashes and 
iron. 
In this operation, 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 sul- 
phide with the iron. (By means of alcohol, potassium cyanide 
* The cyanogen group, CN, is usually designated by the letters Cy. IRON. 
403 
can be extracted from the fusion.) Upon treating the fusion 
with water, the potassium cyanide and iron sulphide react upon 
each other, and ferrocyanide of potassium results and is purified 
by recrystallization : 
FeS + 6KCy = K4FeCy6 + K2S. 
It crystallizes from water in large, yellow, monoclinic 
prisms, having three molecules of water, and soluble in 3-4 
parts H20. The crystals lose all their water at ioo°, and are 
converted into a white powder. At a red heat the ferrocyanide 
breaks up into potassium cyanide, nitrogen, and iron carbide 
(FeC2). When the salt is warmed with dilute sulphuric acid, 
half of the cyanogen escapes as hydrogen cyanide; concen- 
trated sulphuric acid decomposes it, according to the following 
equation : 
K4FeCy6 + 6H2S04 + 6H20 = FeS04 + 2K2S04 + 3S04(NH4)2 
+ 6CO. 
When strong hydrochloric acid is added to a concentrated 
solution of potassium ferrocyanide hydrogen ferrocyanide, 
H4FeCy6, separates as a white crystalline powder, which soon 
turns blue in the air. It has the nature of an acid. Its salts 
with the alkali and alkali ne earth metals are very soluble in water. 
The sodium salt crystallizes with difficulty. The salts of the 
heavy metals are insoluble, and are obtained by double decom- 
position. When potassium ferrocyanide is added to the solu- 
tion of a ferric salt a dark blue cyanide (FeCy6)3(Fe2)2, called 
Prussian Blue is precipitated : 
This is the ferric salt of hydroferrocyanic acid ; and if po- 
tassium or sodium hydrate is poured over it, it is converted 
into ferrocyanide of potassium and ferric hydroxide: 
3K4FeCy6 + 2Fe2C16 = (FeCy6)3(Fe2)2 + 12KCI. 
VI VI 
(FeCy6’>3(Fe2)2 + 12KOH = 3K4FeCy6 + 2Fe2(OH)6. 
Potassium ferrocyanide produces a reddish-brown precipi- 
tate of FeCy6Cu2 in copper solutions. 
Oxidizing agents convert the ferro- into potassium ferri- 
cyanide—K6Fe2Cy12—redprussiateof potash. This conversion 
is most conveniently effected by conducting chlorine into the 
solution of the yellow prussiate : 
2K4l'eCy6 -f- CI2 — K6Fe2Cyj2 -j- 2KCI INORGANIC CHEMISTRV. 
The ferrocyanogen group, FeCy6, is then changed to the 
ferri-, Fe2Cy12. 
The red prussiate crystallizes from water in red rhombic 
prisms. The free hydro-ferricyanic acid, H6Fe2Cy12, is precip- 
itated upon the addition of concentrated hydrochloric acid. 
It is rather unstable. 
With ferrous solutions potassium ferricyanide affords a dark- 
blue precipitate, Fe3Fe2Cy12, very similar to Prussian Blue, and 
called Turnbull's Blue : 
This blue is the ferrous salt of hydroferricyanic acid. Al- 
kalies convert it into ferricyanide of potassium and ferrous 
hydroxide: 
K6Fe2Cyi2 + 3FeS04 = Fe2Cy12Fe3 + 3^2S04. 
Fe2Cy12Fe3 + 6KOH = Fe2Cy,2K6 -f- 3Fe(OH)2* 
Potassium ferricyanide does not cause precipitation in fer- 
ric solutions. Ferrocyanide yields Prussian blue, while it 
forms a bluish-white precipitate in ferrous solutions. 
By these reactions, ferric salts may be readily distinguished 
from the ferrous. Potassium sulphocyanide (CNSK) produces 
a dark-red coloration in ferric solutions, while it leaves the 
ferrous unaltered. 
3. COBALT. 
Co= 58.6. 
Occurs in nature as smaltite (CoAs2) and cobaltite (CoAs2. 
CoS2). The metal is obtained by the ignition of cobaltous 
oxide with carbon, or in a current of hydrogen. 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. It is only slightly attacked by hydrochloric 
and sulphuric acids; nitric acid dissolves it readily, forming 
cobaltous nitrate. 
The predominating compounds have the form CoX2, and 
are called cobaltous. They are very stable, and generally iso- 
morphous with the ferrous salts. The hydrous cobaltous com- 
pounds have a reddish color, the anhydrous are blue. 
* According to recent investigations it appears that Turnbull’s blue and 
VI 
( Fe 
Prussian blue possess the same composition (FeCy6)2 -J 2. The simpler 
relations are retained here. '■ COBALT. 
405 
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 anhydrous and blue in color. Characters made 
with this solution upon paper are almost invisible, but when 
warmed they become distinct and blue (sympathetic ink). 
Cobaltous Hydroxide—Co(OH)2—is a reddish precipi- 
tate produced by the alkalies in hot, cobaltous solutions. 
When exposed to the air, it browns by oxidation. Basic salts 
are precipitated from cold solutions. When heated out of air 
contact, the hydroxide passes into green cobaltous oxide, CoO. 
Cobaltous Sulphate—S04Co -f 7H20—crystallizes in 
dark red monoclinic prisms; the hydrated sulphate, CoS04 
-f 6H20, separates from hot solutions. It is isomorphous 
with ferrous sulphate, and yields double salts with alkaline 
sulphates. 
Cobaltous Nitrate—Co(N03)2 + 6H20—forms red deli- 
quescent prisms. 
Cobaltous Sulphide—CoS—is a black precipitate, pro- 
duced in neutral cobalt solutions by alkaline sulphides. It is 
insoluble in dilute acids. 
Cobalt Silicates.—When glass is fused with a cobalt 
compound it is colored a dark blue, and when reduced to a 
powder is used as a paint, under the name of smalt. 
Smalt is prepared commercially by fusing cobalt ores with potashes and 
quartz. The cobalt forms a silicate (smalt) with the Si02 and potassium, 
while the other metals accompanying it 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, Co203, with alumina, a dark-blue 
mass is produced—cobalt ultratnarine or Thenard's Blue. 
When zinc oxide and cobalt oxide are ignited a green color— 
green cinnabar or Rimnan's Green—is obtained. 
COBALTIC COMPOUNDS. 
Cobaltic Oxide—Co203—is left as a black powder on the 
ignition of cobaltous nitrate. It becomes cobaltous-cobaltic 
oxide, Co304, at a red heat, and cobaltous oxide at a white heat. 
The hydroxide—Co2(OH)6—separates as a dark-brown powder, 
if chlorine be passed through an alkaline solution containing 
a cobaltous salt. 
A cobaltous salt is produced, and hydrogen set free, when 406 
INORGANIC CHEMISTRY. 
sulphuric acid acts upon the oxide or the hydroxide. Chlorine 
is generated when it is heated with hydrochloric acid : 
Co2Os -f* 6HC1 = 2CoC12 + 3H20 -f- Cl,. 
The cobaltic hydroxide dissolves in dilute, cold hydrochloric 
acid, with scarcely any liberation of chlorine; the solution 
probably contains Co2Cl6, which decomposes into 2CoC12 and 
Cl2 on evaporation. 
Cobaltous-Cobaltic Oxide—Co304 = Co203CoO—cor- 
responding to magnetite, Fe304, is formed upon the ignition of 
the oxygen cobalt derivatives, and is a black powder. 
Only a few salts of cobalt in the ic state are known. The 
most interesting of these is potassio-cobaltic nitrite. 
When potassium nitrite, KN02, is added to a cobaltous solu- 
tion acidified with acetic acid, nitrogen is set free, and in 
course of time Co2(N02)6.6KN02 + nH20, the double salt, 
separates as a yellow crystalline powder. This reaction is 
very characteristic for cobalt, and serves to separate it from 
nickel. 
Ammonio-Cobalt Compounds.—Cobalt is capable of forming a 
series of peculiar compounds with ammonia, in which the metal appears in 
its highest state of oxidation; the structure of these derivatives has not 
yet been explained. On adding ammonium hydrate to a cobaltous chlo- 
ride solution, the precipHate first formed dissolves in the excess of the re- 
agent, and when this liquid is permitted to stand exposed to the air, the 
color, which is brown at first, gradually passes into red. On adding con- 
centrated hydrochloric acid to this solution, a brick-red, crystalline powder, 
of the composition — Co2Cl6 ioNH3 2H20—called roseocobaltic chloride 
—is precipitated. If, however, the red solution be boiled with hydrochloric 
acid, a red powder—purpureocobaltic chloride, Co2C16.ioNH3—separates 
out. If the ammoniacal red solution contain much ammonium chloride, 
hydrochloric acid will precipitate a yellowish-brown compound—luteo- 
cobaltic chloride, Co2Cl6 I2NH3. 
The other salts of cobalt, such as the sulphate and nitrate, yield similar 
compounds, e.g., Co2{ N03'6. ioNH3, roseocobaltic nitrate. 
Cyanogen Cobalt Compounds.—In solutions of cobaltous salts, po- 
tassium cyargde produces a bright brown precipitate of cobalto-cyanide 
Co(CN)2, soluble in an excess of the reagent The solution absorbs 
oxygen from the air, and is converted into potassium cobalticyanidc, 
KBCo2(CN)i2, corresponding to potassium ferricyanide. When the solu- 
tion is evaporated the cobalticyanide crystallizes in colorless rhombic 
prisms, very soluble in water. Sulphuric acid precipitates hydrogen co- 
balticyanide, H6Co2(CN)i2, from the concentrated solution. This acid 
crystallizes in needles. NICKEL. 
407 
3. NICKEL. 
Nickel exists in native condition in meteorites; its most 
important ores are Niccolite—NiAs—and Gersdorffite, NiS2. 
NiAs, (constituted like cobaltite). It is always accompanied 
in its ores by cobalt, and vice versa, cobalt usually by nickel. 
The isolation of the latter from its ores and from speiss cobalt 
(p. 405) is very complicated. Nickel usually appears in com- 
merce in cubical forms, which in addition to the chief ingre- 
dient always contain some copper, bismuth, and other metals. 
Chemically pure nickel is procured by igniting the oxalate or 
carbonate in a current of hydrogen. Nickel is almost silver- 
white in color and is very lustrous, and very tenacious. 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 it is 
attracted by the magnet. It is not altered in the air; it dis- 
solves with difficulty in hydrochloric and sulphuric acids, but 
readily in nitric acid. 
Its derivatives are almost exclusively of the ous form NiX2; 
nickelic oxide behaves like a peroxide, and does not afford 
corresponding salts. 
Nickelous Hydroxide—Ni(OH)2—is a bright-green pre- 
cipitate produced by alkalies in nickelous solutions. It dis- 
solves in ammonium hydrate, with a blue color. When heated 
it passes into gray nickelous oxide, NiO. 
Nickelous Chloride—NiCl2 + 6H20—consists of green, 
monoclinic prisms. When heated they lose water and become 
yellow. 
Nickelous Cyanide—Ni(CN)2—is precipitated by potas- 
sium cyanide as a green colored mass from nickel solutions. It is 
soluble in excess of the precipitant. The double cyanide, NiCy2 
2KCy -f H20, crystallizes from the solution. This salt is read- 
ily decomposed by acids. Cyanogen compounds of nickel, con- 
stituted like those of iron and cobalt, are not known. 
Nickelous Sulphate—NiS04+7H20—appears in green, 
rhombic prisms, isomorphous with the sulphates of the magne- 
sium group, and forms analogous double salts. 
Nickelous Sulphide—NiS—is precipitated, black in 
color, by alkaline sulphides from nickel solutions. 
Nickelic Oxide—Ni2Os—and Hydroxide—Ni2(OH)6— 
are perfectly similar to the corresponding cobalt salts ; when 
warmed with hydrochloric acid they liberate chlorine. 
Nickel is used for certain alloys. Argentan consists, ordin- 
arily, of 50 per cent, copper, 25 per cent..nickel and 25 per 
Ni = 58.6. 408 
INORGANIC CHEMISTRY. 
cent. zinc. The German nickel coins consist of 75 per cent. 
Cu and 25 per cent. Ni. The alloy will be whiter in color 
and harder, and also receive a higher polish in proportion to 
the amount of nickel that it contains. At present, cast-iron 
ware is coated with a layer of nickel to prevent it from rusting 
and to impart to it a beautiful white surface. This is accom- 
plished in an electrolytic manner, or by boiling the iron ware 
in a solution of zinc chloride and nickel sulphate. 
In the electrolytic method the solution of the double sul- 
phate of nickel and ammonium is employed ; the positive 
electrode consists of a pure nickel plate, while the object to be 
coated is attached to the negative electrode. 
GROUP OF THE PLATINUM METALS. 
Besides platinum, this group comprises palladium, rhodium, 
ruthenium, osmium and iridium—the constant companions of 
the first in its ores. On page 393 we observed that these 
metals are divided into two groups ; the group of light platinum 
metals, and the group of heavy platinum metals which have 
higher atomic weights and specific gravities: 
Ru, 103. Rh, 104. Pd, 106.2 Os, 195* Ir, 192.5 Pt, 194.3 
Sp. gr. “ 12.26 “ 12.1 “ 11.8 “ 22.4 “ 22.38 “ 21.4 
The relations of the metals of this group to each other are 
perfectly similar to those of the iron group ; and they show in 
their physical and chemical properties a great resemblance to 
the corresponding members of the iron group. Osmium and 
ruthenium, like iron, have a gray color, are very difficultly 
fusible and readily oxidized in the air. Palladium and 
platinum, on the other hand, have an almost silver-white color 
like nickel, are more fusible, and are not oxidized by oxygen. 
In chemical respects osmium and ruthenium, like iron, also 
show a metalloidal nature, inasmuch as their highest oxygen 
compounds form acids. Their derivatives show a complete 
parallelism with those of iron. 
II 
Ill 
IV 
VI 
OsO 
Os203 
0s02 
(OsOs) 
Osmous 
Osmic 
Osmium 
Osmic 
oxide. 
oxide. 
dioxide. 
trioxide. 
RuO 
Ru20, 
Ru02 
(Ru03) 
Ruthenous 
Ruthenic 
Ruthenium 
Ruthenium 
oxide. 
oxide. 
dioxide. 
trioxide. 
* See note (page 393). PLATINUM. 
409 
The acid oxides 0s03 and Ru03 are unknown, but the cor- 
responding acids,- H20s04 (osmic acid) and H2Ru04 (ruthenic 
acid), and their salts have been obtained. Besides the deriva- 
tives already mentioned we find that osmium and ruthenium 
are capable of still higher oxidation, yielding 0s04, per-osmic 
oxide, and Ru04, per-ruthenic oxide—which is not the case 
with iron; in these compounds the metals appear-to be octads, 
yet these oxides do not afford corresponding acids or salts. 
Rhodium and iridium, like cobalt, do not yield acid-like 
derivatives. Their salts correspond to the forms: 
II 
RhO 
Ill 
Rh2o3 
IV 
Rh02. 
Rhodous 
oxide. 
Rhodic 
oxide. 
Rhodium 
dioxide. 
The rhodic compounds are the more stable. 
Palladium and platinum, finally, are relatively of more 
basic nature, as their ous derivatives, PdX2 and PtX2, are pro- 
portionally 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 found in nature almost exclusively 
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 especially in 
the Urals). The platinum ore, like that of gold, is obtained 
by the elutriation of the platiniferous sand with water whereby 
the lighter particles are carried away. Platinum ore usually 
contains 50-80 per cent platinum, besides palladium (to 2 per 
cent.), iridium (to 7 per cent.), osmium (1 per cent.), and 
ruthenium per cent.), and different other metals, as gold, 
copper, and 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, when platinum, palladium, rhodium, ruthenium, and 
a portion of iridium are dissolved. Metallic grains or leaflets, 
an alloy of osmium and iridium—platinum residues—remain. 
Ammonium chloride is then added to the solution and plati- 
num and iridium precipitated as double salts. When the pre- 
cipitate is ignited a spongy mass of iridium-bearing plati- 
num (platinum sponge) is obtained, which is applied directly 410 
INORGANIC CHEMISTRY. 
in the manufacture of platinum vessels. The filtered 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. 
Formerly spongy platinum was employed almost exclusively 
for the manufacture of platinum objects; it was pressed into 
moulds, then ignited and hammered out. Now the fusibility 
of Pt in the oxy-hydrogen flame is resorted to, and the fused 
metal run into moulds. 
Platinum containing both iridium and rhodium may be 
fused directly out of the platinum-ore by means of the oxy- 
hydrogen blowpipe. The greater portion of the osmium and 
ruthenium is consumed in this operation. The presence of 
iridium and rhodium makes platinum harder and less readily 
attacked by many reagents. 
RUTHENIUM AND OSMIUM. 
Ru = 103.6. 
Os = 195.* 
Ruthenium has a steel-gray color; it is very hard, brittle, and difficultly 
fusible (about 1800°). When pulverized and ignited in the air it oxidizes 
to RuO and Ru2Og. It is insoluble in acids, and only slowly dissolved by 
aqua regia. When fused with potassium hydroxide and nitrate, it forms 
potassium rutheniate, K2Ru04. 
Ruthenium heated in chlorine gas yields ruthenium dichloride, RuC12, a 
black powder, insoluble in acids. The sesquichloride, Ru2C16, is obtained 
by the solution of Ru2 OH)g in hydrochloric acid, and is a yellow, crystal- 
line mass, which deliquesces in the air. It yields crystalline double chlo- 
rides with potassium and ammonium chlorides, eg., Ru2Clg.'KCl. The 
tetrachloride, RuC14, is only known in double salts. Ruthenious oxide, 
RuO, the sesquioxide, Ru208, and dioxide, Ru02, are black powders, in- 
soluble in acids, and are obtained when ruthenium is roasted in the air. 
The hydroxides, Ru2(OH)6 and Ru(OH)4, are produced by the action 
of the alkalies upon the corresponding chlorides, and are very readily 
soluble in acids. Ruthenic acid, II2Ru04, is not known in a free condi- 
tion. Its potassium salt, K2Ru04, is formed by fusing the metal with po- 
tassium hydroxide and nitre. It dissolves in water with an orange-yellow 
color. When chlorine is conducted through the solution ruthenium tetrox- 
ide, Ru04. separates as a gold-yellow crystalline mass. It fuses at 40°, 
boils about ioo°, and yields a yellow vapor, the odor of which is similar 
to nitrogen dioxide, N()2. It decomposes with explosion at 108°. Water 
breaks it up with formation of Ru2(OHl6. It dissolves to Ru04K2 in con- 
centrated potassium hydrate. When less chlorine is introduced into the 
solution of RuC>4K2, greenish-black crystals separate out, which are isomor- 
phous with potassium permanganate, and appear to be Ru04K. 
* Compare note (p. 393). The atomic weight of osmium, ascertained 
from the vapor density of the tetroxide, 0s04, is 193. RHODIUM AND IRIDIUM—PALLADIUM. 
411 
Osmium is very much like the preceding. It is not even fusible in the 
oxyhydrogen flame ; it only sinters together. According to Violle it fuses 
at 2500°. Reduced to a fine powder it will burn when ignited in the air 
to 0s04. Nitric acid and aqua regia convert it into the same oxide. The 
compounds, OsCl2 and OsO, Os2Cl6 and 0s203,0s02 and OsCl4, are very simi- 
lar to the corresponding ruthenium derivatives. By fusion with potassium 
hydroxide and nitre we get potassium osmate—K20s04—which crystallizes 
from aqueous solution with 2H20 in dark-violet octahedra. The most 
staTle and a very characteristic derivative of osmium is the tetroxide, 0s04, 
which is produced by igniting the metal in the air, or by the action of 
chlorine on osmium in the presence of water. It crystallizes in large col- 
orless prisms, which fuse below ioo° and distil at a somewhat higher tem- 
perature. It has a very sharp, piercing odor, similar to that of sulphur 
chloride. Reducing and organic substances precipitate pulverulent osmium 
from it. This is the basis of its application in microscopy. 
Os04 and Ru04 do not afford corresponding salts. 
RHODIUM AND IRIDIUM. 
These metals are lighter in color and are more easily fusible than ruthe- 
nium and osmium. (Iridium fuses at 19500.) When pure they are not 
attacked by acids or aqua regia; but when alloyed with platinum they dis- 
solve in the latter. 
Rhodium forms three oxides: RhO, Rh203 and Rh02, of which the 
second forms salts with acids. Rh02 results when rhodium is heated with 
nitre. 
Of the chlorides only Rh2Cl6 is known. It results when the metal is 
heated in chlorine gas. It is a brownish-red mass. It forms readily 
crystallizing, re'd-colored double salts with alkaline chlorides. 
Iridium has perfectly analogous derivatives : IrO, lr203, Ir02 and IrCl2, 
IraCl6, IrCl4. The sesquichloride, Ir2Cl6, formed by heating Ir in chlo- 
rine, is an olive-green, crystalline mass, insoluble in water and acids. It 
affords double salts with the alkaline chlorides, e.g., Ir2Cl6, 6KC1 -j- 6II20, 
which crystallizes from water in green crystals. They are also produced 
by the action of S02 upon the double salts of IrCl4. 
Iridium Tetrachloride—IrCl4—is produced in the solution of iridium 
or its oxide in aqua regia, and remains, on evaporation, as a black mass, 
readily soluble in water (with red color). When alkaline chlorides are 
added to the solution double chlorides are precipitated, eg., IrCl4.2NH4Cl, 
isomorphous with the double chlorides of platinum. When a solution of 
IrCl4 is boiled with KOII, IrvOH)4 will be precipitated. 
Rh = 104. 
Ir = 192.5. 
PALLADIUM. 
Pd == 106.2. 
Palladium, in addition to occurring in platinum ores, is 
found alloyed with gold (Brazil), and in some selenium ores 
(Hartz); it has a silver-white color, and is somewhat more 412 
INORGANIC CHEMISTRY. 
fusible (about 1500°) than platinum. When finely divided it 
dissolves in boiling concentrated hydrochloric, sulphuric, and 
nitric acids. When ignited in the air it at first becomes dull 
by oxidation, but at a higher temperature the surface again 
assumes a metallic appearance. 
Palladium absorbs hydrogen gas (occlusion) to a much 
greater extent than platinum or silver. Freshly ignited pal- 
ladium leaf absorbs upwards of 370 volumes of hydrogen at 
ordinary temperatures, and about 650 volumes at 90-100° C. 
A greater absorption may be effected at ordinary temperatures 
in the following manner: 
Water is decomposed by the electric current, palladium foil 
being used as negative electrode. The liberated hydrogen is 
then taken up by the palladium (to 960 volumes); the metal 
expands its volume), becomes specifically lighter, but re- 
tains its metallic appearance entire. According to the inves- 
tigations of Debray, the compound Pd.2H is produced, which 
contains dissolved hydrogen, and deports itself similarly to an 
alloy (compare p. 45). Palladium charged with hydrogen 
usually remains unaltered in the air, and in a vacuum ; it, how- 
ever, sometimes becomes heated in the air, as the hydrogen is 
oxidized to water. The same occurs when palladium hydride is 
heated to ioo°; in vacuo, all the hydrogen escapes as gas. 
Palladium hydride is a strong reducing agent, like nascent 
hydrogen. Ferric salts are reduced to the ferrous state ; chlo- 
rine and iodine in aqueous solution are converted into hydro- 
chloric and hydriodic acids. 
Palladium black absorbs hydrogen more energetically than 
the compact variety (at ioo° upwards of 980 volumes). This 
substance is obtained by the reduction or electrolysis of pal- 
ladic chloride. If palladium sponge be heated in the air until 
the white metallic color becomes black, in consequence of the 
superficial oxidation, it will absorb hydrogen very energetically 
at ordinary temperatures, and partially oxidize it to water. 
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. 
There are two series of palladium compounds: the palla- 
dious, PdX2, and palladic, PdX4. The first are well character- 
ized and are distinguished by their stability. 
Palladious Chloride—PdCl2—remains as a brown, deli- 
quescent mass, on evaporating the solution of palladium in PLATINUM. 
413 
aqua regia. It yields easily soluble crystalline double salts, 
with alkaline chlorides, e.g, PdCl2,2KCl. 
Palladious Iodide—Pdl2—is precipitated from palla- 
dium solutions by potassium iodide as a black mass, insoluble 
in water. 
Palladious Oxide—PdO—is a black residue left upon 
careful ignition of the nitrate. It is difficultly soluble in acids. 
When heated, it loses oxygen, and forms palladium suboxide, 
Pd20. 
When palladium is dissolved in sulphuric or nitric acids, the 
corresponding salts are produced. 
The sulphate, PdS04 + 2H20, is composed of brown crys- 
tals, readily soluble in water. Much of the latter decomposes 
it. 
Palladic Chloride—PdCl4—is formed when the metal is 
dissolved in aqua regia. It decomposes, on evaporation, 
into PdCl2 and Cl2. When potassium or ammonium chloride 
is added to its solution, red-colored, difficultly soluble double 
chlorides crystallize out; they are analogous to the corres- 
ponding salts of platinum. 
PLATINUM. 
Pt= 194-3-* 
The separation of platinum from the ore was described on 
page 409. The metal has a grayish-white color, and a spe- 
cific gravity of 21.4. It is very tough and malleable, and may be 
drawn out into very fine wire and rolled into foil. It becomes 
soft without melting at an intense heat. It fuses in the oxy-hy- 
drogen flame (about 1770°—Violle), and is somewhat volatile. 
On fusion, it absorbs oxygen, which it gives up again on cooling 
(like silver). At ordinary temperatures, it also condenses oxy- 
gen upon its surface, especially when in a finely divided state, 
as platinum black or sponge. The first is obtained, if reducing 
substances, like zinc, be added to solutions of platinic chloride 
or upon boiling the solution with sugar and sodium carbonate ; 
it absorbs as much as 250 volumes of oxygen. Platinum sponge 
is obtained by the ignition of PtCl42NH4Cl. The production 
of various reactions is due to this power of platinum to condense 
oxygen; thus hydrogen will inflame in the air, if it be con- 
* Until recently, the atomic weight of platinum was considered 196.7. 
According to the periodic system, it ought to be less than that of gold 
(196.2). The latest investigations in this direction confirm this view. 414 
INORGANIC CHEMISTRV. 
ducted upon platinum sponge; sulphur dioxide combines with 
O at ioo° to form the trioxide. At a red heat platinum per- 
mits free passage to hydrogen, while it is not permeable by 
oxygen and other gases (pp. 93 and 260). 
Platinum is not attacked by acids; it is only soluble in 
liquids generating free chlorine, e.g., aqua regia. In conse- 
quence of this resistance to acids, and 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 hydroxides, sulphides, and cyanides attack it 
strongly at a red heat. It forms readily fusible alloys with 
phosphorus, arsenic, and many heavy metals, especially lead, 
it also reduces many heavy metals from their platinum salts. 
Therefore such substances must not be ignited in platinum 
crucibles, etc. 
Platinum, like palladium, affordspiatinous, PtX2,andplatinic, 
PtX4, derivatives ; in the first it is more basic, in the latter 
more acidic. 
Platinic Chloride—PtCl4—is obtained by the solution of 
platinum in aqua regia, and when the solution is evaporated, 
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 imme- 
diately separate out as a crystalline yellow powder. Ignition 
completely decomposes the ammonium salt, leaving spongy 
platinum. 
Platinum chloride affords similar insoluble double chlorides 
with those of rubidium, caesium, 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- 
cipitate of platinic hydroxide, 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 hydrox- 
ide has also an acidic character {platinic acid), and dissolves 
in alkalies, yielding salts with them. These, also, result on 
fusing platinum with potassium and sodium hydroxide. The 
barium salt, Pt j Q^qT2 + 3H20, is precipitated from platinic 
chloride, by barium hydrate, as a yellow, crystalline com- PLATINUM. 
415 
pound. The acidic nature of its hydrate allies platinum to 
gold. If hvdrogen sulphide, be conducted through platinic 
solutions, black platinum disulphide, PtS2, is precipitated'; it 
is soluble in alkaline sulphides, with formation of sulpho-salts. 
Platinous Chloride—PtCl2—is a green powder, insoluble 
in water, remaining after heating PtCl4 to 200°. It affords 
double salts with alkaline chlorides, e.g., PtCl2,2NaCl. When 
digested with potassium hydroxide it yields the hydroxide, 
Pt(OH)2. . 
Cyanogen Compounds.—Like iron and cobalt, plati- 
num affords double cyanides corresponding to the ferrocyan- 
ides. When platinous chloride is dissolved in potassium 
cyanide platinum-potassium cyanide, K2PtCy4 + 3H20, crystal- 
lizes on evaporation in large prisms exhibiting magnificent 
dichroism; in transmitted light they are yellow and in reflected 
light blue. This salt must be viewed as the potassium com- 
pound of hydro-platino-cyanic acid, H2PtCy4. When separated 
from its salts it crystallizes in gold-yellow needles. Its salts 
with heavy metals are obtained by double decomposition, 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 according 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 platinous 
chloride. The bases are obtained by substituting hydroxyl 
groups for the acid radicals, e.g., Pt(NH3)4(OH)2. They 
resemble alkaline hydroxides in their chemical properties. 
The other platinum metals afford similar amine derivatives. 
The nature and chemical constitution of these interesting 
compounds is, however, not fully explained. 416 
INORGANIC CHEMISTRY. 
SPECTRUM ANALYSIS.* 
We observed that various substances, when introduced into 
a non-luminous flame, imparted to it a characteristic colora- 
tion. Thus, the sodium compounds color it yellow; the po- 
tassium, violet; thallium, green ; etc. The decomposition of 
the light thus obtained, and, indeed, of every light, by means 
of a prism and the study of the resulting spectrum form the 
basis of spectrum analysis, established in 1859 by Kirchhofif 
and Bunsen. Its important applications and universal use 
constitute one of the gieatest of scientific achievements of all 
ages. 
As we well know, every substance, solid or liquid, heated 
to white heat (e.g., molten platinum ; lime heated in the oxy- 
hydrogen flame; the ordinary flame containing glowing par- 
ticles of carbon ;) emits rays of every refrangibility; and hence, 
furnishes a continuous spectrum, which brings to view all the 
colors of the rainbow, from red to violet, if the light be con- 
ducted through a prism. Glowing gases and vapors, on the 
contrary, whose molecules can execute unobstructed oscilla- 
tions, emit light of definite refrangibility, and, therefore, af- 
ford spectra, consisting of single, bright lines. Thus, the 
spectrum of the yellow sodium flame is recognized as composed 
of one very bright yellow line, which by increased magnify- 
ing power is shown to consist of two lines lying very near each 
other. This reaction is so very delicate, that of a 
milligram of sodium may be detected by it. The violet po- 
tassium light affords a spectrum, consisting of a red and a blue 
line. The crimson strontium light shows in the spectrum 
several distinct red lines and a blue line. (See the spectrum 
plate.) Each of these lines corresponds to a definite coefficient 
of refraction, and therefore occupies a definite relative position 
in the spectrum. 
If substances affording different colors be introduced into a 
flame, the most intense color generally obscures the others ; in 
the spectrum, however, each individual substance shows its 
peculiar bright lines, which appear simultaneously or succeed 
each other, according to the volatility of the various sub- 
stances. 
The spectrum apparatus or spectroscope, figured in Fig. 89, 
serves for the observation of the spectra. 
* A more exhaustive, concise, and distinct presentation of the spectrum 
phenomena maybe found in Herman W. Vogel’s “ Practische Spec ral- 
analyse irdischer Stoffe,” 1882. SPECTRUM ANALYSIS. 
417 
In the middle of the apparatus is a flint-glass prism P. At 
the further end of the tube A is a movable vertical slit, in 
front of which is placed the light to be investigated. The en- 
tering light rays are directed by a collecting lens into the 
tube A, upon the prism, and the refracted rays (the spectrum) 
are observed by the telescope B. The tube C is employed to 
ascertain the relative position of the spectrum lines. This is 
provided at its outer end with a transparent horizontal scale. 
Fir. 89. 
When a luminous flame is placed before the scale its divisions 
are reflected from the prism surface and thrown into the tele- 
scope B. We then see the spectrum to be studied and the 
scale divisions in B at the same time, and can readily deter- 
mine the relative position of the lines of the spectrum. To 
study two spectra at the same time, and compare them, a three- 
sided, right-angled glass prism is attached in front of one-half 
(the lower or upper) of the slit of the tube A; this directs 
the rays of a light placed at the side (/, Fig. 89) through A 
upon the prism P. By means of B, two horizontal spectra 
will be observed, one above the other, and between are the 
bright divisions of the scale. 
Adjustment of the Spectroscope.—To observe the spectra in the appa- 
ratus described, it is necessary to first adjust the same correctly. The tube 
A contains, besides the slit, also a lens (Collimator Lens), which serves to 418 
INORGANIC CHEMISTRY. 
render parallel the bunch of rays proceeding from the slit; hence, the latter 
must be accurately placed in the focus of the lens. This is best accom- 
plished as follows: The telescope (JS) is drawn out and adjusted for some 
distant object, that it may be adapted for the reception of parallel rays; it 
is then replaced in the stand, pointed toward the slit, illuminated by a 
sodium chloride flame, and the slit then removed far enough to appear 
perfectly distinct in the telescope. To have the spectrum lines sharply 
defined, the slit must be made quite narrow; for feebly luminous lines it 
must, however, be widened. The horizontal black lines that appear in the 
spectrum arise from dust particles adhering to the slit. 
The proper position of the tube with the slit with respect to the prism is 
usually fixed by the frame on which they rest. It must be so adjusted that 
the refracted rays pass through the prism as symmetrically as possible, i.e., in 
the minimum of their deviation, otherwise the spectrum will be less distinct 
and (owing to unequal refraction) will appear distorted. The symmetrical 
passage of the spectrum rays is approximately attained when the intermed- 
iate green rays pass through the prism symmetrically. Such a position 
must, therefore, be given the prism, with reference to the tube carrying the 
slit, that the middle green rays (line E of the sun spectrum) pass through 
it in the minimum of their deviation. It is then 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). This arrangement is such 
(according to Bunsen) that the yellow sodium line coincides with the line 
50 of the scale; then the red potassium line (<z) will lie at 17, and the 
violet (/?) at 152 (apparatus of Desaga). Since the refraction and disper- 
sion of the rays are influenced by the quality of the glass of the prism, the 
scale indications of different forms of apparatus are not directly com- 
parable. They must, therefore, be referred to an absolute measure. This 
is most conveniently attained by reduction to the sun spectrum, which may 
be rendered visible in the telescope at the same time by means of a com- 
parison-prism attached to the tube carrying the slit. We then determine 
the dark lines of the sun with which the lines of the flame under investi- 
gation coincide. In accurate determinations, the spectrum lines are repre- 
sented in wave lengths, according to the millionth of a millimeter. 
The apparatus described above is that usually employed in chemical 
laboratories. There are other forms adapted for special purposes—for 
investigation under the microscope, for the observation of the sun and 
stars. In accurate observations, where clear, broad spectra are required, 
the light is permitted to pass through several (3-9) prisms, consisting of 
hollow glass filled with carbon disulphide, which refracts light very 
strongly. Diffraction spectra may be advantageously applied in many 
instances in the place of the prism spectra. 
The direct line (d vision directe) spectroscopes are very excellent for 
laboratory purposes. With these the spectra may be viewed in the direc- 
tion in which the luminous objects really are; there are no deflections. 
This is accomplished by a combination of several prisms of crown and 
flint glass, whereby dispersion is attained, together with the simultaneous 
removal of deflection. 
To observe the spectra of metals, it is only necessary, in 
many instances, e.g.f with the alkalies and alkaline earths, to SPECTRUM ANALYSIS. 
419 
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 first 
into HC1 and NaOH, which is then reduced by the carbon of 
the glowing gases to metallic sodium, and colors the flame 
yellow. The compounds difficult to decompose (as the barium 
salts), often afford independent spectra, differing from those 
of the free metal; this is plainly recognizable in the copper 
compounds. x 
Most metals, however, require a much higher temperature 
than that of the gas flame for their conversion into gases. To 
vaporize them and observe their spectra, the electric spark is 
made to pass from electrodes constructed of them. In this 
manner we can study all metals, even the most non-volatile, 
like gold, iron, and platinum. Their spectra are generally 
quite complicated, and exhibit a great number of single lines. 
Thus, over 450 lines have been established for iron. 
Instead of making the electrodes of the metals we wish to study, we 
can, according to Bunsen (Poggend, Ann., 155), employ Carbon points 
saturated with solutions of the metallic salts. A Ruhmkorf induction 
apparatus, with a dip battery of 4 elements, and an attached Leyden jar, 
will suffice to produce the electric spark. Such spark spectra are often 
quite different from the flame spectra obtained in the gas flame. 
A peculiar, very interesting and practical procedure for the production 
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 salt that is being examined. 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 wire connected with the positive 
pole of the induction spiral. The spectra of all the metals may be easily 
obtained by this contrivance, and indeed, Lecoq de Boisbaudran discov- 
ered gallium by it. 
The spectra of the elementary gases may be obtained by 
passing electric sparks through them ; these will be variously 
colored. Hydrogen shines with a red light, which is decom- 
posed and gives a spectrum consisting of a bright red, a blue, 
and a green line. Nitrogen shines with a violet light and 
affords a spectrum of many lines, chief of which are the 
violet. The spectra of gas may be more conveniently ob- 
served by aid of Geissler’s tubes, which are filled with very 
dilute gases and the induction stream then passed through 
them. 
These methods give us a means of readily distinguishing the 
individual chemical elements, and even detecting them in 420 
INORGANIC CHEMISTRY. 
traces. Since the year i860 various new elements, e.g., caesium, 
rubidium, thallium, indium, scandium, gallium, and several 
others, not accurately studied, have been discovered by their 
aid. 
In addition to the direct, bright spectra just described there 
are yet dark absorption spectra. If a white light giving an un- 
interrupted spectrum be allowed to pass through different 
transparent bodies, the latter will absorb rays of definite re- 
frangibility, allowing all others to pass. Therefore we observe 
the sun spectrum interrupted by dark lines or bands in the 
spectroscope. The solutions of didymium and erbium absorb 
certain rays, and show corresponding dark lines in the spec- 
trum. The gases deport themselves similarly. White light 
that has traversed a broad layer of air shows in the spectrum 
several dark lines peculiar to nitrogen, oxygen, and steam. 
This power of absorption is possessed to a much higher degree 
by 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), a dark line 
will appear in the rainbow spectrum of the white light, and its 
position will correspond exactly to that of the yellow sodium 
line; the latter thus appears converted into a dark line. If 
white light be passed through the potassium flame, two dark 
lines will be visible in the spectrum, corresponding to the red 
and blue lines of the potassium spectrum. Such spectra are 
designated the inverted spectra of the corresponding metals. 
The inverted spectra of all elements may be obtained in 
this way, and they correspond accurately to the direct 
bright spectra. The cause of these phenomena lies in the 
proposition deduced by Kirchhoff from the undulatory theory 
of light, that the ratio between the emissive and absorptive power 
is the same for almost all bodies at like temperatures. Accord- 
ing to this, incandescent gases only absorb rays of just the 
same refrangibility as those which they emit. For example, 
when bright white light is passed through the yellow sodium 
flame the yellow rays of the former are absorbed and retained, 
while all others pass on almost entirely unaltered. Therefore, in 
the rainbow spectrum of white light the yellow rays of definite 
refrangibility will be absent; and if the other refracted rays of 
the white light are brighter than the yellow rays emitted from 
the sodium flame, the latter will be relatively darker; a dark 
line will therefore make its appearance. 
These phenomena have presented a new and wide province 
to spectral analysis, inasmuch as they open up avenues for SPECTRUM ANALYSIS. 
421 
the investigation of the chemical nature of the sun and other 
bodies. 
It is known that the bright rainbow sun spectrum is inter- 
sected by a number of dark lines which have been called the 
Frauenhofer lines from their discoverer. Kirchhoff has shown 
that these lines can be easily accounted for, after what has 
already been said, by the following hypothesis upon the nature 
of the sun : The latter consists of a solid or liquid luminous 
nucleus, surrounded by an atmosphere of incandescent gases 
and vapors. Then the uninterrupted spectrum of the glowing 
nucleus must be intersected by the dark lines of the inverted 
spectra of those gases and vapors which occur in the sun’s 
atmosphere. An accurate comparison of the Frauenhofer 
lines with the spectrum lines of the various elements, has re- 
vealed the fact that iron, sodium, magnesium, calcium, chro- 
mium, nickel, barium, copper, zinc, and hydrogen are present 
in the sun’s atmosphere. Thus, dark lines have been found 
in the sun’s spectrum corresponding to all of the 450 lines of 
the iron spectrum. The inferences upon the chemical consti- 
tution of the sun possess as much and, indeed, a higher de- 
gree of probability than falls to many other deductions. 
The investigation of the sun’s spectrum has cleared up many 
other changes occurring there, and this has led to a complete 
sun-meteorology. All the fixed stars thus far investigated possess 
a constitution like that of the sun. They give spectra inter- 
sected by dark lines, and therefore consist of incandescent 
nuclei surrounded by gaseous atmospheres. The spectra of 
nebulae, however, only show bright lines; hence these consist 
of uncondensed, incandescent masses of vapor. 
Periodicity of the Spectrum Lines.—Since all other properties of 
the elements and their compounds have proved therrtselves to be periodic 
functions of the atomic weights, we may expect the same with reference to 
the spectrum phenomena. But few such regularities have occurred thus 
far. Most of the elements, especially the metalloids and the difficultly 
volatile metals, afford very complex spectra, which often vary at the same 
time with the temperature, so that spectra of ist, 2d, etc., order are distin- 
guished for some of them. It appears that not all spectral lines are of the 
same importance, since it has been possible, in some instances, to refer 
the various lines of a spectrum to particular fundamental lines whose rela- 
tions to each other are comparable with those of the harmonics to primary 
tones. Thus the four lines of the hydrogen spectrum may be looked upon 
as the harmonics of a single wave. Therefore, only individual lines are 
to be regarded in the comparison of spectra. This is clearly observed 
with the easily volatile metals belonging to the homologous groups, K, Rb 422 
INORGANIC CHEMISTRY. 
Cs; Ca, Sr, Ba; Ga, In, T1; whose lines, lying in the violet portion of the 
spectrum, advance more towards the red end as the atomic weights increase ; 
the wave lengths (in millionths of a millimeter) become successively 
greater with the increase of the latter or that of the atomic volumes: 
K 39 Wave length, 404 
Rb 83 “ “ 420.421 
CSI32 “ “ 456-459 
Ca 40 Wave length, 422 
Sr 87 “ “ 461 
Ba 137 “ “ 525-550 
Ga 69 Wave length, 403.417 
In 113 “ “ 410.450 
T1209 “ “ 535. 
These relations are made apparent by the arrangement of the spectra in 
the spectrum plate. 
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 it 
seems possible to draw a conclusion upon the spectra of the succeeding 
elements. Indeed, the element scandium (45) succeeding calcium, shows 
intense violet lines of the wave length 425-440. HEAT OF FORMATION OF THE METALS. 
423 
a 
b 
a 
b 
a 
b 
a 
b 
a 
b 
a 
b 
a 
b 
KC1 
105.6 
97.6 
93-8 
194.2 
KBr 
90.2 
K.I 
80.1 
75-o 
KOH 
103.9 
II6.4 
KSH 
(64)t 
65.1 
K2S04 
Na2S04 
344-5 
337-2 
kno3 
119-4 
III .4 
NaCl 
96-5 
NaBr 
85-7 
85.6 
Nal 
69.1 
70*3 
NaOH 
102.0 
in .8 
NaSH 
(55) 
60.4 
328.5 
329.C) 
NaNd3 
hi .2 
106.2 
LiCl 
LiBr 
9I*3 
Lil 
76.1 
LiOH 
H7-4 
LiSH 
66.0 
Li2S04 
334M 
34°-i 
LiN03 
hi .6 
112.2 
BaCl2 
196.3 
BaBr2 
SrBr2 
169.4 
157*7 
174-4 
173-8 
Bal2 
Srlo 
144.0 
143-3 
BaO-)H2 
Sr02H2 
216.3 
216.4 
226.6 
226.1 
BaS2H2 
SrH2S2 
124. I 
123.6 
BaS04 
SrS04 
337-5 
320.8 
BaN2( )g 
SrN206 
225.7 
219.8 
2l6.3 
215.2 
CaCf2 
MgCI* 
MnCU 
170.2 
I5I.C> 
187.6 
186.9 
CaBr2 
MgBr2 
MnBr2 
AlBr3 
141.2 
165.7 
165.0 
105.1 
Cal2 
Mgl., 
Mnl.) 
107.6 
135-3 
134-6 
75-7 
CaG2H2 
MgG2H2 
aio3h3 
214.7 
217.2 
2Q6.Q 
217.8 
CaSoHo 
MgSoH., 
MnS, nHoO 
46-3 
H5-3 
114.8 
CaS04 
MgS04 
A1qS3012 
3I9*9 
302.2 
323.0 
878.9 
CaN2Oe 
MgN208 
MnN2Oe 
203.2 
206.3 
147-5 
A1C1S‘ 
I2Q.6 
204. Q 
All, 
70.3 
159-3 
MnO, H.,0 
94-7 
ZnS, nHoO 
4i-5 
JVlnS04 
249.8 
263.5 
ZnN2O0 
132.3 
97.2 
90.9 
Znl2 
49*2 
60.5 
Z11O, H.,0 
82.6 
CdS, nHjO 
33-9 
ZnS04 
230.0 
248.4 
CdN2Os 
115.8 
CdCl." 
96.2 
CdBr» 
73-9 
74-3 
Cdl.> 
44-9 
43-9 
H20 (liquid) 
68 
3 
FeS, nHaO 
23.7 
CdS04 
221.1 
231.8 
FeN206 
CoN208 
ii9-5 
FeCI2 
99-9 
FeBr3 
78.0 
Fel2 
47.6 
FeO, H..O 
68.2 
CoS, nH20 
21.7 
b eS04 
235.6 
H4-3 
CoCI2 
NiClo 
CoBro 
72.9 
Col2 
42-5 
FeoOs. 3 H .,0 
191.1 
NiS, nH20 
19-3 
CoS04 
230.4 
NiN206 
113-2 
Wo 
93:7 
NiBr2 
FeBr3 
71.8 
Nil2 
41.4 
CdO, HoO 
65.6 
Ti.,S 
21.6 
NiS04 
229.3 
hdjo3 
49.1 
FeClg 
CoO, HoO 
63-4 
PbS 
20.4 
ii2so4 
221.0 
212.7 
PbN208 
TINOa 
HN03 
105.5 
97-9 
SnCl2 
HOI 
80.8 
8l . I 
39-3 
SnBr2 
TIBr' 
41.4 
... 
Snl2 
Til 
40.1 
NiO, H20 
SnO, H20 
60.8 
68.1 
CuS 
CuoS 
IO.O 
20.2 
H-.>S04 
PbS04 
216.2 
210.7 
58.1 
41-5 
48 . I 
127.2 
HBr 
28.3 
Pblo 
39-6 
Sn(J2, H2U 
H2O “(vapor) 
133-5 
HgS 
16.8 
H2S04 
192.9 
198.3 
CuN2C)8 
82.2 
T1C1 
PbBr2 
64.4 
54-4 
HgTo 
34-3 
58 
O 
Ag-,S 
5-3 
CuS04 
182.5 
HgN03 
38.9 
PbCl2 
HgCl 
82.7 
75-9 
HgBr 
HgBr2 
34-1 
Hgl 
24.2 
PbO 
AS203 
50-3 
147.0 
H-.y 
4-5 
9.2 
Ag2S04 
167.2 
162.7 
HgN208 
AgN03 
28.7 
41.2 
5° *5 
(Jill 
154.6 
KoS 
102.4f 
K„C03 
HgCl.j 
59-8 
CuBr 
24.9 
Cub, 
rioO 
42.2 
39-2 
113.2 
z79.5|2oo.o 
32.8 
CuBr2 
32-5 
20.4 
Agl 
13-8 
HgO 
30.6 
Na2S 
(88.4) 
103.9 
N a2CC3 
271 .O 
276.5 
CuCl2 
SbCl3' 
62.7 
SbBr3 
HI 
I3-1 
HgoO 
42.2 
LioS 
115.2 
BaC03 
281.3 
AgBr 
27.7 
ills 
10.5 
Cud 
37-1 
BaS 
(99-o) 
IO7. I 
SrC03 
279.6 
BiCl3 
90.6 
TlBra 
AsBr, 
56.1 
12.6 
CuoO 
40.8 
SrS 
(99-2) 
I06.6 
CaC03 
26^.2 
AgCl 
Aul 
—5-5 
TloOa, 3H2d 
86.0 
CaS 
(92.0) 
98-3 
MnC03 
209.2 
AsC13 
71-4 
HBr 
8.4 
HI 
—6.0 
PdO, H.,0 
22.7 
MgS 
(79-6) 
(no) 
CdC03 
179.9 
TiCl3 
89.0 
AuBr 
-0.1 
PtI, 
AgoO 
5-9 
A12S3 
(124.4) 
PbC03 
168.2 
HOI" 
AuCl 
22.0 
AuBr3 
PtBr2 
8.8 
5-o 
Au203, 3H.2O 
—13.2 
Ag2CC>3 
W.3 
5*8 
LioO 
I166.5 
22.8 
27.2 
(99.1)* 
PtCl2 
KoO 
164.5 
NaaO 
(100.2 
155.2 
BaO 
130.4 
158.2 
SrO 
130.9 
H57-7 
CaO 
i3i-7 
1149•4 
* According to Beketoff. 
f According to Sabatier, 
Heat of Formation of the Most Important Compounds of the Metals (according to f. Thomsen). 
In usual state of aggregation (columns a), and in dilute aqueous solution (columns b). INDEX. 
A, 
Absorptiometer, 119 
Acetylene, 151 
Acid radicals, 175 
Acids, 57, 82, 174 
Active oxygen, 83 
chlorine, 303 
Affinity, chemical, 23 
Agate, 235 
Alabandite, 388 
Algaroth, 144 
•Alkali earths, 299 
metals, 273 
Allotropy, 86 
Alloys, 259 
Alum, 354, 390 
burned, 354 
cubic, 354 
Roman, 354 
’ . shales, 354 
Alumina, 352 
Aluminates, 352 
Aluminite, 353 
Aluminium, 348 . 
bronze, 349 
chloride, 349 
oxide, 351 
sulphate, 353 
Alunite, 354 
Amalgams, 260 
Ammonia, 123 
thermo-chemistry, 145 
Ammonium, 127 
chloride, 296 
compounds, 295 
• salts, 127 
Anatase, 236 
Anglesite, 371 
Anhydrides of acids, 173 
Anhydrite, 304 
Anhydro-acids, 191, 212, 235 
Animal charcoal, 147 
Anthracite, 148 
Antimony, 142 
butter, 144 
chloride, 143 
hydride, 143 
oxide, 221 
sulphides, 223 
Antimonic acid, 222 
Antimonyl, 222 
Apatite, 304 
Apparatus of Carrd, 125 
Marsh, 140 
Aqua regia, 202, 345 
Argentan, 337, 407 
Arseniates, 220' 
Arsenic, 139 ' 
acid, 219 
chloride, 142 
sulphides, 220 
sulpho-salts, 221 
Arsenious oxide, 218 
Arsenites, 219 ' 
Arsine, 139 
Asbestos, 315 
Atmosphere, 116 
Atomic compounds, 171 
heat, 254 
volume, 252 
weights, 26, 164 
thermal, 257 
Atomicity. (See Valence,) 
Atoms, 68, 75, 163 
Aurates, 346 
Auric acid, 346 
Auric compounds, 345 
Auripigment, 220 
Aurous compounds, 345' 
Avogadro, law of, 75 
Azote, 117 
Azurite, 336 
B. 
Barium, 308 
chromates, 380 
peroxide, 309 426 
INDEX. 
Barium sulphate, 309 
Baryta water, 309 
Bases, 57, 82 * 
Basicity of acids, 174, 214 
Bauxite, 351 
Bell metal, 337 
Beryllium, 316 
Bessemer steel, 397 
Bismuth, 144, 372 
Bismuthic acid, 373 
Bismuthinite, 372 
Bleaching, 49, 185 
Bleaching lime, 302 
Blue, Prussian, 403 
Boiler incrustation, 306 
Boracite, 238 
Borax, 293 
Boron, 237 
fluoride, 239 
Brass, 337 
Braunite, 387 
Bromic acid, 179 
Bromine, 50 
hydrate, 50 
hydride, 59 
Bronze, 337 • 
Brown coal, 148 
Burned alum, 354 
C. 
Cadmium, 319 
Caesium, 284 
Calamine, 319 
Calcite, 305 
Calcium, 300 
carbonate, -305 
manganite, 391 
oxide, 300 
silicate, 306 
sulphate,-303 
sulphides, 307 
Calomel, 324 
Calorie, 64, 88 
Caput mortuum, 192, 400 
Carbon, 146 
chlorides, 157 
dioxides, 226 
disulphide, 231 
gas, 147 
group,147 
monoxide, 229 
oxychloride, 230 
Carbonates, 229 
Carbonic acid, 229 
Carry’s apparatus, 125 
Carnallite, 277 
Cast iron, 395 
steel, 396 
Cassiterite—Tin stone, 364 
Catalysis, 101 
Caustic potash,.276 
soda, 285 
Celsius’s thermometer, 87 
Cement, 301 
Cepientation steel, 396 
Cerite, 357 
Cerium, 357 
Cerussite, 368, 371 
Chalcedony, 235 
Chalk, 305 
Chamber acid, 189 
Charcoal, 147 
animal, 147 
Chemical affinity, 23 
elements, 20 
energy, 23 
equations, 28 
formulas and symbols, 25 
structure, 169 
Chemistry, definition of, 19 
Chili, saltpetre, 292 
Chloranhydrides, 192 
Chloric acid, 177 
Chlorides of sulphur, 109 
Chlorine, 46 
. gas, 48 
hydrates, 48 
oxides, 174 • 
water, 48 
Chloroform, 157 
Chlorous acid, 176 
Chlorthionyl, 185 
Chromates, 378 
Chromic acid, 379 
Chromite, 380 
Chromium alum, 378 
group, 373, 375 
oxychloride, 381 
yellow, 381 
Chrysoberyl, 352 
Cinnabar, 328 
green, 405 
Coal, anthracite, 148 
bituminous, 148 
brown’ 148 INDEX. 
427 
Cobalt, 404 
cyanides, 406 
rose, 406 
ultramarine, 405 
Cobaltic compounds, 405 
Cobaltite, 404 
Cobaltous compounds, 405 
Coke, 147 
Colcothar, 192, 400 
Colloids, 235 
Combination, types of, 247,322, 329, 
340, 374, 387 
Compounds, atomic, 171 
determined, 90 
endothermic, 65 
exothermic, 65 
molecular, 171 
Condensation of gases, 42 
Conservation of energy, 21 
Constant proportions, law of, 67 
Copper, 332 
alloys, 337 
hydride, 334 
vitriol, 335 
Corrosive sublimate, 326 
Corundum, 351 
Critical pressure, 88 
temperature, 227 
Crocoisite, 368, 375 
Cryolite, 291, 352 
Crystal combinations, 31 
Crystallization, water of, 91 
Crystallography, 28 
Crystalloids, 235 
Cubical alum, 354 
Cupellation, 338 
Cuprammonium compounds, 336 
Cupric compounds, 335 
Cuprite, 333 
Cuprous compounds, 333 
Cyanogen, 232 
compounds, 232, 278, 342, 402 
D. 
Davy’s lamp, 157 
'Decipium (Samarium), 357 
Dialysis, 234 
Diamond, 147 
Dicyanogen, 233 
Didymium, 357 
Diffusion of gases, 121 
Disodium phosphate, 292 
Dissociation, 92, 104, 137, 190, 204 
Disulphuric acid, 191 
Dithionic acid, 195 
Dcebereiner’s lamp, 45 
Drummond’s light, 81 
E. 
Earth metals, 348 
Earths, alkaline, 299 
rare, 356 
Ekaboron, 356 
Eka-aluminium, 359 
Electrolysis of salts, 267 
Elements, 20 
classification of, 37 
heterologous, 242 
homologous, 241 
Emery, 351 
Energy, conservation of, 21 
degradation of, 25 
entropy of, 25 
Epsom salt, 314 
Equivalence, 168 
of transformation, 22 
Equivalent weights, 168 
Erbia, 356 
Erbium, 357 
Ethane, 150 
Ethylene, 150 
Eudiometer, 119 
Euxenite, 356 
F. 
Feldspar, 348, 355 
Ferric acid, 401 
compounds, 401 
Ferricyanogen, 402 
Ferrocy anogen ,• 402 
Ferrous compounds, 398 
Flame, 151 
Bunsen, 155 
illuminating power of, 153 
non-luminosity of, 155 
oxidizing, 155 
reduction, 155 
Flint glass, 306 
Flores zinci, 318 
Fluorine, 53 428 
INDEX. 
Fluorite, 362 
Fulminating gold, 346 
silver, 339 
G. . 
Gadolinite, 356 
Gahnite, 352 
Galenite, 337, 367 
Gallium, 358 
group, 358 
Gas densities, 70 
laughing, 209 
volume, 71 
Gases, condensation of, 42 
diffusion of, 121 
kinetic theory of, 71 
. measuring of, 120 
Gersdorffite, 407 
Glass, 306 
Glauber’s salt, 287 
Glover’s tower, 188 
Glucinum, 316 
Gold, 343 
purple, 346 
fulminating, 346 
sulphide, 346 
Granite, 355 
Graphite, 147 
Greenockite, 320 
Guignet’s green, 336, 377 
Gypsum, 304 
* H. 
Halogen compounds of metals, 260 
hydrides, 54 
Halogens, 46 
Haloid acids, 57 
salts, 57 
Hausmannite, 387, 389 
Heat, atomic, 254 
latent, 88 
modulus, 64 
of solution, 91 
specific, 254 
unit of,-88 
Hematite, 395, 400 
Ilomologues, 150 
Hornsilver, 337, 341 
Hydrargylite, 351 
Hydrates, 82, 180, 261 
Hydrates of acids, 190 
Hydrobromic acid, 59 
Hydrocarbons, 148 
Hydrochloric acid, 54 
Hydrocyanic acid, 232 
Hydrofluoric acid, 62 
Hydrogen, 38 
chloride, 54 
cyanide, 232 
fluoride, 62 
iodide, 60 
peroxide, 97 
persulphide, 108 
sulphide, 105 
Hydroxides, 82, 180/261 
Hydroxylamine, 128 
Hypobromous acid, 178 
Hypochlorites, 176 
Hypochlorous acid, 176 
Hyponitrous acid, 210 
Hypophosphites, 212 
Hypophosphorous acid, 212 
Hyposulphites, 186 
Hyposulphurous acid, 185 
I. 
Ice machine of Carre, 125 
Indium, 359 
Iodic acid, 179 
Iodine, 51 
Iodine chloride, 66 
Iridium, 409, 411 
Iron, 395 
alum, 401 
cast, 395 
fused, 397 
group, 394 
pyrite, 40 x 
pyrophoric, 397 
vitriol, 399 
wrought, 396 
Isomerism, 86 
Isomorphism, 258 
K. 
Kaolin, 348, 355 
Kelp, 51 
Kermes mineral, 223 
Kieserite, 314 - INDEX. 
429 
■L. . 
Lanthanum, 347, 337 
Lapis lazuli, 355 
Laughing gas, 209 
Lead, 362, 367 
alloys, 368 
chamber crystals, 189 
oxide, 369 
peroxide, 369 
sulphate, 371 
sulphide, 371 
tree, 368 
white, 371 
Lime, 300 
chloride of, 301 
stone, 305 
water, 301 
Litharge, 369 
Lithium, 294 
Lunar caustic, 342 
Luteocobaltic chloride, 406 
M. 
Magisterium bismuthi, 373 
Magnesia, 312 • 
Magnesite, 315 
Magnesium, 312 
group, 310 
Magnetite, 400 
Malachite, 336 
Manganese, 386 
Manganic acid, 387 
salts, 391 
Manganite, 389 
Marble, 305 
Marsh gas, 148 
Marsh’s apparatus, 140 
Massicot, 369 
Maximum valence, 170 
Meerschaum, 315 
Mercuric compounds, 326 
Mercurous compoifnds, 324 
Mercury, 321, 323 
Meta-acids, 1*98 
Meta-antimonic acid, 222 
Metahydrates, 262 
Metalloids, 20, 37 
Metals, 251 
Metaphosphoric acid, 215 
Metastannic acid, 367 
Methane, 148 
Mineral waters, 90 
Molecular theory, 75 
compounds, 171 
Molecule, 74, 75, 163 
Molecules of elements, 75, 85, 104, 
132 
Molybdenite, 382, 383 
Molybdenum, 382 
Molybdic acid, 383 
Mortar, 301 
Mosaic gold, 367 
Multiple proportions, 69 
N. 
Nickel, 407 
plating, 408 
Niobium, 224 
Nitrates, 199 
Nitric acid, 199 
anhydride, 202 
fuming, 201 
oxide, 207 • 
Nitrogen, 114 
chloride, 129 
iodide, 130 
pentoxide, 202 
tetroxide, 204 
! Nitrosulphonic acid, 205 
Nitrosylchloride, 202 
Nitrosylsulphuric acid, 205 
Nitrous acid, 203 
Nordhausen sulphuric acid, 191 
.O. 
Olefiant gas, 150 
Olivine, 315 
Opal, 235 
Organic chemistry, 148 
Ores, 317 
Orthite, 356 
Ortho-acids, 198 
nitrioacid, 199 
phosphoric acid, 214 
Orthoclase, 355 
Osmium, 408, 410 
Oxidation, 81 
Oxides, 261 
acid-forming, 82 
basic, 82 430 
INDEX. 
Oxygen, 78 
Oxyhydrogen flame, 81 
Ozone, 83 
P. 
Palladium, 408, 411 
hydride, 46, 412 
Passive iron, 398 
Pattison’s method, 338 
Pentathionic acid, 196 
Perchloric acid,' 178 
Periodates, 180 
Periodic acid, 180 
Periodic system, 240 
Periodicity of spectral lines, 415 
thermo-chemical phenom 
ena, 249 
. valence, 247 
Permanent white, 309 
Permanganic acid, 386, 391 
Peroxides, 262 
Phlogiston, 116 
Phosgene, 230 
Phospham, 218 
Phosphates, 212 
Phosphine, 133 
Phosphites, 212. 
Phosphonium, 135 
iodide, 135 
Phosphoric acid, 214 
Phosphorite, 304 
Phosphorous acid, 213 
Phosphorus, 130 
bromides, 138 
chlorides, 136 
iodides, 138 • 
oxychloride, 137 
pentoxide, 212, 216 
sulphides, 217 
Photography, 341 
Pink salt, 366 
Platinum, 413 
bases, 415 
metals, 408 * 
ores, 409 
sponge, 45,409, 413 
Plumbic acid, 370 
Polyacids, 191, 235 
silicic acid, 235 
thionic acid, 194 
Porcelain, 355 
Potashes, 281 . 
Potassium, 275 
amide, 283 
aurate, 346 
bromide, 277 
chloride, 277 
cyanide, 278 
fluoride, 278 
hydrate, 276 
iodide, 277 
manganate, 391 
nitrate (saltpetre), 280 
oxide, 276. 
plumbate, 370 
Prussiate of potash, yellow, 402 
red, 403 
Puddle process, 396 
Puddle steel, 397 
Purpureocobaltic chloride, 406 
Pyrites, 187, 401 
Pyroacids, 191, 235 
Pyrolusite, 387, 390 
Pyrophosphoric acid, 214 
Pyrosulphuric acid, 191 
Q- 
Quadrant oxides, 263 
Quartz, 233 
Quicksilver, 323 
R. 
Radicals, 175 
Rare metals, 356 
Realgar, 220 
Reaumur’s.thermometer, 87 
Red chromate of potash, 380 
Red lead, 369 
Reduction, 82 
Respiration, 229 
Rhodium, 411 
Rhodochroisite,*387, 388 
Rock salt,- 286- 
Roseo-cobaltic chloride, 406 
Rpse’s metal, 373 
Rubidium, 284 
Ruby, 348 
Ruthenium, 408,410 
Rutile, 236 
Roman alum, 354 INDEX. 
431 
S. 
Safety lamp, 157 
Saltpetre, 280 
Salt producers, 46 
springs, 286 
Salts, 58,82, 175, 264 
basic, 264 
double, 264 
Samarium, 357 
Sapphire, 348 
Sassolite, 239 
Scandium, 356 
Scheele’s green, 336 
Scheelite, 384 
Schlippe’s salt, 224 
Schweizer’s reagent, 335 
Sea water, 92 
Selenic acid, 196- 
Selenium, no 
Silica, 233 
Silicates, 235 
Silicic acid, 234 
Silicon, 158 
chloride, 160 
chloroform, 160 
dioxide, 233 
disulphide, 231 
fluoride, 161 
hydride, 159 
Silver, 337 
cyanide, 342 
nitrate, 342 
oxides, 339 
Silvering, 343 
Slags, 396 
Smalt, 405 
Smaltite, 404 
Smithsonite, 318 
Soda, 290 
caustic, 285 
Sodium, 284 
carbonate, 290 
chloride, 286 
hydroxide, 285 
hyposulphite, 289 
iodate, 287 
nitrate, 292 
silicate, 294 
sulphate, 287 
oxide, 28.5 
phosphates, 292 
Soft solder, 368 
Solutions, 90 
supersaturated, 288 
thermo-chemistry of, 91 
Specific gravity of gases, 43, 70 
metals, 251 
heat, 254 
volume, 252 
Spectrum analysis, 416 
Spinels, 352, 389 
Stalactites, 306 
Stannates, 363 
Stannic compounds, 366 
Stannous compounds, 365 
Stassfurt salts, 275 
Status nascens, 50, 76 
Steel, 395 
Bessemer, 397 
cast, 396 
ingot, 397 
weld, 397 
Stibine, 143 
Stibnite, 142 
Stolzite, 384 
Stone coal-, 148 
Strass, 306 
Strontianite, 307 
Strontium, 307 
Structure, chemical, 169 
Substitution, 157 
Sulphates, 182 
Sulphides, 107 
Sulphites, 185 
Sulp'ho-stannates, 366 
Sulphur, 102 
chlorides, 109 
dioxide, 183 
heptoxide, 186 
liver, 283 
milk, 104, 307 
sesquioxide, 186 
trioxide, 186 
Sulphur auratum, 223 
Sulphuretted hydrogen, 105 
Sulphuric acid, 187 
anhydride, 186 
English, 189 
fuming, 191 
Sulphurods acid, 185 
Sulphuryl, 193 
chloride, 193 
Supersaturated solutions, 288 
Sylvanite, 343 432 
INDEX. 
Sylvite, 277 
Sympathetic ink, 405 
T. 
Talmi gold, 337 
Tantalum, 224 
Telluric acid, 197 
Telluride, hydrogen, 111 
Tellurium, ill 
Tension of vapors, 89 
Terbium, 357 
Tetrathionic acid, 195 
Thallic acid, 362 
Thallium, 360 
glass, 362 
Thenard’s blue, 405 
Thermo-chemistry, elements of, 25, 
64, 91, TOI, 113, .145. 158, 181, 
211, 224, 233, 249, 266, 274, 32O, 
331. 349. 361.363. 369 
Thionyl chloride, 185, 193 
Thiosulphuric acid, 195 
Thorium, 237 
Thulium, 26, 347 
Tin, 364 
and lead, 362 
salt, 365 
stone, 364 
Tinkal, 293 
Titanium, 236 
Tombac, 337 
Trithionic acid, 195 
Tungsten, 374, 384 
steel, 384 
Tungstic acid, 384 
Turnbull’s blue, 404 
Turpeth mineral, 328 
Type metal, 368 
U. 
Ultramarine, 355 
Uranates, 386 
Uranium, 374, 385- 
Uranium yellow, 386 
glass, 386 
Uranyl, 385 
V. 
Valence, 166 
Vanadium, 224 
Vapor density, 70 
Varec, 51 
Vitriol, iron, 399 
Vitriol, oil of (See Sulphuric acid), 
.>87 . 
Vivianite, 399 
W. 
Water, 87 
glass, 282, 294 
hard and soft, 90 
mineral, 90 
natural, 89 
of constitution, 314 
of crystallization, 91 
Weld iron, 397 
steel, 397 
White iron, 395 
lead, 371 
Woifram (Tungsten), 384 
Wollastonite, 306 
Wood’s metal, 373 
Wrought iron, 396 
Wtilfenite, 368 
Y. 
Yellow prussiate of potash, 402 
Ytterbium,'357 
Yttrium, 357 
Z. 
Zinc, 317 
blende, 318 
dust, 317 
sulphate, 318 
white, 317 
Zircon, 237 
Zirconium, 237 Catalogue No. 5. 
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OUR HOMES. With Illustrations. By Henry Hartshorns, m.d., of Phila- 
delphia, formerly Professor of Hygiene in the University of Pennsylvania. 
THE SKIN IN HEALTH AND DISEASE. By L. D. Bulkley, m.d., of 
New York, Physician to the Skin Department of the Demilt Dispensary and of 
the New York Hospital. 
SEA AIR AND SEA BATHING. By John H. Packard, m.d., of Phila- 
delphia, Surgeon to the Pennsylvania and to St. Joseph’s Hospitals. 
SCHOOL AND INDUSTRIAL HYGIENE. By D. F. Lincoln, m.d., of 
Boston, Chairman Department of Health, American Social Science Association. 
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3 
CHEMISTRY AND PHYSICS. 
RICHTER’S INORGANIC AND ORGANIC CHEMISTRY. 
Inorganic Chemistry, a Text-book for Students. By Prof. Yictor 
vox Richter, University of Breslau. Second American from 
Fourth German Edition. Authorized Translation. By Edgar 
F. Smith, m.a., ph.d., Prof, of Chemistry, Wittenberg College, 
formerly in the Laboratories of the University of Pennsylvania. 
With 89 Illus. and a Colored Plate of Spectra. 12mo. Cloth, $2.00 
Edition has been thoroughly revised, in many parts rewritten, 
and is handsomely printed. 
From F. A. Genth, Prof, of Chemistry,and F. A. Genth, Jr., Ass't Prof, of Chemistry, University of Penn- 
sylvania. 
“ We have examined with much care the ‘Inorganic Chemistry ’ of Prof. Victor von Kichter, re- 
cently translated by Dr. E. F. Smith. Both theoretical and general chemistry are treated in such a 
clear and comprehensive manner that it has become one of the leading text-books fora University 
course in Germany. We are indebted to Dr. Smith for his translation of this excellent work, which 
may help to facilitate the study of chemistry in this country.” 
From Prof. B. Silliman, Yale College, New Haven, Conn.—“It is decidedly a good book, and in some 
respects the best manual we have.” 
From Prof. Sam'l. S. Green, Swarlhmore College, Penn'a.—“ I am of the Opinion that it is the best text- 
book of the kind I have seen. I shall recommend it to my classes.” 
From Prof. A. A. Bennett, Chicago University.—-“I am satisfied this work is the best that I have yet 
seen, and that it will in a high degree till the want.” 
From E. H. S. Bailey, University of Kansas, Lawrence.—“ Dr Smith has, by his excellent translation 
brought into prominence one of the best and most recent books upon the science of chemistry.” 
CHEMISTRY OF THE CARBON COMPOUNDS, or, 
Organic Chemistry ; a complete Text-book and Laboratory Guide 
for Students. Authorized Translation from the Fourth German 
Edition. Illustrated. Cloth, $3.00 
Richter’s Chemistry is recommended at a number of prominent 
Schools and Colleges. Complete Descriptive Circulars sent free, upon 
application. 
BLOXAM. CHEMISTRY; Inorganic and Organic. Fifth Edition. 
With Experiments. By Charles L. Bloxam, Professor of Chem- 
istry in King’s College, London, and in the Department for Artillery 
Studies, Woolwich. Fifth Edition. With nearly 300 Engravings. 
Cloth, $3.75 ; Leather, $4.75 
BLOXAM’S LABORATORY TEACHING. Fourth Edition. 
Progressive Exercises in Practical Chemistry. By Charles L. 
Bloxam, Professor of Chemistry, in King’s College, London, etc. 
Fourth Edition. With 89 Engravings. 12mo. Cloth, $1.75 
For Students commencing the study of Practical Chemistry. It 
contains:— 
1. A series of simple Tables for the analysis of unknown substan- 
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the aid of the Blowpipe. 2. A brief description of all the practically 
important single substances likely to be met with in ordinary analy- 
sis. 3. Simple directions and illustrations relating to Chemical 
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ration of the tests. 
STUDENT’S MANUAL OF PHYSICS. 
By Sylvanus P. Thompson, b.a., d.s.c., f.r.a.s., Professor of 
Experimental Physics in University College, Bristol. Preparing. 
BY THE SAME AUTHOR AND TRANSLATOR. 4 
P. BLAKISTON, SON & CO.'S 
WATTS’ MANUAL OF CHEMISTRY. 
Physical and Inorganic. By Henry Watts, b.a., f.r.s., Editor 
of the Journal of the Chemical Society ; Author of “ A Dictionary 
of Chemistry,” etc. With Colored plate of Spectra and 150 other 
Illustrations. 12mo. 595 pages. Cloth, $2.25 
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STAMMER. CHEMICAL PROBLEMS. 
By Karl Stammer. Translated from 2d German Edition, with 
explanations and answers added, by Prof. W. S. IIoskinson, a.m., 
Wittenberg College, Springfield, O. 12mo. Cloth, .75 
VALENTIN. QUALITATIVE CHEMICAL ANALYSIS. 
By Wm. G. Valentin, f.c.s. Revised and edited by W. R. 
IIodgkinson, pn.D., Professor of Chemistry, Royal Military Acad- 
emy, and H. M. Chapman, Assistant Demonstrator of Chemistry, 
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SUTTON’S SYSTEMATIC HANDBOOK OF VOLUMETRIC 
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stances by Measure, applied to Liquids, Solids, and Gases. Adapted 
to the requirements of Pure Chemical Research, Pathological Chem- 
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Public Analyst for the County of Norfolk. Fourth Edition. Re- 
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ALLEN. COMMERCIAL ORGANIC ANALYSIS. 
A Treatise on the Modes of Assaying the Various Organic Chemi- 
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“We doubt if any other chemical work containing so large an amount of information could be 
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BOWMAN. PRACTICAL CHEMISTRY. 
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MUTER’S MEDICAL AND PHARMACEUTICAL CHEMISTRY. 
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TRIMBLE. PRACTICAL AND ANALYTICAL CHEMISTRY. 
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BEASLEY’S BOOK OF THREE THOUSAND PRESCRIPTIONS. 
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BEASLEY’S POCKET FORMULARY AND SYNOPSIS OF TIIE 
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By Professor Friedrich Fluckiger, of Strasburg. Translated 
by Frederick B. Power, pii.d., formerly Professor of Chemistry, 
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and Pharmacy, University of Wisconsin. With 8 Lithographic 
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*** “The Cinchona Barks are perhaps the most prominent medicinal agents now employed, and 
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SWERINGEN’S DRUGGISTS’ READY REFERENCE BOOK. 
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tions, etc., etc. By Hiram Y. Sweringen, m.d. 8vo. 
Cloth, $3.00 ; Leather, $4.00 
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Journal. 
“ It will prove of great service to the pharmaceutical student, apprentice, pharmacist, druggist and 
physician, as a book of ready reference and as an aid to the study of scientific works.”—American 
Journal of Pharmacy. 
BENTLEY AND TRIMEN’S MEDICINAL PLANTS. 
Price Reduced.—In order to bring this Valuable Work more icithin the 
reach of Pharmacists, the publishers have determined to reduce the price to 
$1.50 per part, in place of $2.00, the old subscription price, and $75.00 
complete, in 4 Vols., half morocco, in place o/$90.00. 
Containing full botanical descriptions, with an account of the 
properties and uses of the principal plants employed in medicine, 
especial attention being paid to those which are officinal in the 
British and United States Pharmacopoeias. The plants which supply 
food and substances required by the sick and convalescent are also 
included. By R. Bentley, f.r.s., Professor of Botany, King’s 
College, London, and H. Trimen, m.b., f.h.s., Department of 
Botany, British Museum. Each species illustrated by a colored 
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familiar with the Botany of Medicinal Plants.”—Druggists' Circular. 
“The work when complete (it is now complete) will be the most valuable compend of Medical 
Botany ever published.”—Boston Journal of Chemistry. 
BIDDLE’S MATERIA MEDICA. Ninth Edition. 
Materia Medica. For the Use of Students and Physicians. By 
the late Prof. John B. Biddle, m.d., Professor of Materia Medica 
in Jefferson Medical College, Philadelphia. The Ninth Edition, 
thoroughly revised, and in many parts rewritten, by his son, Clem- 
ent Biddle, m.d., Assistant Surgeon, U. S. Navy, assisted by 
Henry Morris, m.d. Containing all the additions and changes 
made in the last revision of the United States Pharmacopoeia. 
The Botanical portions have been curtailed or left out, and the other 
sections, on the Physiological action of Drugs, greatly enlarged. 
Octavo. Cloth, $4.00; Leather, $4.75 
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already so useful a work, both to the student and physician.”—Phila. Medical atid Surgical Reporter. 
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and yet sufficiently comprehensive. His style is clear and yet. succinct. He covers the ground—covers 
it well, and cumbers his work with nothing superfluous.”—Atlanta Medical and Surgical Journal. 
“One thing that particularly recommends this work to the student is, that the book is not so large 
as to discourage and cause him to feel that it is impossible for him to get over it and so much else in 
the short time before him.”—St. Louis Medical and Surgical Journal. 
“ It contains, in a condensed form, ail that is valuable in materia medica, and furnishes the medical 
student with a complete manual on this subject.”—Canada Lancet. 8 
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PROCTER’S LECTURES ON PRACTICAL PHARMACY. 
By Prof. Bernard S. Procter. Second Edition, with addi- 
tions and corrections. 44 Wood Engravings and 32 Fac-simile 
Prescriptions. Octavo. Cloth, $4.50 
ROBERTS’ MATERIA MEDICA AND PHARMACY. 
Containing many Valuable Tables. Directions for Preparing, 
with the Constituents of, numerous Modern and Useful Prepara- 
tions, Extracts, etc., etc. By Frederick Roberts, Examiner of 
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etc. 12mo. * Cloth, $2.00 
MERRELL’S DIGEST OF MATERIA MEDICA AND PHAR- 
MACY. 
Forming a complete Pharmacopoeia for the Use of Physicians, 
Druggists and Students. By Albert Merrell, m.d., Member of 
the State Board of Health of Missouri; Professor of Chemistry, 
Pharmacy and Toxicology in the American Medical College. St. 
Louis. Octavo, 512 pages. Half dark Calf, Red edges, $4.00 
“In clearness of style, simplicity of classification, and preciseness of directions, the treatise is 
indeed a model, and will be found alike instructive and useful for the physician and druggist.”— 
American Chemical Review. 
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of arrangement lends worth to the practitioner, dispensing pharmacist and the reviewer.”—SI. Louis 
Medical Journal. 
TUSON’S VETERINARY PHARMACOPOEIA. 
Including the Outlines of Materia Medica and Therapeutics ; for 
the Use of Practitioners and Students of Veterinary Medicine. 
By Richard J. Tuson, Professor of Chemistry, Materia Medica 
and Toxicology, at the Royal Veterinary College of England. 
Third Edition. Cloth, $2.50 
PEREIRA’S PHYSICIANS’ PRESCRIPTION BOOK. 16th Edition. 
Containing Lists of Terms, Phrases, Contractions and Abbrevia- 
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ceutical Terms, etc. By Jonathan Pereira, m.d. Sixteenth 
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OLDBERG. PRESCRIPTION BOOK. 300 New Prescriptions. 
Three Hundred Prescriptions, Selected Chiefly from the Best Col- 
lections of Formulae used in Hospital and Out-patient-practice, with 
a Dose Table, and a Complete Account of the Metric System. By 
Oscar Oldberg, phar.d., Late Medical Purveyor, United States 
Marine Hospital Service; Professor of Materia Medica, National 
College of Pharmacy, Washington, D. C. ; Member of the American 
Pharmaceutical Association, and of the Sixth Decennial Committee 
of Revision and Publication of the Pharmacopoeia of the United 
States. 12mo. Paper Covers, .75 ; Cloth, $1.25 
The prescriptions given in this work are selected from the Pharma- 
copoeias and formularies of the great Hospitals of New York, Phila- 
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medical officers of the United States Service. The Dose Table in- 
cludes nearly all of the remedies that have a place in the current 
Materia Medica. GENERAL AND SCIENTIFIC BOOKS. 
9 
WYTHE’S DOSE AND SYMPTOM BOOK. Eleventh Edition. 
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Cloth, $1.00 ; Leather, with Tucks and Pocket, $1.25 
“ The chapter on Dietetic Preparations will be found useful to all practicing physicians, most of 
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sick.”—Boston Medical and, Surgical Journal. 
“ Many a hard-worked practitioner will find it a useful little work to have on his study table.”— 
Canada Medical and Surgical Journal. 
A PHARMACOPCEIA OF SELECTED REMEDIES, with Thera- 
peutic Annotations, Notes on Alimentation in Disease, Air, Mas- 
sage, Electricity and other Supplementary Remedial Agents; and 
a Clinical Index ; arranged as a Handbook for Prescribers. By 
Edmund A. Kirby, m.d. Sixth Edition, Revised and Enlarged. 
With Illustrations. Cloth, $2.25 
ON THE VALUE OF PHOSPHORUS AS A REMEDY FOR LOSS 
OF NERVE POWER, MELANCHOLIA, NEURALGIA and 
other Functional Disorders of the Nervous System induced by 
Overwork and the Exigencies of Modern Life. With Formulse and 
Directions for Treatment. Fifth Edition. Octavo. Boards, $1.00 
TANNERS’ MEMORANDA OF POISONS and their Antidotes and 
Tests. Fifth American, from the Last London Edition. Revised 
and Enlarged. Cloth, .75 
This most complete Toxicological Manual should be within reach 
of all physicians and pharmacists, and as an addition to every family 
library, would be the means of saving life and allaying pain when 
the delay of sending for a physician would prove fatal. 
THE SAME AUTHOR 
YEO’S MANUAL OF PHYSIOLOGY. 
A Text-book for Students. By Gerald F. Yeo, m.d., f.r.c.s., 
Professor of Physiology in King’s College, London. With over 300 
carefully printed Illustrations and a Complete Glossary and Index. 
Crown Octavo. Cloth, $4.00 ; Leather, $5.00 
“ After a careful examination of this manual of Physiology, I can truthfully say that it is a most 
valuable addition to the list of text-books upon this subject. That it should and will receive a welcome 
from both students and teachers there can be no doubt: for, in addition to the familiar but well 
presented facts of most text-books, it contains all the more important facts of physiological science 
which have been established in the last few years. The author presents his subject in a manner that 
is clear, concise and logical. Each section has had a careful revision, and reveals the author’s famili- 
arity with the scope and tendencies of modern physiology. It will prove an interesting and instruc- 
tive book to those commencing the study of this subject.”—A. P. Brubaker, Jefferson College, Philadel- 
phia. 
“ We have pleasure in recommending this book, as a most excellent manual, being what it pretends 
to be—elementary, and yet containing all that is really of importance to the student.”—Medical Times 
and Gazette. 
“There are many points in physiology that are either not comprehended or are misunderstood by 
the great majority of students. In this work these points are made especially clear. We have had 
long experience in teaching this branch of medical science, and unreservedly commend this work to 
the student of physiology.”—Archives of Dentistry. 
“For students’ use it is one of the very best text-books in Physiology.”—Prof. L. B. How, Dart- 
mouth College, Hanover, N. H. 
PHYSIOLOGY. 10 
P. BLAKISTON, SON & CO.'S 
THE MICROSCOPE. 
BEALE’S HOW TO WORK WITH THE MICROSCOPE. 5th Ed. 
A Complete Manual of Microscopical Manipulation, containing a 
full description of many new processes of investigation, with direc- 
tions for examining objects under the highest powers, and for taking 
photographs of microscopic objects. Fifth Edition. Containing 
over 400 Illustrations, many of them colored. 8vo. Cloth, $7.50 
“The Encyclopaedic character of Dr. Beale’s well known work on the Microscope renders it im- 
possible to present an abstract of its contents; suffice it to say, that anything in his department upon 
which the physician can desire such information will be found here, and much more in addition. It 
is, moreover, a storehouse of facts, most valuable to the physician, and is indispensable to every one 
who uses the microscope.”—American Journal of Medical Science. 
BEALE’S USE OF THE MICROSCOPE IN' PRACTICAL MEDI- 
CINE. Fourth Edition. 
For Students and Practitioners, with full directions for examining 
the various secretions, etc., in the Microscope. Fourth Edition. 
500 Illustrations. Much enlarged. 8vo. Cloth, $7.50 
“As a microscopical observer, and a histological manipulator, his (Dr. Beale) skill and eminence 
are generally conceded.”—Popular Science Monthly. 
CARPENTER ON THE MICROSCOPE. Sixth Edition. 
The Microscope and its Revelations. By W. B. Carpenter, m.d., 
f.r.s. Sixth Edition. Revised and Enlarged, with over 500 Illus- 
trations. Several Lithographic Plates. Cloth, $5.50 
“As a text-book of Microscopy in its special relation to natural history and general science, the 
work before us stands confessedly first, and is alone sufficient to supply the wants of the ordinary 
student.”—American Journal of Microscopy. 
MACDONALD’S MICROSCOPICAL EXAMINATION OF WATER 
AND AIR. 
A Guide to the Microscopical Examination of Drinking Water, 
with an Appendix on the Microscopical Examination of Air. By 
J. D. Macdonald, m.d. With Twenty-five Full-page Lithographic 
Plates, Reference Tables, etc. Second Edition, 8vo. Cloth, $2.75 
THE MICROTOMIST’S VADE-MECUM. 
A Handbook of the methods of Microscopic Anatomy, comprising 
upwards of Five Hundred Formulae and Methods, collected from 
the practice of the best workers. By Arthur Bolles Lee. 
Crown 8vo. Cloth, $3.00 
WYTHE, ON THE MICROSCOPE. 
The Microscopist. A Manual of Microscopy and Compendium 
of the Microscopic Sciences, Micro-Mineralogy, Micro-Chemistry, 
Biology, Histology, and Practical Medicine. By Joseph II. Wythe, 
a.m., m.d. Fourth Edition. 252 Illustrations. 8vo. 
Cloth, $3.00 ; Leather, $4.00 
An Index and Glossary, with notices of recent additions to the 
microscope, together with the genera of microscopic plants, have 
been given in an Appendix. 
“Tho author very carefully brings out every necessary fact and principle relating to the use of the 
microscope, and now that this instrument has become an essential part of every practitioner's arma- 
mentarium, a practical guide and reference book is also a necessity, and we are fully warranted in 
reiterating the statement that this is one of the most valuable text-books ever offered to students and 
practitioners of medicine.”—The Cincinnati Lancet and Clinic. GENERAL AND SCIENTIFIC BObKS, 
11 
MARTIN’S MANUAL OF MICROSCOPIC MOUNTING. 
With Notes on the Collection and Examination of Objects, and 
upwards of 150 Illustrations. By John H. Martin. Second Edi- 
tion, Enlarged. 8vo. Cloth, $2.75 
INDIGESTION, HEALTH, HEADACHES, ETC. 
4®“See also List of “Health PrimersPage 
BEALE ON SLIGHT AILMENTS. New Edition. Just Ready. 
Slight Ailments, Their Nature and Treatment. By Lionel S. 
Beale, m.d., f.r.s., Professor of Practice of Medicine in King’s 
College, London. Second Edition. Enlarged and Illustrated. 
Cloth, $1.25 ; Paper Covers, .75 
Better Edition, Heavy Paper. Extra Cloth, $1.75 
Introductory. The Tongue in Health and Slight Ailments. Appetite. Nausea. Thirst. Hunger. 
Indigestion, its Nature and Treatment. Constipation, its Treatment. Diarrhoea. Vertigo. Giddi- 
ness. Biliousness. Sick Headache. Neuralgia. Rheumatism. The Feverish and Inflammatory 
State. Of the Actual Changes in Fever and Inflammation. Slight Inflammations, etc., etc. 
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GILL. ON INDIGESTION. Third Edition. 
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These are invaluable little treatises upon subjects that enter pain- 
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WRIGHT, ON HEADACHES. Ninth Thousand. 
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POTTER ON SPEECH, AND ITS DEFECTS. 
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ANSTIE, STIMULANTS AND NARCOTICS. 
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ACTON. THE REPRODUCTIVE ORGANS. 
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HYGIENE. 
PARKE’S PRACTICAL HYGIENE. Sixth Edition. 
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FRANKLAND’S WATER ANALYSIS, 
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BIBLE HYGIENE ; 
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WILSON, ON DRAINAGE. 
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copy ."—Builder and Wood Worker. 
GENERAL AND SCIENTIFIC BOOKS. 
15 
BY SAME AUTHOR. 
NAVAL HYGIENE. 
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SEPULTURE : 
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Introduction. History of Sepulture. Ancient Customs and Methods. Sepulchres. Interments 
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GENERAL AND SCIENTIFIC BOOKS. 
WILSON’S TEXT-BOOK OF DOMESTIC HYGIENE AND SANI- 
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HEALTH RESORTS. 
MADDEN’S HEALTH RESORTS FOR CHRONIC DISEASES. 
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WILSON’S SEA VOYAGES FOR HEALTH. 
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THE CARE OF CHILDREN. 
HALE. OH THE MANAGEMENT OF CHILDREN IN HEALTH 
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